Basic chemistry - Ross University

Transcription

Department of Biochemistry, Ross University School of Medicine
Preparation for incoming students
Basic Chemistry: A Reminder
Engelbert Buxbaum,
Dr. rer. nat., Dipl. Biol.
February 2010
c 2002-2010 by Dr. Engelbert Buxbaum. It may be freely used for
This text is educational and non-commercial purposes. Distribution is allowed as long as the
text is distributed completely, including this copyright notice, and without fee. All
other rights reserved.
All trademarks are acknowledged as trademarks of their respective owners.
Although great care has gone into the preparation of this text, it is not intended
to give guidance for the treatment of humans. Any information given should be
confirmed independently before being used for this purpose.
Please report any errors found in this text or any suggestions for its improvement
to me ([email protected]).
The contents of this text should be familiar to you from high school and your
premedical training. If you have problems with this material, textbooks on general
and organic chemistry cover it in more detail.
This text was created with LATEX using the MikTEX system
(http://www.dante.de). Chemical structures were created with ISIS-Draw
(http://www.mdli.com), drawings with XFig (http://www.xfig.org) under Linux. Gnuplot (http://www.cs.dartmouth.edu/gnuplot_info.html)
was used for plotting mathematical functions. For conversion between graphic formats the Gnu Image Manipulation Package (Gimp,
http://www.xcf.berkely.edu/∼gimp/gimp.html) was used. Three-dimensional
graphics were rendered with the Persistence Of Vision Ray-tracer (POV-Ray
http://www.povray.org.
Many thanks to all those who made these free tools available on the net.
Contents
1 Physical chemistry
1.1 Thermodynamics . . . . . . . . . . . . . . . . .
1.2 The law of mass action . . . . . . . . . . . . . .
1.2.1 Maintaining the equilibrium constant . .
1.2.2 Le Chatelier’s principle . . . . . . . .
1.3 Energy conversion during chemical reactions . .
1.3.1 Free energy . . . . . . . . . . . . . . . .
1.3.2 Enthalpy and entropy . . . . . . . . . .
1.4 Electrochemistry and Redox-reactions . . . . . .
1.4.1 Thermodynamics of electrochemical cells
1.4.2 General Redox-reactions . . . . . . . . .
1.4.3 Biological applications . . . . . . . . . .
1.4.4 Technical applications . . . . . . . . . .
1.5 Water, acids, bases, buffers and pH . . . . . . .
1.5.1 The auto dissociation of water . . . . . .
1.5.2 The pH of a solution . . . . . . . . . . .
1.5.3 Strong and weak acids and bases . . . .
1.5.4 Buffers . . . . . . . . . . . . . . . . . . .
1.5.5 Indicators . . . . . . . . . . . . . . . . .
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2 Atom structure and chemical bonds
2.1 Atoms and elements . . . . . . . . . . . . . . . . . . . . . . . . . . .
2.2 Basic quantum theory . . . . . . . . . . . . . . . . . . . . . . . . . .
2.2.1 The first row of the periodic system: H and He . . . . . . . . .
2.2.2 The second row of the periodic system: Li, Be, B, C, N, O, F, Ne
2.2.3 Electronegativity . . . . . . . . . . . . . . . . . . . . . . . . .
2.3 Quantum theory II: Orbital hybridisation . . . . . . . . . . . . . . . .
2.3.1 The sp-hybrid orbital . . . . . . . . . . . . . . . . . . . . . . .
2.3.2 The sp2 -hybrid orbital . . . . . . . . . . . . . . . . . . . . . .
2.3.3 Nitrogen and oxygen: Hybrid orbitals with 2 electrons . . . . .
2.3.4 σ and π-Bonds . . . . . . . . . . . . . . . . . . . . . . . . . .
2.3.5 Delocalised π-bonds: Mesomery . . . . . . . . . . . . . . . . .
2.4 Weak interactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
2.4.1 Hydrogen bonds . . . . . . . . . . . . . . . . . . . . . . . . . .
2.4.2 van der Waals bonds . . . . . . . . . . . . . . . . . . . . .
2.4.3 Hydrophobic interactions . . . . . . . . . . . . . . . . . . . . .
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Contents
2.4.4
Transition metals and complex bonds . . . . . . . . . . . . . . 42
3 Basic organic chemistry
3.1 Important functional groups . . . . . . . . . . . . .
3.1.1 Alkanes . . . . . . . . . . . . . . . . . . . .
3.1.2 Oxidation products . . . . . . . . . . . . . .
3.1.3 Reaction products of alcohols, aldehydes and
3.1.4 Acetals and ketals . . . . . . . . . . . . . . .
3.1.5 Nitrogen containing compounds . . . . . . .
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4 Stereochemistry
4.1 Chirality . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
4.1.1 Chiral compounds and polarised light . . . . . . . . . .
4.1.2 Biological and medical importance of chiral compounds
4.1.3 Nomenclature of chiral compounds . . . . . . . . . . .
4.2 Double bonds and cyclic compounds: cis/trans isomery . . . .
4.2.1 Triple bonds . . . . . . . . . . . . . . . . . . . . . . . .
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5 Appendix
5.1 List of symbols . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
5.2 Periodic System of the Elements . . . . . . . . . . . . . . . . . . . . .
5.3 Acronyms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .
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Thermodynamic properties are those features of a chemical reaction that are
related to its energy balance and equilibrium. Kinetic properties are those related
to the rate (velocity) of the reaction. Only the kinetic properties of reactions are
changed by enzymes.
1.1 Thermodynamics
The following definitions of important terms are oversimplified, but sufficient for the
purpose of medical biochemistry. It is important that you know them as they will
be frequently used to describe reactions that go on in our body.
Temperature is the random movement of molecules, and is measured in K. 0 K is
the lowest theoretically possible temperature, where not only all molecules, but even
the electrons of their atoms, would be at rest. In practice, this temperature can never
be reached. Note that temperatures are measured in Kelvin (K), but temperature
differences in ◦ . There is no such thing as ◦ K. Alternatively, temperatures may also
be measured in degrees Celsius ◦C (confusing, isn’t it?). You convert from ◦C to K
by adding 273.16. In formulas, the capital T represents temperature. In the US the
unit ◦ F was used in the past, but you should avoid any non-SI units, they make a
lot of additional work and are a source of calculation error.
Heat (q) flows between bodies of different temperature, with q = C ×∆T . C is the
heat capacity of the bodies (amount of energy you need to increase the temperature
by a given amount), and ∆T the difference in temperature between them. The Greek
letter ∆ (Delta) always symbolizes differences. In formulas, the heat leaving a body
is denoted with a negative, heat entering a body with a positive sign. For an example
see fig. 1.1. Heat is measured in Joule (J).
Work w is anything that “somehow” can be converted to the lifting of weights.
When work is transferred across system boundaries then work done on a system gets
a positive, system done by a system a negative sign. The unit for work is the Joule
(J).
Energy is the sum of heat and work:
E = q + w = const
(1.1)
In other words, energy can not be created nor destroyed. This statement is called
the first law of thermodynamics. Since both heat and work are measured in J,
energy of course has that unit too.
5
Figure 1.1: Heat flowing between bodies of different temperatures. For explanations
see text.
40 °C
60 °C
20 °C
0 °C
30 °C
30 °C
Figure 1.2: From the left to the right: open, closed and insulated (adiabatic) systems.
open
6
closed
insulated
1.1 Thermodynamics
A system is some part of the universe, which is separated from the rest (see fig.
1.2). We distinguish:
Open systems can exchange both matter and energy with the rest of the environment. Example: an open test tube.
Closed systems can exchange energy, but not matter with the environment. A
stoppered test tube would be an example, matter can no longer leave or enter,
but we may still heat its content.
Insulated systems (also called “adiabatic” systems) can exchange neither matter
nor energy with the environment. If we perform our reaction in a Dewarvessel, energy exchange would be prevented. Such vessels have a double-layer
construction, with vacuum as insulator between the inner and outer glass bottle to minimize conduction and convection. The bottles have a silver coating
on the vacuum side, to limit heat transfer by radiation. You may have used
such vessels to keep your coffee hot.
Standard conditions: Some thermodynamic properties of a system depend on
the conditions that system is in. Chemists have agreed to reference their measurement to a pressure of 1014 hPa (standard atmospheric pressure), a temperature of
25 ◦C (= 298 K) and a concentration of 1 M for all reactants. The latter point is an
inconvenience for biochemists, because if [H+ ] = 1 M, then pH= 0 (see later). Very
few biochemical reactions occur at such low a pH, so biochemical standard conditions are pH 7. Biochemists also define standard values for ionic strength (250 mM)
and [Mg2+ ] (3 mM), since these influence the activity of many enzymes. The values
chosen simply reflect those that we find in living cells. If a parameter X is measured
under chemical standard conditions, we write X 0 , if it was measured under biological
standard conditions we write X 00 . In older books you may find X 0 ’ instead, because
IUPAC changed the nomenclature several years ago.
Enthalpy (H) is the heat of a chemical reaction. It is expressed in Joule/mol
(J/mol). Enthalpy itself can not be measured, but changes in enthalpy (∆H) can.
To do these measurements, we have to assign the value 0 J/mol to some arbitrary
state. By convention, this value is assigned to pure elements in their stable state
under standard conditions.
If heat is produced by a reaction, we give ∆H a negative sign and call the reaction
exothermic. If enthalpy is consumed during a reaction, we give ∆H a positive sign,
such reactions are called endothermic.
The change in enthalpy is given by:
∆H = ∆E + P × ∆V
(1.2)
where P is the pressure (in Pa) and ∆V the change in volume. The product P × ∆V
is the mechanical work done during a reaction. Because biochemical reactions
usually happen in dilute aqueous solutions, ∆V ≈ 0 and hence ∆H ≈ ∆E. This is
a useful result because it is difficult to measure a change in energy, but much easier
7
to measure changes in enthalpy. For biochemical reactions we can thus use ∆H as
an approximation for ∆E, which is the much more important parameter.
Entropy (S) is a measure of chaos. If you look at fig. 1.1, you see two bodies
of different temperature. If you bring them close together, heat could theoretically
flow from the cold body to the hot, making the cold body even colder and the hot
body even hotter. The first law of thermodynamics would allow that, because the
energy lost by the cold body exactly balances the energy gained by the hot. However,
everyday experience shows us that this is not what will happen. Instead, heat will
flow from the hot body to the cold, until both have the same temperature. This is
an example for a general principle called the second law of thermodynamics:
The direction of a reaction is the one where the entropy (disorder) of the universe
is maximized, that is ∆S ≥ 0.
Note: The second law of thermodynamics allows for a local reduction of entropy
as long as the global entropy increases. Life does exactly that: Our sun sends out
energy in form of photons, constantly increasing universal entropy. Plants catch
this orderly stream of photons and dissipate it, again increasing the entropy of the
universe. However, part of this entropy difference is not given up to the universe, but
saved in the form of organic compounds like sugar, which animals then consume.
There is a third law of thermodynamics which states that the entropy of an
ideal crystal at a temperature of 0 K is 0 J mol−1 K−1 . This law allows us to measure
entropies, but is of no other consequence to biochemistry. You may safely assign it
to passive knowledge.
Water below 100 ◦C is liquid, above that temperature it is gaseous, but we have to
expend energy to convert water of 100 ◦C into steam of 100 ◦C (the heat of evaporation). Gases, of course, have a much higher entropy than liquids. Thus energy,
temperature and entropy of a reaction must be linked. This linkage is provided by
the Gibbs free energy ∆G through the Gibbs-Helmholtz-equation:
∆G = ∆H − T ∆S
(1.3)
∆G is a very important parameter, which we will encounter again and again in this
course. It determines the direction of a chemical reaction:
∆G < 0: reaction proceeds from left to right (v+ > v− )
∆G = 0: reaction is at equilibrium (v+ = v− )
∆G > 0: reaction proceeds from right to left (v+ < v− )
By convention, chemical reactions are always written in the direction where ∆G is
negative.
∆G is that part of ∆H of a reaction that can be used to do useful work. T ∆S
is the part of ∆H which is lost as heat. Therefore ∆G determines the energy efficiency of a reaction. This is, of course, a key driving force in evolution since any
energy not lost as heat may be used for growth and reproduction. Thus it is not
8
1.2 The law of mass action
surprising that the energy efficiency of living organisms is close to the thermodynamically allowed maximum, and often much higher than that of man-made devices.
Physical processes in which chemical bond energies don’t change (∆H = 0, for
example diffusion across a membrane) are also driven by ∆G via an increase in
entropy (i.e., a negative value for [T × ∆S]). The body has to use chemical bond
energy to combat the tendency for increasing entropy. That’s what metabolic energy
is good for!
Summary:
Exothermic reaction:
∆H < 0, heat is released.
Endothermic reaction: ∆H > 0, heat is absorbed
Exergonic reaction:
∆G < 0, reaction can proceed
Endergonic reaction:
∆G > 0, reaction cannot proceed.
Equilibrium:
∆G = 0 concentration of reactants doesn’t change over time
For a chemical reaction to occur between two molecules, they have to come together
first. This is the more probable the higher the concentration of the reactants is. For
example for the reaction
A+B *
(1.4)
)C +D
the reaction velocity would be proportional to the concentration of A and B.
By convention concentration of substances are indicated by putting square brackets around their symbols, instead of “concentration of A” we can write “[A]”. Thus
v1 ∝ [A]
v1 ∝ [B]
v1 = k1 ∗ [A] ∗ [B]
(1.5)
(1.6)
(1.7)
Similarly, we can write for the back-reaction:
v2 ∝ [C]
v2 ∝ [D]
v2 = k2 ∗ [C] ∗ [D]
(1.8)
(1.9)
(1.10)
If we mix the reactants A and B, the initial velocity of the forward reaction will
be high. As the reaction proceeds, [A] and [B] will decrease, this will also decrease
the reaction velocity. On the other hand, [C] and [D] will increase, this will increase
the velocity of the reverse reaction (which was 0 initially, because the concentrations
of C and D were 0). After some time, the velocities of the forward and back reaction
will become equal:
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k1 ∗ [A] ∗ [B] = k2 ∗ [C] ∗ [D]
k1
[C] ∗ [D]
=
= Keq
k2
[A] ∗ [B]
(1.11)
(1.12)
In other words, the concentrations of the reactants A, B, C and D will no longer
change, even though reactions still occur. This situation is called equilibrium. The
ratio k1 /k2 is called equilibrium constant Keq .
In theory, all chemical reactions are reversible. If only non-covalent interactions are involved, as in protonation-deprotonation, antigen-antibody binding and
hormone-receptor binding, the reactions are always found at equilibrium concentrations. Chemical reactions in which covalent bonds are formed and broken are
also reversible, but there is an energy barrier to the reaction. Therefore the reactions
are not always found at equilibrium concentrations.
The exact form of the equation depends on the number of reactants on either side
of the reaction. For example:
* B
2A )
v1 = k1 ∗ [A] ∗ [A] = k1 ∗ [A]2
v2 = k2 ∗ [B]
[B]
Keq =
[A]2
(1.13)
(1.14)
(1.15)
aA + bB + cC + ... *
) zZ + yY + xX + ...
[Z]z ∗ [Y ]y ∗ [X]x ∗ ...
Keq =
[A]a ∗ [B]b ∗ [C]c ∗ ...
(1.17)
(1.16)
In general:
(1.18)
Equation 1.18 is called the law of mass action and of fundamental importance
for our understanding of chemical reactions. You will encounter it again and again,
and it is worthwhile to consider the consequences of this law.
1.2.1 Maintaining the equilibrium constant
Example 1 Let us consider the bimolecular reaction which was our first
example:
A+B *
)C +D
(1.19)
Let us assume we had mixed the reactants A and B, and we allowed the reaction to proceed until equilibrium was obtained. Than we add some more A. What
will happen? Obviously, [C]∗[D]
6= Keq . Therefore more C and D will be formed,
[A]∗[B]
10
until a new equilibrium is obtained. The concentration of B will be lower, and the
concentrations of C and D higher than in the initial equilibrium.
This has important applications in technical chemistry. Let us assume that B is a
very expensive substance, while A is cheap. Normally one would add equal amounts
of A and B into the reaction vessel, but under these circumstances it makes sense
to use much more A than B, so that only little B will be left when equilibrium has
been obtained.
Example 2 Let us assume now that C is a volatile compound, and evaporates
rapidly from the reaction mixture. Under these conditions, [C] will be very low, and
the reaction will continue to proceed much further than when all of the reactant C
stayed in the reaction vessel.
Thus by adding one of the reactants in excess, or by removing a product from
the reaction, we can make better use of expensive compounds in chemical reactions.
These ideas are used not only in technical chemistry, but also in chemical pathways
in our own bodies.
Note that removal of a product can be achieved also by a second reaction, that
converts it into some other compound.
1.2.2 Le Chatelier’s principle
In the last section we have defined the equilibrium constant Keq . We have seen that
this constant is independent of the initial ratio of reactants. However, it is influenced
by environmental conditions.
Let’s consider example 2 a little further. If we let the reaction proceed in a closed
vessel, rather than an open test tube, the evaporating product C will cause an
increase in pressure. Under these conditions, Keq will be more on the side of the
educts, rather than the products. By artificially increasing the pressure further we
can shift the equilibrium even more to the educts, by applying a vacuum on the
other hand we can shift it to the products.
Many reactions cause a change in the temperature of the reaction system. Some
salts for example will cause a considerable decrease in temperature, when they are
dissolved in water. If we prepare a saturated solution of such a salt at room temperature and then heat it, solubility of the salt will increase. If on the other hand we
cool the solution in an ice bath, some of the salt will precipitate.
A substance that raises the temperature when dissolved in water will behave in
the opposite way.
These examples show a general phenomenon, known as Le Chatelier’s principle:
If the conditions of a chemical system are changed, the equilibrium constant will
change in such a way as to counteract the change in conditions.
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1.3 Energy conversion during chemical reactions
1.3.1 Free energy
In physics you have learned that energy, like mass, is invariant, that it can not be
destroyed nor created (First main law of thermodynamics). However, different
forms of energy can be converted into each other. Chemical energy is a form of
potential energy. For example a drop of methanol contains chemical energy, when it
is ignited this energy is converted into heat and light:
3
CH3 OH + O2 → CO2 + 2H2 O
2
∆G0 = −1656 kJ/mol
(1.20)
Under the special conditions of an engine the chemical may be converted into kinetic
energy. In a fuel cell, it can be converted into electrical energy.
Assume we mix two reagents, A and B, which react to products C and D. As
we have learned, the reaction will proceed and convert chemical energy into, say,
heat until equilibrium is obtained. No energy conversion occurs once equilibrium is
obtained. Thus we see that Gibbs’ free energy ∆G, depends on how far away
a system is from equilibrium.
The change in free energy (∆G) during a chemical reaction in an open system is
connected to the concentration of reactants by Gibbs’ equation:
C ∗D
0
(1.21)
∆G = ∆G + RT ln
A∗B
R is the universal gas constant 8.31 JK−1 mol−1 , K the absolute temperature (in
Kelvin, 25 ◦C is equivalent to 298.4 K). The sign of ∆G depends on whether energy
has to be spend to make the reaction happen (∆G is positive), or whether energy
is liberated during the reaction (∆G is negative). By convention, chemical reactions
are usually written in the direction that liberates energy.
1.3.2 Enthalpy and entropy
As we have seen the total amount of energy in an insulated system remains constant. However, not all forms of energy were created equal. If we burn methanol, its
chemical energy is converted completely into heat. If we use methanol as fuel in an
engine, some, but not all of its chemical energy can be converted into mechanical
energy. The rest is converted into heat. This is the reason that the engine in your
car needs to be cooled. It is on principle not possible to construct an engine, that
could convert all the chemical energy contained in its fuel into mechanical energy,
some energy “loss” in form of heat is inevitable.
If we heat a mixture of CO2 and H2 O we can produce some methanol. However, we
would have to spend more energy on heating this mixture, than we could claim from
the amount of methanol formed. So, although we can convert all forms of energy
into heat with 100% efficiency, in the reverse process invariably some of the heat will
12
1.4 Electrochemistry and Redox-reactions
be lost to the environment, and only the remainder can be converted into chemical
energy.
In some way energy in the form of heat is “less valuable” than other forms of
energy. “Value” here means the ability to do useful work. What is the reason for this
difference?
An every day example may clarify the situation: If you leave your room for a
couple of days without clearing up, it will look a mess. To make it look tidy again,
you have to spend energy. This is a general rule: Maintaining or increasing order
requires energy, if that energy is not spend, disorder will increase (all by itself,
unfortunately). This rule is called the second main law of thermodynamics.
There is a more scientific name for disorder: Entropy S, which has the unit of
energy divided by temperature, J/K.
Let us now return to the different forms of energy. When we heat up a system, we
increase the average speed with which the atoms or molecules in that system move.
Because this movement is random (Brownian motion), increasing the temperature
means increasing entropy.
On the other extreme, electrical energy consists of the flow of electrons from a
negative to the positive pole of a circuit.
You now see why it is impossible to convert energy in the form of heat quantitatively into electricity: Some of the energy has to be spend to increase order (or
reduce entropy).
In other words, the total free energy (∆G) of a chemical reaction consists of two
different items: Entropy (∆S) and enthalpy (∆H, the part of ∆G that can do useful
work). The exact relationship between the three is:
∆G = ∆H − T ∆S
(1.22)
with T again being the absolute temperature.
Most chemical reactions in the lab perform work by changing the pressure(P ) and/or
volume (V ) of the system:
P
(1.23)
∆G = nRT ln
P0
P V = nRT
(1.24)
with R = gas constant and T = absolute temperature. An important exception is
the electrochemical cell, which generates electrical current instead.
Each electrochemical cell consists of two half cells or electrodes:
D
+
A + e−
D + A+
*
) D + + e−
*
)A
* D+ + A
)
(1.25)
(1.26)
(1.27)
13
0.76 V
-
+
H2 , 1024 HPa
Bridge
(KCl agar)
Zn-electrode
1 M HCl
Hydrogen standard
electrode
{
Zn electrode
{
{
1 M ZnCl2
Pt / Pt-sponge electrode
electrochemical element
Figure 1.3: An electrochemical element consisting of a hydrogen standard electrode
(H/H+ ) (to which a standard potential of 0 V has been arbitrarily assigned) and a zinc electrode (Zn/Zn2+ ). Each element is in contact with
a solution of its ion which has a concentration of 1 M. The bridge is filled
with an inert solution (usually KCl or KNO3 ), solidified with agar. The
standard potential of Zn is −760 mV.
14
Reactions which involve the uptake of an electron are called reductions, those that
involve the removal of an electron are called oxidation. Both have to occur at the
same time, thus we speak of redox-reactions. This is one of two major classes of
reactions in chemistry, the other being acid-base reactions.
In our example, H+ is the acceptor and Zn the donor of electrons. If instead of
Zn we were to use for example a copper electrode, Cu2+ would be the acceptor and
H the donor. Thus we see that different elements have a different desire to donate
electrons: Zn > H > Cu. On the other hand Cu2+ has a higher desire to accept
electrons than H+ , we say that Cu is nobler than H. Metals which are nobler than
hydrogen are called noble metals (for example Cu, Hg, Au, Pt). Their electrochemical
potential, measured against a standard hydrogen electrode, is positive. Metals like
Zn, which can donate electrons to the H+ -ion, have negative standard potential.
Standard electrode-potentials can be measured not only for metals. Halogens for
example have a very high tendency to accept an electron, for example E(Cl2 , Cl− )
= 1360 mV.
1.4.1 Thermodynamics of electrochemical cells
The electromotive force (emf ) between two electrodes can be calculated as
the difference between the standard potential of the electron acceptor (say, Cu
+340 mV) and that of the electron donor (say, Zn −760 mV). Thus an emf of
340 mV − −760 mV = 1100 mV could be measured between Zn- and Cu-electrodes
under standard conditions (i.e. Concentration = 1 M, temperature = 25 ◦C).
The emf (E) is the pressure with which electrons are pressed through a circuit.
The work done in the circuit is the product of the emf and the amount of electrons
actually flowing, that is the charge Q:
w = −EQ
(1.28)
If one mol of ions are converted, the charge can be calculated from the number of
charges per ion (z), the number of ions per mol (NA = 6.022 × 1023 mol−1 ) and the
charge of the electrons (e = −1.602 × 10−19 C). The product NA * e is the charge
of one mol of electrons, this is called 1 Faraday (1 F = 96 480 C/mol). Thus:
w = −zF E
(1.29)
If the element is operated under fully reversible conditions (i.e. current is infinitely
low), w is equal to ∆G. As we have seen previously,
[C][D]
0
(1.30)
∆G = ∆G + RT ∗ ln
[A][B]
thus
RT
∗ ln
E =E −
zF
0
[C][D]
[A][B]
(1.31)
15
This is called the Nernst-equation and describes how the voltage across a cell
changes as we change the concentration of the ions in the solutions. If the reaction
is at equilibrium, E = 0 mV and
E0 =
RT
∗ ln(Keq )
zF
(1.32)
This allows Keq to be determined experimentally. Also, other important thermodynamical parameters can be determined in electrochemical experiments (provided
we carry out our experiments in such a way as not to change the concentration of
reactants):
∆S =
∂E
∂∆G
=
zF
∂T
∂T
∆H = ∆G + T ∆S = −zF E +
(1.33)
∂E
zF T
∂T
(1.34)
1.4.2 General Redox-reactions
So far we have looked at cells which are formed by an element and its salt (e.g.
Cu/Cu2+ ). However, a redox-system like Fe2+ /Fe3+ can also be used, if an inert electrode (for example platinum) is inserted into the solution. For example the reaction
Fe3+ + e− *
) Fe2+ can occur in an electrical electrode, the potential of this electrode
(with respect to the standard hydrogen electrode) depends on the ratio of Fe2+ /Fe3+
according to the Nernst-equation:
2+ [Fe ]
RT
0
∗ ln
E =E −
(1.35)
F
[Fe3+ ]
E 0 for this reaction is 771 mV.
1.4.3 Biological applications
Since we obtain the energy we need for our daily life from the oxidation of food,
redox-reactions are of considerable importance in biochemistry.
In mitochondria (and also in chloroplasts) energy is converted into a proton gradient across the membrane. Thus on one side of the membrane there is an excess of
positive charges from the protons, on the other an excess of negative charges from
the counter-ions. This tiny battery drives an electric motor (consisting of several protein molecules), which produces ATP from ADP and Pi . In other words: Electrical
energy is converted into mechanical, and mechanical into chemical energy.
This process is reversible, other enzymes can create a potential difference across
the plasma membrane. For humans the most important example is the Na/K-ATPase
which pumps 3 Na+ -ions out of and 2 K+ -ions into the cell for each molecule of
ATP hydrolysed. Thus the interior of the cell becomes negative with respect to the
outside (≈ 70 mV). When these ions are allowed to move freely across the membrane
16
O
H3C
N
H3C
N
N
N
O
H3C
N
H3C
N
N
R
H
R
FAD
H
H
+
N
N
O
FADH2
O
C
H
O
H
2 H+, 2 e-
2 H+, 2 e-
NH2
H
H
H
O
C
NH2
+ H+
H
R
NAD+
H
N
H
R
NADH + H+
Figure 1.4: NAD+ and FAD can accept activated hydrogen in one reaction and transport it to another. You will encounter them very often.
of a nerve cell to their equilibrium concentrations, the electrical potential collapses,
resulting in a nerve impulse.
By the way, 70 mV are not as meager as it may sound: The membrane is only
3.5 × 10−7 cm thick, so the field strength is about 200 kV/cm!
Redox-reaction in the cell often involve hydrogen. Cells therefore have developed
special carrier molecules to transfer activated hydrogen:
N AD+ + 2H + + 2e− *
) N ADH + H +
F AD + 2H + + 2e− *
) F ADH2
E = −320mV
E = −219mV
(1.36)
(1.37)
Standard potentials for free redox-reagents like FAD/FADH2 or Fe3+ /Fe2+ can be
quite different from those bound to proteins, this gives living cells the opportunity to
adapt a redox-reagent to the specific needs of a reaction. For example free Fe3+ /Fe2+
has a standard potential of 771 mV, for protein bound iron the potential can reach
from 365 mV (cytochrome f) down to −432 mV (ferredoxin).
17
1.4.4 Technical applications
Batteries
Batteries are packs of electrolytic cells. They are used to power electrical apparatus
independently from the grid. We distinguish primary batteries which can be used
only once (like in the definition of torch: container for dead batteries) from secondary
batteries, which can be recharged.
Primary cells Zinc/carbon and alkali/manganese cells are commonly used as batteries. They consist of a Zn anode, an electrolyte (sulphuric acid or KOH) and a
(inert) carbon electrode covered with MnO2 :
Zn → Zn2+ + 2 e−
MnO2 + 2 H+ + 2 e− → MnO + H2 O
This cell gives about 1.5 V. Alkaline cells have a higher capacity than those with
acidic electrolyte. Their chief disadvantage was until very recently that the MnO2
had to be mixed with ≈ 0.8 % mercury to prevent self-discharge. This meant that
used alkaline batteries had to be disposed of as toxic waste. Modern alkalines no
longer contain Hg (“green alkaline”), however, the laws in most developed countries
require batteries to be collected and recycled.
For long life applications with occasional high power demands (e.g. photography)
lithium batteries are used. They use lithium metal as anode and an inert carbon
electrode as cathode. Because Li is has a very negative standard potential and a
very low density, these cells offer an excellent power density. The cells are filled with
LiAlCl4 (as electrolyte) dissolved in thionylchloride as oxidant:
Li → Li+ + e−
4 SOCl2 + 4 e− → 4 Cl− + SO2 + S
Each cell gives about 3 V.
Secondary cells Lead/acid batteries are used in cars, lead/gel batteries (where
the acid has been converted into a gel for increased safety) in miners lamps. They
consist of a lead anode and a PbIV O2 electrode as cathode. Both are converted to
PbII SO4 during discharge :
−
* II
Pb0 + SO2−
4 ) Pb SO4 ↓ +2e
2− *
II
−
+
Pb O2 + 6 e + 4 H + SO4 ) Pb SO4 ↓ +2 H2 O
IV
Each cell gives 2 V.
Ni/Cd batteries use NiIV OH4 (or NiIII OOH) as cathode and Cd as anode, the electrolyte is KOH:
Cd0 + 2 OH− *
) CdII (OH)2 + 2 e−
− *
Ni (OH)4 + 2 e ) NiII (OH)2 + 2 OH−
IV
18
NiMH (Nickel metal hydride) batteries work similarly, but use a metal alloy as
anode, which can dissolve a large amount of hydrogen, thus the anode reaction
becomes:
H + OH− *
) H2 O + e−
The cathode reaction is the same as in NiCd-batteries. Both give about 1.2 V per
cell. NiMH-batteries do not use the toxic Cd and do not suffer from the “memory
effect” seen in NiCd batteries: When NiCd batteries are repeatedly only partially
discharged, Cd micro-crystals on the anode fuse into larger crystals, which do not
participate in further reaction cycles. This results in a loss of capacity. On the other
hand, NiCd batteries have a lower internal resistance.
Recently, rechargeable alkali-manganese cells have come on the market.
The disadvantage of all these cells is their high self-discharge rate (10–15 % per
week).
Electrolysis
Plating Electrolytic cells are essentially batteries run in reverse. They are used to
cover parts with a thin layer of noble metals (like gold). This process was invented
by the Parthians some 2000 years ago. They used a Fe/Cu element with vinegar as
electrolyte as battery to gold-plate items of silver.
Sometimes items are covered with metals that have a thin, impermeable oxide
layer on their surface (like chromium). This oxide layer prevents further corrosion.
This is called passivation.
Purification of metals Electrolysis can also be used to purify elements like copper.
Bars of impure copper form the positive pole (anode) of the cell. The copper dissolves
(Cu → Cu2+ + 2e− ), the resulting ions move to the negative pole (cathode), where
they form copper again (Cu2+ + 2e− → Cu). Metals that are less noble than copper
will dissolve, but will not be precipitated again at the negative pole and stay in
solution. Elements that are nobler than Cu (Ag, Au, Bi, Se, Te) will not dissolve, but
form a fine sediment below the anode from which they can be purified.
Polarography If the voltage between two electrodes is increased slowly, precipitation of elements will occur at their respective electrochemical potential. This reaction
is accompanied by current flowing, which can be detected and proves the presence
of a particular element. The charge required to completely precipitate that element
from a solution is proportional to its concentration. This can be used for qualitative
and quantitative analysis.
19
1.5 Water, acids, bases, buffers and pH
1.5.1 The auto dissociation of water
Water is essential for all life on earth, our own bodies are about 70 % water. As a solvent, water has some unusual physicochemical properties. One of them is that water
molecules can act both as acid (= proton donors) and base (= proton acceptors1 ).
This allows the following reaction to occur:
2H2 O *
) H3 O+ + OH−
H3 O+ will combine with further water molecules to form aggregates, the size of
which depend on temperature and other conditions. It is therefore easier (if slightly
incorrect) to write:
H2 O *
) H+ + OH−
Like for all chemical reactions we can define an equilibrium constant for this one:
Kd =
[H+ ] × [OH− ]
[H2 O]
(1.38)
One litre of an aqueous solution will always contain ≈ 1 kg of water (actually slightly
less depending on the concentration of solutes and temperature, but we will ignore
this imprecision for the moment). Because water has a molecular weight of 18.0 Da,
the concentration of water in an aqueous solution is:
1000 g/l
= 55.5 mol/l = 55.5 M
18.0 g/mol
(1.39)
Because the water concentration is constant, we can replace it in eqn. 1.38 and get:
Kd =
[H+ ] × [OH− ]
55.5 M
⇔
Kd × 55.5 M = [H+ ] × [OH− ] ≡ Kw
(1.40)
Thus the product of the proton and hydroxide ion concentration in aqueous solutions
is constant and called ion product of water or Kw for short. At 25 ◦C Kw =
1 × 10−14 M2 .
This has an important consequence: Whenever we know either the concentration of
protons or the concentration of hydroxide ions in any aqueous solution we can easily
calculate the other. It is therefore customary to give only the proton concentration.
In pure water the concentration of protons and hydroxide ions must of course be the
same, and we can write:
p
p
[H+ ] = [OH− ] = Kw = 10−14 M2 = 10−7 M
(1.41)
1
There are more general definitions of acids and bases, but for our purpose this one will do.
20
1.5.2 The pH of a solution
It is somewhat inconvenient to say something like 10−7 M. We therefore define:
pH ≡ − log([H+ ])
pOH ≡ − log([OH− ])
pK w ≡ − log(Kw )
(1.42)
(1.43)
(1.44)
The p is an abbreviation for the Latin pondus = tension. Because we can multiply
figures by adding their logarithms, the following two equations are equivalent:
Kw = [H+ ] × [OH− ]
⇔
pK w = pH + pOH
(1.45)
Thus we remember the following:
• A neutral solution has equal concentration of protons and hydroxide ions,
namely 1 × 10−7 M each at 25 ◦C, and therefore a pH of 7. Because the ion
product of water is temperature dependent, the pH of a neutral solution will
also change with temperature (7.5 at 0 ◦C, 6.5 at 60 ◦C).
• In acidic solutions the proton concentration is higher than 1 × 10−7 M, and the
pH is less than 7. To maintain Kw constant, the hydroxide ion concentration
must be smaller than 1 × 10−7 M, and the pOH therefore larger than 7.
• In basic solutions it is the other way round, [H+ ] is less than 1 × 10−7 M, the
pH therefore larger than 7.
• Because we can calculate pOH from pH, it is sufficient to give the pH of a
solution, pOH is redundant.
This is of great physiological significance, as even small deviations of blood pH from
the normal value of 7.4 will cause coma and death in animals and humans (acidosis
and alkalosis, respectively). We will see in the next sections how the pH is normally
maintained at the correct value in our bodies.
If we dissolve HCl gas in water, the following reaction will occur:
HCl *
) H+ + Cl−
Dissociation into H+ and Cl− will be almost complete. Thus, in 0.1 M HCl [H+ ] will
be 1 × 10−1 M and the pH will be 1. pOH will be pKw - pH = 13. On the other
hand, if we dissolve 0.01 M NaOH in water, [OH− ] will be 1 × 10−2 M, corresponding
to pOH = 2. Thus pH will be 14 - 2 = 12.
A famous trick question in chemistry exams is the following: What is the pH of
a 1 × 10−8 M solution of HCl? You should work this out first, and then look at the
footnote2 for the solution. Remember: This is a trick question, so think carefully!
2
Dissolving 1 × 10−8 M HCl in water will add 1 × 10−8 M H+ to the 1 × 10−7 M already present
from the auto-dissociation of water. Thus [H+ ] will be 1.1 × 10−7 M corresponding to a pH of
6.96 (and not 8, because we have added an acid, so pH must drop).
21
1.5.3 Strong and weak acids and bases
Acids and bases which dissociate almost completely into their ions in aqueous solution are called strong.
Note that we have defined acids as proton donors and bases as proton acceptors.
Thus HCl is an acid, because it can donate a proton. By the same token, Cl− is a
base, because it can accept a proton. We say, Cl− is the corresponding base to
HCl (or HCl is the corresponding acid to Cl− ). Lets take another example of this:
−
NH3 + H2 O *
) NH+
4 + OH
Clearly, NH3 is a base, because it can accept a proton, and NH+
4 is the corresponding
acid, because it can donate one. We also see that water acts as a proton donor (acid)
in this reaction, and that OH− is the corresponding base.
The unusual thing about water is that it can act both as an acid and a base,
because the dissociation of HCl should actually be written as:
HCl + H2 O *
) H3 O+ + Cl−
In this case water acts as proton acceptor (base), its corresponding acid is H3 O+ .
Whether water acts as acid or base therefore depends on circumstances: In the
presence of a strong acid water will behave as a base, while in the presence of a
strong base it will behave as an acid.
So far we have used the terms “strong” and “weak” in a purely qualitative sense.
It is however possible to quantify this. Let’s look at the general equation for the
dissociation of an acid:
HA *
) H+ + A−
This leads to the following equilibrium constant:
[H+ ] × [A− ]
[HA]
pK a = − log(Ka )
Ka =
(1.46)
(1.47)
Clearly, the stronger an acid is, the lower will be [HA] and the higher will be the
concentrations of both H+ and A− . Thus for a strong acid Ka will be large and pKa
will be small. In the same way we can define for a base:
B + H+ *
) HB+
[HB+ ]
Kb = +
[H ] × [B]
pK b = − log(Kb )
(1.48)
(1.49)
(1.50)
Again, the stronger a base is, the more the equilibrium will be shifted to the products,
leading to a high Kb and a low pKb .
22
Acetic acid
7
titration curve
pH
6
5
4
3
0
0.25
0.5
OH-Equivalents (mol/mol)
0.75
1
Figure 1.5: Neutralisation of acetic acid (a week acid) with NaOH (a strong base).
For each corresponding acid/base pair the following equations hold:
Ka ∗ Kb = 10−14
⇔
pK a + pK b = 14
(1.51)
Like with pH and pOH it is therefore sufficient to know either pKa or pKb and
usually you will find pKa listed.
Remember: pKa is inversely proportional to the strength of an acid, the stronger
the acid the smaller pKa .
How can we determine pKa experimentally? Lets have a look at fig. 1.5. It shows
how the pH changes, when a sample of a weak acid like acetic acid (pKa = 4.76) is
neutralised with a strong base like NaOH. We know that at each point of the curve,
both of the following equations must be true:
[H+ ] ∗ [OH− ] = Kw
[H+ ] × [A− ]
= Ka
[HAc]
(1.52)
(1.53)
As more and more NaOH is added, more and more protons must combine with
hydroxide ions to form water, in order to maintain Kw constant. To maintain Ka
23
Acetic acid
7
titration curve
pH
6
5
4
3
0
0.25
0.5
0.75
1
Figure 1.6: Addition of a certain amount of NaOH to acetic acid both at the beginning and the end of the experiment leads to considerable changes in pH.
The same amount of base added in the middle of the curve, where pH
≈ pKa leads to a much smaller change in pH.
constant, acetic acid molecules then need to dissociate to acetate and protons, until
essentially all the acetic acid is converted into ions.
Let us now look what happens if we neutralise exactly half of the acid present.
Under these conditions [HA] = [A− ] = 0.5 mol/mol and the second equation simplifies
to:
+
+
[H ] ∗ 0.5 mol/mol
[H ] × [A− ]
= − log
= − log([H+ ]) = pH
pK a = − log
[HA]
0.5 mol/mol
(1.54)
Thus if we neutralise exactly half of a given amount of acid (or if we mix equal
amounts of an acid and its corresponding base) the pH of the solution will give us
the pKa of the acid.
1.5.4 Buffers
As mentioned before, the pH of blood (and cytosol) needs to be carefully controlled
to prevent death. With the knowledge of the last paragraphs we can now try to
24
understand how this is achieved.
If you look at fig. 1.6, you will see that adding a certain amount of base at the
beginning of the neutralisation curve will lead to a large increase in pH. The same
is true at the end of the curve, when most of the acid is neutralised. However, in
the middle of the curve, where pH ≈ pKa , the change in pH caused by the addition
of the same amount of base is much smaller. A mixture of a weak acid (or a weak
base) and its salt is able to resist changes in pH caused by the addition of acids
or bases. We say: The mixture acts as a buffer. Addition or removal of protons
in this situation will change the HA/A− (or B/HB+ ) equilibrium, rather than the
concentration of protons.
Starting with the definition of Ka we can derive:
[H+ ] × [A− ]
[HA]
[HA]
[H+ ] = Ka ∗ −
[A ]
[HA]
pH = pK a − log
[A− ]
Ka =
(1.55)
(1.56)
(1.57)
The last equation is known as the Henderson-Hasselbalch buffer equation. It
allows to calculate the required buffer composition for each pH desired.
We have seen that the pH changes in the solution upon addition (or removal)
of protons depend on how close the pH is to pKa . The ability of a buffer to resist
changes in pH is called buffering capacity. The buffering capacity is defined as the
amount of protons or hydroxide ions you have to add to a buffer in order to change
the pH by 1 unit and can be calculated from the buffer concentration (c) and the
pH by the following equation:
c × Ka × [H+ ]
Kw
d[B]
+
= 2.303
+ [H ] + +
(1.58)
dpH
Ka + [H+ ]
[H ]
If 3 < pH < 11, that is in most biological systems, the term in round brackets can
be ignored, simplifying the equation somewhat.
Several buffer systems maintain the pH of blood at the correct level:
2− *
3−
+
* +
H3 PO4 *
) H+ + H2 PO−
4 ) 2H + HPO4 ) 3H + PO4 (pKa = 2.14, 6.86 and
12.4)
− *
+
*
*
CO2 + H2 O ) H2 CO3 ) H + HCO3 ) 2H+ + CO2−
3 (pKa = 3.77 and 10.2)
As we will see later, proteins also have considerable buffer capacity.
If you have understood the preceding section, you should now ask a question: How
can the carbonic acid / bicarbonate / carbonate system act as buffer at pH 7.4, if
the pKa values are 3.77 and 10.2?
The answer is that this system is at equilibrium with a high concentration of CO2
in our tissues and lungs. If the pH drops, carbonate is converted to bicarbonate,
25
Acetic acid
0.6
7
pH
0.4
5
0.2
buffer capacity (mol / mol*pH)
6
4
titration curve
buffer capacity
3
0
0
0.25
0.5
0.75
1
Figure 1.7: The buffering capacity is highest near the pKa value and declines to
either side. In practice buffers are good to within ± 1 pH-unit of the
pKa .
26
and bicarbonate to carbonic acid, which is hydrolysed to CO2 . Thus the amount of
anions in the blood is reduced, which increases pH.
If on the other hand the pH is increased, bicarbonate is converted to carbonate. To
maintain equilibrium, carbonic acid is converted to bicarbonate and CO2 to carbonic
acid. This will tend to lower pH.
This is only one example for a general rule: The law of mass action is directly
applicable only for reactions in a single phase. If more than one phase is involved
(here: liquid and gas), things become more complicated.
1.5.5 Indicators
Assume we have an acid which has a different colour (say, red) than its corresponding
base (say, blue). Upon addition of a base the pH of a solution of this acid will change,
more and more of the acid will be converted to the corresponding base and the colour
of the solution will change too. If the pH is well below pKa the solution will be pure
red, and if pH is well above pKa it will be pure blue. But if pH is near the pKa the
solution will assume different purple shades depending on the relative concentrations
of the acid and its corresponding base. By comparing the colour of the solution with
the colour of solutions with known pH we can estimate the pH of our solution. Such
acids do indeed exist, they are called indicators.
Obviously we must add only a very small amount of our indicator acid, so small
that it does not noticeably change the pH of the solution.
Also we must select our indicator so that its pKa value is near the pH of the
solution we want to measure. This is of course inconvenient, because we do not
know this pH (otherwise we would not need to measure it). The solution for this
problem is to use a mixture of different indicators, with different pKa and different
colours. Say we use one indicator which changes from yellow to red at a low pH and
another one which changes from red to blue at a higher pH. Initially the solution will
be a mix of yellow and red (=orange), as the pH increases it will turn first red and
later purple (blue + red). Obviously one can use more than 2 acids. Such mixtures
are called universal indicators.
It is possible to soak paper with an indicator solution and dry it. If a piece of
such paper is then held into a sample, the colour change will indicate its pH. Such
papers are clinically used for example to measure the pH of urine samples. Today
pieces of such indicator paper are mounted onto a plastic stripe with other pieces
of paper which show the concentration of glucose, protein, ketone bodies and other
interesting parameters. Such CombiSticks are invaluable for the quick diagnosis of
diseases like diabetes.
27
2 Atom structure and chemical
bonds
2.1 Atoms and elements
In the 4th century BC the Greek philosopher Demokrit made a Gedankenexperiment (experiment in the mind rather than the test tube): What would happen if
one cut a piece of, say, copper in half, took one of the pieces, cut it in half and so
on. Would it be possible to continue to do so ad infinitum? Demokrit came to
the conclusion that this was not the case, that there must be a point were further
division is no longer possible. This smallest particle he called atom (from the Greek
word for un-dividable). The Greeks believed that there were only four possible sorts
for such atoms, called elements: Fire, earth, air and water.
The idea of smallest particles has survived into modern science, but today we
know 104 elements stable enough for their chemical properties to be investigated.
Although the atoms that make up these elements are the smallest structures handled
by chemists, nuclear physicists have learned to break them up into smaller units,
called elementary particles. For our purposes, only 3 of those are important: Protons, neutrons and electrons. The positively charged protons and the uncharged
neutrons together form the nucleus of an atom, which is surrounded by negatively
charged electrons.
Earlier models assumed that the electrons circled around the nucleus like planets around the sun. However, this model led to consequences which do not fit the
experimental observations. Schrödinger has introduced a mathematical concept,
called wave equation, that can make successful predictions on how atoms behave.
Note that chemical compounds are formed by interactions between the outer electrons of atoms, thus it is important to understand the behaviour of these electrons.
We will now study the basic properties of the Schrödinger-model in nonmathematical terms.
2.2 Basic quantum theory
For each positively charged proton in the nucleus of an atom there is a negatively
charged electron in orbit. Since there is an electrostatic attraction between nucleus
and electrons, energy is required to move an electron away from the nucleus. This
energy can be added only in packets, such a packet is called a quantum. As a result,
electrons can move only at certain distances from the nucleus, called orbitals.
29
Figure 2.1: The 4 elements of the Greek philosophers (Fire, soil, air and water) reflected two pairs of properties: hot ↔ cold and wet ↔ dry. Fire is hot
and dry, soil is cold and dry, water cold and wet and air hot and wet.
All other substances can be mixed from these 4 elements in varying proportions. These 4 elements correspond to 4 of the 5 Platonic bodies:
Tetrahedron, cube, octahedron and icosahedron. The 5th , the dodecahedron, represents the cosmos, as all the others can be included into it.
The alchemist symbols for the elements are shown on their respective
bodies.
Figure 2.2: Energy-levels of electron orbitals. Electrons fill the orbitals with the lowest possible energy first. Absorbtion of energy (UV- or visible light) can
push an electron to a higher orbital.
30
2.2 Basic quantum theory
2.2.1 The first row of the periodic system: H and He
Orbital 1 is closest to the nucleus, and can contain 2 electrons. Because it is impossible to measure both location and speed of an electron at the same time (Heisenberg
relationship), we can calculate only the probability of finding an electron at a specific
time in a specific place (this is achieved by the wave function mentioned above). If
we connect all the points that contain the electron with high probability (say, 90 %),
we obtain the shape of the orbital. For orbital 1, this is a sphere.
Hydrogen has only one electron, helium has two. Thus in helium the first orbital
is completely filled, this is a particular stable state. For this reason helium is an
unreactive “noble gas”. In hydrogen on the other hand the outer layer is only partially
filled. For this reason hydrogen is very reactive. It will react with another hydrogen
atom, forming a H2 molecule. In this molecule the two electrons are shared between
the atoms, resulting in a covalent bond.
Hydrogen can also accept an electron from another atom:
H + X → H− + X+
or donate one to another atom:
H + X → H+ + X−
H− is called the hydride ion and has a complete outer orbital (at the expense of
another atom, which donated an electron). H+ , the hydrogen cation or proton, we
have met already in the section on acids and bases. Here too the hydrogen has
obtained a full outer electron layer (after a fashion) by dumping the unpaired outer
electron onto another atom.
Thus atoms can obtain a complete outer electron orbital by dumping excess electrons onto other atoms, or by accepting missing electrons from other atoms. In either
case, they become electrically charged ions. This is called an ionic bond.
2.2.2 The second row of the periodic system:
Li, Be, B, C, N, O, F, Ne
The second electron layer is further away from the nucleus, has a larger diameter
and can accept more electrons than the first, namely 8.
However, all electrons in an atom need to be different. Thus each orbital can hold
only 2 electrons, and those need to have different spin (angular momentum). The
spin of an electron is indicated by an arrow, which can point either up or down (↑
or ↓).
Thus the second electron layer is divided into 4 orbitals. The one with the lowest
energy is the 2s-orbital, which as a spherical shape like the 1(s)-orbital discussed
above. Li has one, Be two electrons in the 2s-orbital.
Further electrons are added into the 2p-orbitals, of which there are three, called
2px , 2py and 2pz . p-Orbitals are shaped like two fused drops, they can be oriented
31
Figure 2.3: The three p-orbitals are oriented orthogonal to each other. Each has
the shape of two drops fused together. The nucleus is located at the
intersection of the orbitals. The spherical 1s- and 2s-orbitals are not
drawn.
either in the x-, y- or z-direction of a coordinate system that originates in the nucleus.
All three have the same energy, thus the p-orbitals are called degenerate.
According to Hund’s rule, the p-orbitals are filled first with one electron each, all
of which have the same spin (traditionally indicated by ↑). Boron has the configuration 1s↑↓, 2s↑↓, 2px ↑, Carbon 1s↑↓, 2s↑↓, 2px ↑, 2py ↑ and Nitrogen 1s↑↓, 2s↑↓,
2px ↑, 2py ↑, 2pz ↑.
Once all p-orbitals have one electron, further electrons are added to them with
opposite spin: Oxygen 1s↑↓, 2s↑↓, 2px ↑↓, 2py ↑, 2pz ↑, fluorine 1s↑↓, 2s↑↓, 2px ↑↓,
2py ↑↓, 2pz ↑ and neon 1s↑↓, 2s↑↓, 2px ↑↓, 2py ↑↓, 2pz ↑↓.
Neon is a noble gas and will not easily participate in chemical reactions, because
its outer electron layer is fully occupied, which is a stable configuration. Li, Be and
B have few electrons in their outer layer, they will tend to give them up and become
positively charged ions (cations). On the other hand N, O and F have already many
electrons, they will tend to accept electrons from other atoms and become negatively
charged ions (anions). Carbon is in between.
2.2.3 Electronegativity
The affinity of an element for electrons is called its electronegativity. Elements with
only one electron in the outer layer (alkaline metals: Li, Na, K, Rb, Cs and Fr) will tend
to give them up (low electronegativity), as they would otherwise need to accept 7
electrons to achieve a full outer layer.
Elements with 7 electrons in the outer layer (halogens: F, Cl, Br, I, At) will tend to
accept one electron from another element, rather than giving up all seven of them.
Elements with a small number of electrons will have their outer layer close to the
nucleus, their electrons will be fully exposed to the electrostatic attraction of the
nucleus. Such elements will tend to accept electrons rather than give them up.
On the other hand, very heavy elements will have their outer electrons far away
32
2.3 Quantum theory II: Orbital hybridisation
Figure 2.4: The shape of a single sp-hybrid orbital. The nucleus would be located at
the node.
from the nucleus, additionally the positive charge of the nucleus will be partially
screened by the inner electrons.
Thus we can explain why fluorine (a light halogen) has the highest electronegativity of all elements (by definition 4.0). On the other hand heavy alkaline metals
like Cs and Fr have the lowest electronegativity (0.7).
Our model so far explains ionic bonding very well. However, further refinement is
needed for covalent bonds. Take the example of beryllium. The electronic configuration is 1s↑↓, 2s↑↓. All three p-orbitals are empty. Thus there are no orbitals with one
electron, which could form a covalent bond. However, Be is known to form covalent
bonds, for example H3 C−Be−CH3 .
Let us assume an electron could jump from the 2s-orbital to one of the 2p-orbitals,
say, 2px . The energy difference between 2s- and 2p-orbitals is not very big, thus such
a jump is feasible. The electronic state would be 1s↑↓, 2s↑, 2px ↑, in this configuration two orbitals with single electrons would be available for bonding. However, the
bonds between the Be-atom and the two methyl groups would have to be of unequal
length, because one would be formed by a s- the other by a p-orbital. Unfortunately
experiment tells us that the bonds in (CH3 )2 Be are of equal length and form an angle
of 180◦ .
2.3.1 The sp-hybrid orbital
To get us out of this mess, we have to assume that the s- and p-orbitals can hybridise
to form 2 sp-hybrid orbitals. These look a bit like p-orbitals, except that the two
“drops” are unsymmetrical, one being much smaller than the other.
Because both orbitals contain negatively charged electrons, the repel each other
and try to stay at maximum distance from each other. This is the case when the
angle between them is 180◦ , thus our theory agrees with the experimental result.
Note that Be in (CH3 )2 Be does not reach a full outer electron layer. This compound
therefore is difficult to prepare and highly unstable.
33
2.3.2 The sp2 -hybrid orbital
Boron, the next element in the periodic system, has the configuration 1s↑↓, 2s↑↓,
2px ↑. Thus there is one orbital with a single electron available for bonding. However,
experiment tells us that boron forms three bonds, for example in BH3 (boran). Again
these bonds are of equal length, BH3 is a flat molecule, the angles between the bonds
is 120◦ .
Again we resort to our old trick and lift one electron from 2s to 2py . Then we
hybridise the 2s orbital with both occupied p-orbitals, forming 3 sp2 -orbitals, each
with a single electron. As these orbitals repel each other, the form angles of 120◦ ,
explaining the experimental observation.
Again the outer electron layer of B in boran is incomplete, and the compound
fairly reactive.
34
Figure 2.5: The four sp3 -hybrid orbitals of carbon form a regular tetrahedron (indicated in red).
The sp3 -hybrid orbital
You should by now be able to guess what happens with the next element, carbon.
Carbon in the ground state has the configuration 1s↑↓, 2s↑↓, 2px ↑, 2py ↑. Thus two
orbitals are available for bonding, but carbon forms compounds like CH4 (methane),
with all bonds of equal length and the hydrogen atoms in the corners of a regular
tetrahedron.
So again we lift one of the 2s-electrons into the 2pz -orbital, then we hybridise the
2s-orbital with all three p-orbitals to get 4 sp3 -hybrid orbitals. The biggest possible
distance between these orbitals is achieved when they point to the corners of a
regular tetrahedron, with bond angles of 109.5◦ .
Carbon in methane achieves a full outer electron layer with 8 electrons. Methane
is a fairly stable compound, although it will react for example with oxygen.
2.3.3 Nitrogen and oxygen: Hybrid orbitals with 2 electrons
The next element in the periodic system is nitrogen (1s↑↓, 2s↑↓, 2px ↑, 2py ↑,
2py ↑). Nitrogen too forms sp3 -hybrid orbitals, but one of them has 2 electrons
and can not form covalent bonds with other atoms. Thus nitrogen can bind three
other atoms covalently, for example in ammonia, NH3 . The fourth orbital, with its
two electrons, is somewhat “fatter” than the other three, and requires more space.
Thus the three bond-forming orbitals are closer together than in methane, with a
bond angle of 107.3◦ rather than 109.5◦ . Additionally, the great density of negative
charges attracts positive ions, for example H+ , to form NH+
4 . Since in this bond all
two electrons are donated by the nitrogen, we call it a dative bond.
Oxygen (1s↑↓, 2s↑↓, 2px ↑↓, 2py ↑, 2py ↑) after sp3 -hybridisation has two hybridorbitals with two electrons, the remaining 2 orbitals have one electron each. This
means that oxygen can form two covalent bonds with other atoms, for example H2 O.
Additionally it can form dative bonds with positively charged ions, like in H3 O+ .
Because of the considerable space occupied by the two orbitals with two electrons
the bond angle is even smaller than in nitrogen, 104.5◦ .
35
H
H
H
C C
H
Ethylene
Figure 2.6: sp2 -hybridised carbon. The three hybrid orbitals (copper, brass and gold)
are in a single plane, and form angles of 120◦ . The remaining p-orbital
(silver) is orthogonal to this plane. This is the electronic configuration
of the carbon atoms in ethylene.
2.3.4 σ and π-Bonds
As we have seen, hybridisation of the 2s orbital of carbon with the three p-orbitals
(each of them occupied by one electron after one of the 2s-electrons moved to 2py )
results in the formation of four sp3 -hybrid orbitals. These four orbitals can react
with 4 other atoms, forming σ-bonds. Carbon is unusual in that it can form long,
stable chains with itself, which is one of the reasons that there are so many different
organic (i.e., carbon-based) compounds.
However, hybridisation can, but does not need to involve all three p-orbitals. It
is also possible to hybridise only one or two p-orbitals with the s-orbital, resulting
in sp- or sp2 -hybrids. The remaining unhybridised p-orbitals can also form covalent
bonds, but only with other unhybridised p-orbitals. These are called π-bonds.
As we have seen with boron, the three sp2 -hybrid orbitals lie in a single plane, and
form angles of 120◦ . In carbon they look the same. However, the remaining p-orbital
is orthogonal (forms an angle of 90◦ ) with this plane. If another sp2 -hybridised
atom approaches, a σ-bond is formed between their hybrid orbitals. Additionally,
the lobes of the p-orbitals lean toward each other, the resulting bond looks like a
sausage. Because each p-orbital has two lobes, a π-bond consists of two sausages. A
double bond is slightly shorter than a single bond, and considerably more reactive.
As in beryllium, the two hybrid-orbitals of a sp-hybridised carbon lie on a straight
line (180◦ bond angle). There are two unhybridised p-orbitals left, which are orthogonal to each other and the hybrid-orbitals. Triple bond formation between
two sp-hybridised atoms proceeds analogous to double bond formation, but involves two p-orbitals. As each p-orbital has two lobes, a triple bond contains four
“sausages”. Triple-bonds are even shorter and more reactive than double bonds,
acetylene (H − C≡C − H) for example can explode upon gentle warming or when
compressed to moderately high pressure.
36
H C C H
Acetylene
Figure 2.7: sp-hybridised carbon. Two sp-orbitals (brass and copper) form σ-bonds,
the two remaining p-orbitals (silver and gold) form π-bonds like in
acetylene.
37
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
H
Benzene
Figure 2.8: The π-electrons in benzene are delocalised, the two structures with double bonds are merely limiting cases, the “true” nature of benzene is in
between. The delocalised electrons are most likely found in two rings, in
front of and behind the carbon ring.
2.3.5 Delocalised π-bonds: Mesomery
A very important compound for organic chemists is benzene (C6 H6 ). For a long time
chemists found it difficult to understand its structure, until Kekulé solved the
riddle. He proposed that benzene is a cyclic compound, where the 6 carbon atoms
form a ring. Every second bond in this ring would be a double bond. Some problems
however remained: Alternating single-and double bonds would have different lengths,
while all C−C bonds in benzene have the same length. Also, as mentioned above,
double bonds are reactive, while benzene is fairly inert.
These problems were solved by the assumption that the π-electrons are not localised in double bonds, but form two rings above and below the carbon ring of
benzene. This delocalisation (mesomery) obviously increases the entropy of the
system, explaining the greater than expected stability of the benzene molecule. Also,
because all C−C bonds would be identical (somewhere in between double- and single
bonds), the geometry of benzene is explained by this model.
Delocalisation can occur whenever double-bonds alternate with single bonds, such
systems are called conjugated double bonds. But delocalisation is not limited to
conjugated double bonds, carboxylates for example are also mesomery-stabilised:
O
H3C C
O
O
H3C C O
O
H3C C
O
Acetate
The two oxygen atoms in a carboxylate are identical, both oxygens have part of the
negative charge, and are bonded to the carbon by a bond with a length in between
that of a single- and a double-bond.
38
2.4 Weak interactions
δ+
H
δ− O
H
δ+
H
O
H
O
H
H
Figure 2.9: Water molecules are polar, because the oxygen pulls the bond electrons
away from the hydrogen (indicated by wedge shaped bonds). The oxygen
atom has a partial negative, the hydrogen atoms partial positive charges.
This results in electrostatic attraction between water molecules, called
hydrogen bonds (dotted lines).
In this section we will discuss a couple of interactions, which lead to weak binding between molecules. “Weak” in this context means that the energy of binding
(5–20 kJ/mol) is similar to the average thermal energy of molecules at room temperature (2–3 kJ/mol). Thus these bonds are easily broken and reformed, unlike
covalent bonds with binding energies of 200–500 kJ/mol, which are much more stable. However, large biomolecules like proteins and DNA are stabilised by thousands
or even millions of weak bonds, and their combined binding energy can be significant.
2.4.1 Hydrogen bonds
We have already seen that different elements have different affinities for electrons,
i.e. electronegativities. Oxygen for example has a high affinity for electrons, while
hydrogen has a low one. Thus the shared electrons of a bond between these elements
will be pulled more to the oxygen than to the hydrogen. As a result the oxygen will
obtain a negative partial charge, the hydrogen a positive one. We say the bond
between oxygen and hydrogen is polarised.
You can think of such a polarised bond as intermediate between an unpolarised
covalent bond between elements of similar electronegativity (C-H for example), and
an ionic bond (like Na+ Cl− ).
Opposite charges of course attract each other. Therefore there will be an interaction between the partially positive charge of a hydrogen in one water molecule
and the partially negative oxygen in another molecule, i.e. there will be a bond between these molecules. Such bonds between partial charges in polarised molecules
39
O O
δ−
δ+
O O
δ−
δ+
O O
O O
O O
δ− δ+
O O
Figure 2.10: van der Waals interaction between oxygen molecules. Top: Both
molecules in ground state. Middle: Random movement of electrons in
the left molecule results in partial electric charges. Bottom: Because of
electrostatic attraction charge separation occurs in the second oxygen
molecule, and a weak, fleeting bond is formed between the molecules.
are called hydrogen bonds.
The strong polarisation of the water molecule has an important consequence:
Water at room temperature is a liquid rather than a gas (like CH4 , NH3 or H2 S).
Hydrogen bonding holds the water molecules together and prevents them from evaporating.
2.4.2 van der Waals bonds
We have seen that we can not predict the position and momentum of electrons, their
movement is random. At any given time the electrons of an atom or molecule may
be distributed more or less evenly, and the molecule will appear neutral. However,
electrons may also by chance concentrate on one side of the molecule, giving this
side a partial negative charge. The opposite side of course will have a partial positive
charge, as it lacks electrons.
Assume a second molecule comes near the temporarily positive side of the first.
The partial positive charge in the first molecule will induce partial charge separation
in the second. Attraction between these molecules results in a temporary bond, the
van der Waals-bond.
Note that all molecules and atoms are, at least in principle, capable of forming
van der Waals-bonds, this force exists even between helium atoms.
van der Waals-bonds are quite weak, but still noticeable in every day life. For
example, if gases are compressed, they become hot. By compression the distance
between gas molecules is reduced, this increases the probability of van der Waalsbonds between them. The heat is simply the binding energy liberated by this process.
Expanding gases on the other hand become colder, as energy is required to break
the van der Waals-bonds. Refrigerators make use of this effect (Linde-process).
40
O
O C
O C
C H
C H
C
H
C
H
H
C
H
H
C
H
H C
H
H
C
H
H
C
H
H
C
H
H
C
H
H
C
H
H
C
H
C H
H
H
C
H
H
H
H
C
H
C H
HH C
C
H H
C
H
H
H
H
H
H
H
H
O
C
O
H
C
H
H C
C H
H
H C C H
H
H
H C C H
H
H
H C C H
H
H
H C C H
H
O
H
H
H H
H C C
C C
H
HH
C C
H
O
C
H
H
C
H C
H
HH
H H
C C
H
HH
C C
HH
H
H
H
C C
H
H
HH
C C
H H
H H
C C
HH
C C
H H
H H
H H H H H H H H H
H H H H H H H
H
O
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
C
H
O H H H H H HH H H H H H H H H H
H
H H H H H H H H HH H H H H H O
C
C
C
C
C
C
C
C
C
O
C
C
C
C
C
C
C
C
H
H
H
H H H H H H H H H H H H H H
H
H
C
Na+
Na-stearate
O
C H
C
H
O
H
C
O
H
C
H C
H
H
C H
H C
H
H
C
C H
H
C
C
O
H
H
H
H
H
C
C O
H
H
C H
H C
H
H
C
H
H
C C
H
H
H
C H
C
H
H
H
H C
C C H
H
H
C
H
H
H
C H
H
H C C
H
H
H C
C
H
H
H
H
H
C C HH
H
C H
H C H
H C
H
H
C H
H
C
C
H
H
H
H
H
C
C
H
H
C C HH
H
H C
H
H
C H
O
H
C C HH
H
H
H C H
H
H
H H
C O
H
H H
H
C C
H C C
HH
C C
H
H H
C C
H
H C
H
H H
C C
H H
H
C C
H
H H H H H
C C
H H
H
C C
H H
H
C
H
C
H H
C
H H
H H H H
H
H
H
H
C
C
C
H
H
H
H
H
H H
H H
C
H
C C
C C
HH
H
C C
HH
H
H
C C
H H
H H
H
H
H
C
H
C
C H
HH
C C
H
H H
H
H
H
C C
C H H H
HH
C C
H
H C
H
O C
H H
H C H
H C C H
H H
C
H
O
H
H
H
H C C H
H C
H
H
H
H C
H C C H
H
C
H
H
H
H
H C C H
C
H
H
H
H
C
H
H C C H
H
C
H
H
H
H
C H
C
H C
H
H
H
H
C
C H
H
H
C
C H
H
H
C
C
O
H
H
O
C
H
H
C
H
H
C
H
H C
H
H C
H H
C
H
H
H
H H H H H H H H HH H H H H H O
C
C
C
C
C
C
C
C
C
C
C
O
C
C
C
C
C
C
H H H H H H H H H H H H H H H H
H
C
H
H H H
H H H
H H H
C
H H H
C
C
C
H H H
C
C
C
O
C
HH H
C
C
C
H H H
C
C
C
H H H
C
H
H
C
H H H
H
O
H
H
H H
H C C
H
C
H
H
C C
C H
H H
H
H
H
H
H
H H C C
H C
H
C H
C C
H
H
H
H
H
H C
C C H H
H
H
H
C H
C C
H
H
H
H
H
H
C C
H C
H
H
C H
H
O
C C
H
H
H
H
H C
H C C
C H
O
H
H
H
H C
C H
H
H
H C
C H
H
H
H C
C H
H
O
C
C
H
H
H
C
C
H
H H
O
C O
O
Figure 2.11: Left: An example of a soap molecule. Soaps are the sodium salts of fatty
acids, here stearic acid with 18 C-atoms. Right: The long hydrocarbon
tail is hydrophobic and can insert into greasy spots. The negatively
charged carboxylic acid group is hydrophilic and interacts with water.
During washing dirt droplets get coated with charged head groups.
The negative charges repel each other, so the droplets tend to break up
(bigger surface area) and disperse in water.
2.4.3 Hydrophobic interactions
Water molecules form hydrogen bonds with each other, or with other polarised
molecules, like ethanol or sugar. Such substances with polar bonds dissolve easily in
water, they are called hydrophilic (water loving).
However, other substances like hydrocarbons (petrol) do not contain polar bonds.
If they were dissolved in water, they could not participate in hydrogen bonding,
and would distort the local structure of the water around them. Additionally, such
molecules force water molecules into highly ordered “cages” around them, this decreases entropy.
Both effects would require energy, and as a consequence such substances do not
dissolve in water in appreciable concentrations. If you mix petrol with water, you get
an emulsion of petrol droplets in water, rather than a solution of petrol molecules.
Substances like hydrocarbons, which do not mix with water, are called hydrophobic (water fearing) or lipophilic (fat loving). If hydrophobic substances are brought
into contact with water, they tend to stick to each other. This effect can be looked
upon as a weak chemical bond, we call it hydrophobic interaction.
Something interesting happens when we construct molecules which are hydrophobic at one end, and hydrophilic at the other (see fig. 2.11). Such molecules are called
amphipatic. A typical example of an amphipatic compound is soap. Soaps are the
sodium salts of long-chain fatty acids. The long hydrocarbon chain is hydrophobic,
41
Ground state
low field
high field
Figure 2.12: Complex formation in iron by energy splitting of d-orbitals.
the negatively charged carboxylic acid head group is hydrophilic. As a result the
tail can insert into a (greasy) dirt particle, the polar head group will stick out and
interact with water. Thus soap can help to remove dirt during washing.
We will see later that amphipatic compounds (phospholipids) are also essential
components of the cell membrane.
2.4.4 Transition metals and complex bonds
The third and fourth electron layers can accept more than 8 electrons each, these are
found in d- and f-orbitals. These orbitals are responsible for the complex reactions
of transition metals like Fe or Zn, which are of great biological significance.
Like p-orbitals, these are degenerate in the sense that all d-orbitals of a layer have
the same energy level (and all f-orbitals have too), see fig. 2.2. The Pauli-principle
and Hund’s rule apply, so iron for example has 5 d-electrons, each in a separate
orbital with spin ↑. If a ligand comes close, orbitals can split into higher and lower
energy levels (see fig. 2.12). Because of energy conservation the gain in the upper
level needs to be as big as the energy loss of the lower level. For some ligands the
difference in energy levels is small (low field ligands, for example [FeF6 ]3− ). Other
ligands (high field ligands, for example in [Fe(CN)6 ]3− ) induce a stronger splitting.
That energy difference is big enough to allow the electrons in the orbitals in the
upper level to drop to the lower ones, which of course requires reversal of spin. This
spin-reversal can be detected by electron-spin resonance spectroscopy (ESR),
an important technique to investigate the reaction mechanism of metallo-enzymes.
42
In many respects organic chemistry is much easier than general and inorganic chemistry are. Very little maths is involved, and you do not have to memorise complicated
formulas.
What you need to know is how to name a compound from a given structure, how
to write the structure when given the name, and how derive chemical properties
from the structures.
As the composition and structure of chemical compounds define their biological
behaviour, we need to be able to describe them unambiguously. For this purpose we
define the following terms:
Composition describes which atoms are present in a compound in what numbers.
Thus ethanol has the composition C2 H6 O.
Constitution describes the pattern of bonds in a molecule. For example, ethanol
(CH3 −CH2 −OH) and dimethyl ether (CH3 −O−CH3 ) have different constitution, but the same composition.
Conformation describes the position of the various atoms in a molecule, as far
as it can be obtained by rotation of chemical bonds. For example, ethane
(CH3 −CH3 ) can occur in eclipsed or staggered conformation. The energy difference between different conformations of a molecule is usually small, allowing
rapid, conversion between them. This conversion is the basis for IR- and Raman-spectroscopy. However, very bulky or charged residues may limit such
interconversions.
Configuration describes the position of substituents on chiral carbon atoms (carbon atoms with 4 different groups). Because interconversion of configuration
isomers would require the breaking and reformation of covalent bonds it can
not occur spontaneously.
3.1 Important functional groups
Several million organic chemicals have been described in the literature. During your
biochemistry training, you will come across several hundred of them, and pharmacology will add some of the 800 or so pharmaceuticals in current clinical use. It
would be very easy to get lost in this variety.
Fortunately, many of the properties of a compound, including its possible chemical reactions, can be gleaned from a short look at its structure, because they are
43
determined by the functional groups that a compound carries. Many students are
afraid of chemical structures. Do not be: They are your friends, not your
enemies. When you look at the structure of a compound, you should identify:
the carbon skeleton: This describes the way different carbon atoms are linked to
each other (chains, rings).
the functional groups: Everything else. Thus the simplest functional group is the
hydrogen atom −H.
We will now look at the most common functional groups which you will encounter
in your further studies.
3.1.1 Alkanes
Straight chains
The simplest organic compound is methane, CH4 . Carbon is unusual in that it
can form long, stable chains with itself. Thus we get ethane (H3 C−CH3 ), propane
(H3 C−CH2 −CH3 ) and so on. The first four elements of this series have trivial names
(methane, ethane, propane and butane). The names of all longer chains are derived
from the Greek words for the number of carbon atoms in the chain: pentane (5),
hexane (6), heptane (7) and so on. All alkanes end on -ane. Their general formula
is Cn H2n+2 .
Alkanes from methane to butane are gases at room temperature, pentane to undecane (C11 ) are liquids, higher alkanes are solid (paraffines).
All alkanes are relatively inert, except that they can be burned. Gaseous alkanes
mixed with oxygen in a proper ratio are dangerous explosives. Incomplete combustion leads to the production of solid carbon (soot) or, worse, highly toxic carbon
monoxide. Other than that relatively drastic conditions are needed to make alkanes
react (for example UV-irradiation).
Alkanes are non-polar, hydrophobic compounds, their hydrophobicity increases
with chain length. There is a medical application for this: A thin film of vaseline
(a semisolid mixture of alkanes) can protect the skin against water and other fluids.
This can be used to prevent diaper rash or rough hands.
Branched chain alkanes
The carbon chain in alkanes need not be straight lines, there can be branch points
on them. This does not change their chemical properties, but presents us with the
problem of naming them unambiguously. This is easy, but requires some practice.
Straight alkanes are often called n-alkanes (for normal), those with a branch ialkanes (for iso, from Greek ‘equal’).
The procedure is always the same: Identify the longest stretch of carbon atoms
in the molecule (see fig. 3.1). This gives you the name of the parent compound (in
this example octane). Then identify the substituents.
44
H3C
HC
CH3 CH3
CH2
H3C C C C C C CH2
H2 H2 H H2 H
3-Methyl-5-isopropyl-octane
Figure 3.1: Do not be misled by the way a structure is drawn. Identify the longest
carbon chain (here drawn in red), this determines the name of the compound (8 = octane).
In the case of fig. 3.1, there are other carbon chains linked to the octane. These
carry the same name as the corresponding alkanes, but the ending -yl. Thus we have
a methyl- and an iso-propyl-group.
To unambiguously name the compound we have to identify the positions where
the functional groups are attached. This we do by numbering the carbon-atoms of
the main chain. Thus in our example we have 3-methyl-5-isopropyl-octane.
You may now ask: Why not 4-isopropyl-6-methyl-octane? In such a situation we
add the substituent positions and choose the option with the lowest result: 3 + 5 <
4 + 6, thus the former option is correct.
Cyclic alkanes
Chains of 3 or more carbon atoms can be ‘made to bite into their own tail’, resulting
in cyclic molecules. They have the same name as their parent compound, but with
the prefix cyclo-. Cyclopropane and cyclobutane are unstable, because the carboncarbon-bonds need to assume angles very different from the tetrahedral angle of
109.5◦ (90◦ in cyclobutane and 120◦ in cyclopropane). This is known as Bayertension.
Because of the Bayer-tension 3- and 4-membered rings are reactive compounds.
Suitably designed members of this group may open in the catalytic centre of an
enzyme and chemically modify the protein molecule, leading to its inactivation.
The most well known example for this is the antibiotic penicillin which contains
a reactive β-lactam ring (see fig. 3.2). It inactivates an enzyme that is responsible
for the synthesis of the bacterial cell wall, thus the bacteria are killed during cell
division.
In structures these rings are often represented by geometric figures (triangle,
square, pentagon etc). Each corner represents a carbon atom, the hydrogens are
often omitted. Remember that each carbon has 4 bonds, so mentally add the required hydrogens.
45
O
H
N
O
H
S
N
CH3
CH3
COOPenicillin
Figure 3.2: Penicillin was the first antibiotic of biological origin. Due to Bayertension the reactive β-lactam ring opens in the catalytic centre of an
enzyme responsible for bacterial cell wall synthesis, killing them during
cell division.
Figure 3.3: C60 , Buckminsterfillerene, affectionately known as buckyball.
46
Figure 3.4: In a hexagonal conformation the carbon bond angle would be 60◦ . This
is unfavourable, thus cyclohexane assumes either the boat- or the chair
conformation, where bond angles are close to 109.5◦ . The energy difference between these two conformations is small, both are present.
Carbon-rings can fuse to form larger structures, the most impressive example is
Buckminsterfillerene (“buckyball”) with 60 carbon atoms arranged in 12 five-carbon
and 20 six-carbon rings. Thus the shape of this molecule resembles a football (and
here I mean the real football, which is kicked with a foot rather than carried around
by hand). The carbon atoms are all sp2 -hybridised, the double bonds tend to concentrate in the hexagons. When doped with other compounds, this molecule displays
a couple of technically useful properties like supraconductivity and ferromagnetism.
Cycloalkanes are not flat rings, but bend in 3 dimensions in order to have bond
angles as close as possible to 109.5◦ , and to prevent crowding of hydrogen atoms
(Pfitzer-tension).
Note that in cyclic compounds the bonds can no longer rotate freely, we will deal
with the consequences of this fact in the chapter on stereoisomery.
Double bonds
So far all carbons were linked by single bonds. However, carbon can form double
bonds (alkenes). Such compounds carry the name of the parent structure, but with
the ending -ene, (e.g. ethylene). When numbering carbon atoms in alkenes the numbers of the carbons carrying the double bond should be as low as possible.
Note that double bonds are not rotatable. This leads to cis-trans-isomery, as we
47
will discuss later.
Double bonds are energy rich and reactive. They will for example react with
halogens like Br2 or hydrogen halides like HCl. In the presence of a catalyst they will
also react with water or hydrogen. Such reactions, where a molecule reacts with a
double bond with no leftovers are called additions.
Alkenes can also react with themselves, forming long chains. Such reactions, where
a long molecule is formed without addition or removal of other molecules, are called
polymerisation. Polyethylene, polypropylene or polystyrene are important examples. Polymerisation is one of three types of reactions, which are used to make
plastic:
Polymerisation Molecules with C=C double bonds react with each other, forming
long chains. No small molecules are produced. We distinguish homo-polymers
(only one type of monomer) from copolymers (two or more different monomers)
like acryl butadien styrole copolymer (ABS).
Polycondensation Molecules with different functional groups react with each other,
producing small molecules as side products. Man-made examples are polyester
(see later), natural polycondensates are proteins and polysaccharides.
Polyaddition Like in polycondensation the side groups, rather than the carbon
chains react with each other, but like in polymerisation no small molecules
are produced as side-products. Polyurethanes are an important example for
polyadducts.
Conjugated double bonds Carbon chains with alternating double- and single
bonds are called conjugated systems. They show mesomery and are more stable
than their unconjugated counterparts (see the section on atomic structure for more
details on mesomery).
Triple bonds
Compounds with triple-bonds are called alkynes, they show reactions similar to
alkenes. Because each of the carbons forming such a triple bond can only form one
other bond, no cis-trans-isomers exist, even though the triple bond can not rotate.
3.1.2 Oxidation products
Oxidation in organic molecules means either the introduction of oxygen or the removal of hydrogen.
Alcohols
If oxygen is inserted into an alkane, we get an alcohol (-OH). Depending on whether
the carbon carrying the OH-group is connected to 1, 2 or 3 other carbon atoms, we
48
H2C CH2
Ethylene
H2O
Br2
HCl
Br
CH2
Br
CH2
Cl
H2
(Pt)
CH2
CH3
Dibromoethane
Chloroethane
(H2SO4 )
CH3
CH3
CH3
HO CH2
Ethane
Ethanol
Addition-reactions of ethylene
H2C
H2C
h*ν
H2C
H2C
catalyst
H2C
+
H2C
H2C
H2C
H2C
+
H2C
H2C
H2C
H2C
H2C
H2C
H2C
H2C
H2C
Polymerisation of ethylene
Figure 3.5: Top: Addition reactions to alkenes (here ethylene). Reaction with halogens and hydrogen halides are spontaneous, however, reactions with water or hydrogen require appropriate catalysts. Bottom: Polymerisation
of ethylene. In a reaction that requires energy, the electrons forming the
double bond separate, so that each C-atom has a single electron. Compounds with single electrons are called radicals and are highly reactive.
Such a single electron can attack the double bond of another ethylene
molecule, forming a new C−C-bond. This leaves a different carbon with
a single electron, so that the reaction continues. The reaction will terminate only if two radicals react with each other, forming a bond. Radical
reactions also occur in our body, and you will learn later how our body
defends itself against these very reactive, and therefore toxic, molecules
(“molecular terrorists”).
49
+ [O]
CH3
OH
O
- [H]
CH2
CH
CH2
CH2
CH2
CH3
CH3
CH3
Propane
Propan-1-ol
(n-Propanol)
CH3
+ [O]
CH2
Propanal
(Propionaldehyde)
- [H]
HC OH
CH3
CH3
Propane
CH3
CH3
CH2
CH3
Propionic acid
CH3
C O
CH3
Propan-2-ol
(i-Propanol)
+ [O]
HO O
C
+ [O]
Propanone
(acetone)
CH3
H3C C OH
H3C CH
CH3
CH3
2-Methylpropane
(iso-butane)
2-methylpropan-2-ol
(tert. butanol)
Figure 3.6: Oxidation products of alkanes.
50
speak of primary, secondary or tertiary alcohols. In IUPAC nomenclature, alcohols
have the name of the parent compound with the ending -ol (methanol, ethanol etc).
The OH-group is polar, thus alcohols have a higher solubility in water than the
corresponding alkanes. Indeed, methanol, ethanol and both propanols are freely
miscible with water. Because of hydrogen bonding, alcohols have a higher melting
and boiling point than their corresponding alkanes.
Because alcohols have a polarity intermediate between water and alkanes, they
make excellent solvents for many organic compounds. Methanol, ethanol and
propanol are also used as fuel for model planes or certain racing cars. Addition
of ethanol to petrol gives gasohol, a less environmentally unfriendly fuel for cars.
The most common alcohol is ethanol (ethylalcohol, “alcohol”). Apart from its industrial uses as solvent (for example in perfumes) and starting material for chemical
synthesis, it is component of many beverages (whine, beer, spirits). This is possible
because ethanol is much less toxic than other alcohols, however, excess consumption
(blood alcohol concentrations > 0.6 %) can lead to death. In medicine it is used as
solvent for tinctures and as a disinfectant. In the later function it is often replaced
by iso-propanol, which is cheaper. Although the consumption of small amounts of
ethanol has been shown to be beneficial to health, increased consumption leads to
dependency, liver damage and possibly death. Many countries control ethanol consumption by heavy taxes, others (often with islamic majority) prohibit it. Alcohol
consumption is associated with about 200 000 annual deaths in the USA (mainly
from alcohol related traffic accidents) and social costs of 120 × 109 US$ annually
for treatment, lost working time etc. At particular risk are unborn children if their
mother consumes even relatively minor quantities of alcohol during pregnancy (fetal
alcohol syndrome).
Another important alcohol is propane-1,2,3-triol, commonly known as glycerol.
Its high affinity to water and its low toxicity make it a valuable moisturising agent,
for example in hand creams.
Phenols Although the OH-group is polar, it is usually not acidic. Phenols are an
exception to this rule, because the vicinity of the π-system increases the acidity of
the OH-group to a pKa of about 10. Phenol (hydroxybenzene) was therefore called
carbolic acid in the older literature. It is a highly toxic compound with a typical smell
that was used as disinfectant during surgery at the end of the 19th century, when
the value of aseptic conditions became appreciated. It was introduced by Joseph
Lister, an English surgeon. Today it has been replaced by safer alternatives (4chloro-3,5-dimethyl phenol, o-phenyl phenol, chlorhexidine and others). The amino
acid tyrosine is a derivative of phenol.
Thiols and selenols In the catalytic centres of some enzymes weak acids perform
an important function. In cases where the phenolic hydroxy-group of tyrosine (pKa
= 10) is to weak, but carboxylic acid residues from glutamate and aspartate (pKa
≈ 4) are to strong, nature uses amino acids with thiol (−SH, pKa = 8.2, cysteine)
51
and selenol (−SeH, pKa = 5.2, selenocystein) groups. This will be dealt with further
in the section on enzyme mechanism.
Thiols are used as antioxidant: Glutathione in living organisms and βmercaptoethanol (HS−CH2 −CH2 −OH) in the laboratory. They react according to
R−SH + HS−R0 + [O] *
) R−S−S−R0 + H2 O. The resulting product is called a disulphide.
The name β-mercaptoethanol contains two outdated but still frequently used
components. Mercapto- means the same as -thiol. Greek letters are sometimes used
to describe the position of a functional group. The carbon with the name-giving
group (here −OH) is called α, the neighbouring β and so on. ω is sometimes used
for the last carbon in a chain, irrespective of the chain length.
Many thiols have strong odours. Derivatives of buthanethiol are the active ingredient in the secretions of skunks, 1-propene-3-thiol gives onions their characteristic
smell. Ethanethiol is added to natural gas so that it can be smelled.
Aldehydes and Ketones
If hydrogen is removed from a primary alcohol, an aldehyde is formed. The corresponding reaction in a secondary alcohol leads to a ketone. Note that it is not
possible to oxidise tertiary alcohols in a similar manner. Aldehydes have the ending
-al, ketones the ending -one. R−CO−R is called a carbonyl-group.
Carbonyl-groups are polarised. This results in lower melting and boiling points
than the corresponding alcohol, but higher than the corresponding alkanes. Because
the carbonyl-group is polarised, both aldehydes and ketones are water soluble (if the
carbon chain is not to long).
Aldehydes, but not ketones, can be further oxidised to carboxylic acid. This can
be demonstrated with Tollens’ reagent (AgNO3 in ammonium hydroxide) where
the Ag+ is reduced to elemental Ag0 , which precipitates as a metallic layer. This
reaction is used to produce mirrors. Benedict’s reagent and Fehling’s reagent
both contain soluble, blue Cu2+ -complexes, which are reduced to insoluble red Cu2 O.
These reactions are used in clinical chemistry to detect aldoses, that is reducing
sugars, for example in the urine of diabetic patients.
Addition of hydrogen in the presence of a nickel or platinum catalyst turns aldehydes and ketones into the corresponding alcohol.
Carboxylic acids
If oxygen is introduced into an aldehyde, a carboxylic acid residue is formed. Note
that ketones can not be oxidised in a similar manner.
Carboxylic acids are abundant in nature. Citric, lactic and acetic acid are present
in our food, formic acid causes the sting in ant bites.
Carboxylic acids are named after the parent compound, with the ending -oic acid.
Many of them have also trivial names (methanoic acid = formic acid, ethanoic acid
52
= acetic acid, butanoic acid = butyric acid). Fatty acids are carboxylic acids with
long (12–22) carbon tails.
Methanoic, ethanoic and propionic acid are liquids with sharp odour and highly
water soluble. Butanoic acid causes the smell of rancid butter and unwashed socks.
Carboxylic acids with higher molecular weight become increasingly insoluble in
water. Their ions form soaps, where the ionic head group interacts with water, the
hydrophobic tail with lipids, for example dirt. Thus soaps can disperse dirt in water
and help in cleaning.
Melting and boiling point of carboxylic acid is higher even than the corresponding
alcohols, reflecting the highly polarised nature of this group and its ability to form
hydrogen bonds.
3.1.3 Reaction products of alcohols, aldehydes and acids
Ethers
If two alcohol molecules react with each other after gentle warming in the presence
of a water binding catalyst (sulphuric acid), water and an ether may be formed, for
example:
H3 C−OH + HO−CH3 *
) H3 C−O−CH3 + H2 O
Because the ether-bond is not polar, their boiling points are about the same as that
of alkanes of the same molecular weight, and much lower than their parent alcohols.
Because of their low boiling points, low molecular weight ethers are highly flammable
and form explosive mixtures with air.
The most well known ether is diethylether (often just called ether), which was used
for general narcosis before it was replaced with safer, less flammable alternatives.
Working with ethers requires great care, because they may react with air oxygen
to form ether peroxides, which are explosive. For this reason, ethers must never be
allowed to evaporate to dryness.
Ether residues occur on many natural compounds. In this case the ending -oxy is
used to describe the functional group (methoxy-, ethoxy- etc).
Alkenes
Water removal from alcohols at higher temperature (> 180 ◦C) leads to the formation
of alkenes. These have been discussed already.
In biochemistry the elimination of water from an alcohol to form a double bond
(or the reverse reaction) is important and catalysed by a special class of enzymes,
the lyases.
Esters
Alcohols and acids react to form esters and water. The OH-group of water comes
from the acid, the remaining hydrogen from the alcohol.
53
H3C O
O
H
N CH3
O
dihydrocodeinone
O
O
H3C C O
H3C O
HO
O
H
Morphine
HO
3-methoxymorphine
(Codeine)
H
N CH3
N CH3
N CH3
HO
O
H
H3C C O
O
3,6-diacetylmorphine
(Heroin)
Figure 3.7: Morphine is an analgetic (pain-killing) drug with high addictive potential. 3-methoxymorphine (codeine) contains an additional methyl-group
bound via an ether-linkage. It is much less addictive than morphine, for
reasons you will learn in second semester. Diacetylmorphine (heroin) contains two acetyl groups in ester-linkage (see later). It is highly addictive
and therefore has no medical use. This is an example how small alterations in the structure of a molecule can change its biological behaviour.
54
Esters can be hydrolysed in the presence of bases (NaOH or KOH) to form acids
and alcohol. This process is called saponification and results in the Na- or K-salts of
fatty acids (i.e., soap).
Esters of low molecular weight acids and alcohols are volatile liquids, many of
them with characteristic, fruity odour. Some of them are used as artificial flavours
(ethyl butanoate = strawberry).
Fats are esters formed from 3 fatty acids and glycerol. They are used by plants
and animals to store energy.
Phospholipids are similar, but instead of the third fatty acid they contain phosphoric acid, which may be further linked to a polar headgroup. They are the main
constituents of the plasma membrane which surrounds the cells of our body. This
will be discussed in more detail in the Biomembrane section of this course.
Esters from bivalent acids (like phthalic or terephthalic acid) and bivalent alcohols
(like ethylene glycol or polyethylene glycol) can form long chains of polyester.
These are used for making artificial fibres (Dacron, Terylene), foil (Mylar), soft
drink bottles (PET) and glass fibre reinforced plastics. Surgical sutures are also
made with polyester fibres.
Oil of wintergreen contains methyl salicylate as main ingredient. It is used as
R
ointment to treat rheumatism. Acetyl salicylate (aspirin
) is an effective pain killer
and anti-inflammatory agent. It also reduces blood clot formation and stroke, some
research indicates that it has anti-cancer properties.
Esters may also be formed from inorganic acids. Reaction of glycerol with nitric
acid produces glycerol trinitrate, more commonly known as nitroglycerol. This is
a very effective explosive, but also used as medicine to treat angina pectoris, chest
pain due to insufficient blood supply to the heart. Why this works you will learn in
the hormone section of this course.
3.1.4 Acetals and ketals
Hemiacetals and hemiketals contain an OH-group and an OR-group on the same
carbon, acetals and ketals two OR-groups. Since ketals and Hemiketals are formed
from ketones, the carbon in question is linked to two other carbons, in case of
hemiacetals and acetals only one.
The most important example of hemiacetals and -ketals are sugars. Sugars are
polyalcohols with one aldehyde or keto-group on the first or second C-atom, respectively. They can react with the OH-group on the penultimate carbon to form a ring
structure. Ring-shaped hemiacetals and -ketals are more stable than their linear
counterparts, so this is the preferred constitution of sugars. More about this you
will learn in the section on carbohydrate metabolism.
55
O
C OH
+
H O C C O H
H2 H2
C OH
ethylene glycol
(n times)
O
Terephtalic acid
(n times)
O
O
C C O
H2 H2
C
O
C O C C O C
H2 H2
O
+
n H2O
O
C
O
O C C O C
H2 H2
polyethylene terephtalat (PET)
C
O
Figure 3.8: Formation of ester bonds from alcohol and carboxylic acid.
56
O
C
O
O
C
OH
OH
salicylic acid
C
OH
O C CH3
O
Acetyl salicylate
(aspirin®)
O CH3
OH
Methyl salicylate
(oil of wintergreen)
Figure 3.9: An extract from willow trees (Salix ssp.) has been used for centuries to
relieve pain and reduce fever. Today we know that the active ingredient is
salicylic acid. Methyl salicylate (oil of wintergreen) and acetyl salicylate
R
(aspirin
) are derivatives of salicylic acid.
R''
O
R C
H
R'
C OH
H2
O H
R C H
O C R'
H2
Aldehyde
Hemiacetal
R''
C OH
H2
H
O CH2
R C H
+
O C R'
H2
+
H
O
H
Acetal
Figure 3.10: Formation of a hemiacetal from an aldehyde and an alcohol. This reaction does not produce water and proceeds in the absence of a catalyst.
The reaction is easily reversible. The hemiacetal can react with a further alcohol molecule to form an acetal, which is much more stable.
However, this reaction requires protons as catalyst and produces water
as side product. Ketones can form similar compounds, which are called
hemiketals and ketals.
57
R'
H
H
N
N
H
CH2
H
H
N
CH2
H
R
ammonia
primary amine
R'
R
CH2
CH2
N
H2C
R''
R'
CH2
R
tertiary amine
CH2
R
secondary amine
+
C N C R'''
H2
H
CH2 2
R''
quartenary amine
Figure 3.11: By replacing the hydrogens of ammonia with organic chains we get
primary, secondary and tertiary amines. Quaternary amines are made
by binding a chain to the free electron pair of nitrogen, they carry a
positive charge.
3.1.5 Nitrogen containing compounds
Amines
Amines are derivatives of ammonia, where one or more of the hydrogen atoms have
been replaced by carbon chains.
In the older literature the organic residues were treated as substituents of ammonia, thus compounds had names like methylamine, ethylmethylamine and so on.
According to IUPAC nomenclature however, the longest carbon chain is name giving,
resulting in names like 2-aminopropane or N,N-dimethyl-2-aminopropane.
Amines are polar and can form hydrogen bonds, but less strong than alcohols.
Thus amines have higher melting and boiling points than alkanes but lower than
alcohols. The exception are tertiary amines, which can not form hydrogen bonds
and have boiling points similar to alkanes of equal molecular weight.
If the carbon chains are not longer than 6, amines are water soluble.
Low molecular weight amines are volatile liquids with fishy odour (indeed the
smell of ‘ripe’ fish is caused by dimethylamine). 1,4-diaminobutane (putrescine) and
1,5-diaminopentane (cadaverine) are responsible for the smell of decaying animals.
Like ammonia, amines (except of course the quaternary) can accept a proton to
become positively charged ions. In this case the -amino is replaced by -ammonium in
the name. In this form amines are more water soluble than their parent compound.
Many pharmaceuticals contain amino-groups, and they are usually given as salts,
rather than free bases.
58
HO
H3C
O
H
N
N CH3
N
HO
Nicotin
Morphine
CH3
N
CH3
N
H
C CH2
OH
O
H2C H C
O
C
HO H
C
O
N
Quinine
Atropine
O
H3C
O
H3C
O
CH3
H
N
N
N
CH3
Caffein
N
C
CH3-CH2
N
CH3-CH2
N
N
lysergig acid diethylamide (LSD)
Figure 3.12: A few examples of alkaloids used as (legal or illegal) drugs.
Quaternary amines are of course always charged. Some of them are used as disinfectants and cationic detergents. One example is the soap that surgeons use for
“scrubbing” before an operation.
The secondary metabolism of plants can produce a wide variety of compounds
with amino groups, these are known as alkaloids because of their ability to accept
protons. Alkaloids can be highly toxic, but in small doses some are important drugs.
Quinine, atropine, caffeïne, nicotine or morphine are but a few examples of alkaloids
(see fig. 3.12).
Amides
Amides can be thought of as reaction products of amines and acids, in a similar
way as esters are formed from alcohol and acid. However, their formation requires
more reactive compounds than acids, usually acid chlorides or anhydrides are used
(R−NH2 + Cl−CO−R0 → R−NH−CO−R0 + HCl).
Technically important are polyamides produced by condensation of diacid chlo-
59
rides with diamines, for example nylon.
Amino acids and proteins
Amino acids contain both a carboxylic acid and a amino-group. Most naturally
occurring amino acids have the amino group on carbon 2, the α-carbon. The simplest
amino acid is glycine NH2 −CH2 −COOH. All other amino acids can be derived from
this compound by replacing one of the hydrogens of the α-carbon with something
else. If this is done, the α-carbon has 4 different substituents, and becomes chiral
(see section on stereochemistry for more explanations). Suffice it here that the amino
acids in our body usually have the L-configuration, but D-amino acids occur in
bacterial cell walls, neurotransmitters and some antibiotics.
If the carboxylic acid of one amino acid reacts with the amino group of another,
a peptide is formed. If more amino acids are added, we get proteins. Because of
their central role in biology an entire section of this course will deal with protein
structure and function.
60
NH2
O
Cl
CH2
O
C C C C C C Cl
H2 H2 H2 H2
+
CH2
CH2
CH2
CH2
n adipoyl chloride
CH2
NH2
n hexamethylenediamine
(1,6-diaminohexane)
O
O
H
C C C C C C N
H2 H2 H2 H2
CH2
CH2
CH2
+
n HCl
CH2
CH2
CH2
O
H
N C C C C C C N
H
H2 H2 H2 H2
O
CH2
Nylon 66
CH2
CH2
CH2
CH2
CH2
N
H
Figure 3.13: Polyamides like nylon 66 contain amide linkages. The technical product
has a molecular weight of about 10 kDa. 3 × 1012 kg nylon are produced
annually worldwide.
61
4 Stereochemistry
4.1 Chirality
If (and only if) all 4 bonds of a carbon atom carry different ligands then there are
two different ways to arrange the ligands on the bonds. This results in two different
molecules with the same total composition, but different 3-dimensional structure.
Each of these molecules looks like the mirror image of the other one, in the same way
as your right hand looks like the mirror image of your left. Therefore the phenomenon
is called chirality (= handedness). The two compounds are called enantiomers.
It is customary to mark enantiomeric carbon atoms in structural formulas with a *.
4.1.1 Chiral compounds and polarised light
Because chiral molecules are unsymmetrical, they can rotate polarised light (see fig.
4.2). The different stereoisomers, when purified, rotate polarised light by different
angles. This was first recognised by Louis Pasteur in 1883. He found that the
crystalline precipitate that forms in wine during maturation (potassium hydrogen
tartrate) contained 2 different sorts of crystals, which have the same shape except
that they are mirror images of each other. Using fine tweezers, he separated these
two forms and showed that they had the same chemical composition, and also gave
the same chemical reactions. Indeed, apart from the different crystal shape the only
other difference was that solutions of these crystals gave different rotation angles
20
= −32◦ and 32◦ for the D and L form respectively). A
with polarised light (αD
mixture with equal concentrations of both forms did not rotate polarised light at
all. Such a mixture of equal amounts of the two forms of a chiral molecule is still
called a racemate today (from the Latin word for grape). Pasteur proposed that
tartaric acid is an asymmetric molecule, and that the different crystals contained
the two different forms of this molecule. It took until 1951 before the structure of
these crystals was solved by X-ray crystallography, proving that Pasteur was right
from the beginning.
4.1.2 Biological and medical importance of chiral compounds
The two enantiomers of a chiral compounds can have substantially different biological effects. A particular infamous example is Thalidomide (trade name Contergan,
fig. 4.3). This substance was marketed as an anti-emetic drug (prevents vomiting)
to pregnant woman suffering from morning sickness from 1958 to 1962, when it
63
4 Stereochemistry
A
A
D
C
B
B
C
D
Figure 4.1: The binding electrons of a carbon atom are arranged in the corners of
a tetrahedron. If all 4 bonds carry different groups 2 different molecules
are possible. These look like mirror images of each other.
Figure 4.2: Schematic diagram of a polarimeter. Unpolarised light from a light source
is passed through a polariser filter, which lets only light swinging in one
direction pass. After interacting with the sample the light goes through
a second filter (called the analyser), which is oriented at right angle to
the first. If the sample did not rotate the light, then all the light which
passed the polariser will be stopped by the analyser, and the observer will
see a dark field. If however the sample is optically active light will pass
the analyser and can be seen by the observer. By turning the analyser
until the field is dark again, the angle of rotation can be determined. By
convention light of 589 nm (Na D-line) is used, the temperature is 20 ◦C
and the sample cell is 10 cm long. If the light is rotated clockwise, the
angle is given a positive, otherwise a negative sign.
64
4.1 Chirality
O
H2 H2
H C C
C O
N C*
C N
H
O
O
Thalidomide (Contergan®)
Figure 4.3: Medicine or poison? In case of Thalidomide, this is not only a question
of dose, but also of using the correct enantiomer. The right-handed is
an anti-emetic, immune-modulatory and analgetic drug used to treat
AIDS, cancer and leprosy. The left-handed causes severe birth-defects
when taken during pregnancy.
transpired that this compound had caused some 10 000 cases worldwide of babies
born with severe defects, in particular phocomelia (“seal-limbs”, hands and feet directly attached to shoulders and hip). It was known that this compound contains
an enantiomeric carbon, but because it is difficult and expensive to synthesise pure
enantiomers, a racemic mixture was used as drug. Later research revealed, that only
one of the enantiomers is a useful anti-emetic (prevents vomiting), anti-inflammatory
and analgetic (pain-killing) drug. The other enantiomers is responsible for the birth
defects. Today the drug is used again against leprosy, certain autoimmune diseases
and the wasting seen in AIDS and some cancer patients. Great care must be taken
that female patients of childbearing age, if given this drug, use appropriate birthcontrol. Unfortunately, under physiological conditions Thalidomide auto-racemises,
there is therefore no point in giving the pure right handed isomer.
4.1.3 Nomenclature of chiral compounds
The +/- system
Because of the high importance of enantiomers in biochemistry it is essential to
name them unambiguously. One possibility is to name them by the angle of rotation
of polarised light (clockwise + or l (laevo = left), counterclockwise - or d (dexter =
right)). However, this does not tell us anything about the structure of the molecules.
Fischer-projections and absolute conformations
Following a suggestion of Emil Fischer, absolute configurations are derived by
comparing a molecule to glyceraldehyde, the most simple sugar, as seen in figure
4.4. D-compounds can be derived from D-glyceraldehyde, L-compounds from Lglyceraldehyde.
65
4 Stereochemistry
O
C
H
HO C H
O
C
H
H C OH
CH2
CH2
OH
OH
L-Glyceraldehyde
O
C
OH
H2N C H
D-Glyceraldehyde
O
C
H C NH2
CH3
CH3
L-Alanine
OH
D-Alanine
Figure 4.4: The D/L-system for stereoisomers. The carboxy-group of alanine is derived from the aldehyde group of glyceraldehyde by oxidation, the βcarbon by reduction. These simple chemical operations allow the alignment of the isomers of alanine with those of glyceraldehyde.
Following Fischer’s suggestion, structures of chiral compounds are drawn in a
particular way to make their identification easier:
• Orient the carbon chain vertically, the carbon atom with highest oxidation
state should be on top.
• Turn the molecule along its long axis, until the substituents of the chiral carbon
face you.
• Project the molecule onto the paper plane
Note that capital letters are used to describe absolute configurations, small letters
are used to describe the rotation of polarised light. This difference is important,
because it is not possible to deduct the rotation of polarised light from the absolute
configuration. Some compounds with L-configuration rotate polarised light to the
right and vice versa.
The R/S-system
In molecules with more than one chiral carbons Fischer’s absolute conformation
is not sufficient for describing a molecule unambiguously. For this reason the R/Ssystem was introduced.
• Number the substituents of a chiral carbon according to their priorities (OH
> NH2 > COOH > CHO > CH3 > H)
66
4.2 Double bonds and cyclic compounds: cis/trans isomery
1
1
4
2
4
3
3
2
Figure 4.5: The R/S system for stereoisomers. Substituents are numbered according
to priority. If the numbers increase clockwise, the substance has the R
(lat. rectus = right) conformation, if the numbers increase counterclockwise, the configuration is S (lat. sinister = left).
H
X
C C
Y
H
H
X
trans
C C
H
Y
cis
Figure 4.6: Rotation around a double bond is not possible. Thus compounds with
several substituents can occur in different configurations, which may have
different biochemical properties.
• Turn the molecule so that the substituent with the lowest priority faces away
from the observer.
• If the priorities increase clockwise, the molecules configuration is R, otherwise
S
• Repeat this process for each chiral carbon.
4.2 Double bonds and cyclic compounds: cis/trans
isomery
A double bond is formed if two sp2 hybridised atoms react with each other, resulting
in bond formation between one hybrid orbital from each of the atoms plus the porbitals.
Rotation around such a double bond, unlike a single bond, is not possible, thus
compounds with several substituents can occur in different configurations.
This is of considerable biological significance. For example in our eyes absorption
67
4 Stereochemistry
of light energy results in the conversion of 11-cis-retinal into all-trans-retinal. This
is the beginning of a signaling cascade that results in vision.
4.2.1 Triple bonds
A triple bond is formed when two atoms with sp-hybridisation react with each
other. Triple bonds, like double bonds, can not rotate. However, since there is only
one substituent on either atom, no cis/trans-isomery can occur.
68
5 Appendix
5.1 List of symbols
c
e
E
~
E
∆G00
h
k
Kd
Keq
Kp
Kw
l
m
M
NA
n
pH
pOH
pKa
pKb
Q
r
R
t
t1/2
T
v
V
z
λ
τ
concentration (mol/l)
elementary charge (1.6022 × 10−19 C/mol)
energy (J)
potential difference (V)
electrical field strength (V/m)
Gibbs free energy (under standard biological conditions, J/mol)
Planck’s quantum (6.6262 × 10−34 J/s)
Boltzmann constant (1.3807 × 10−23 J/K )
reaction velocity constant (unit depends on order of reaction)
dissociation constant (M−1 )
equilibrium constant (unit depends on number of reactants)
partitioning coefficient (pure number)
ion product of water
length (m)
mass (kg)
molecular mass (pure number, but Da is often used)
Avogadro’s number (6.022 × 1023 mol−1 )
number
hydrogen ion tension (pure number)
negative logarithm of [OH- ]
strength of an acid (pure number)
strength of a base (pure number)
electrical charge (C)
radius (m)
universal gas constant (8.3143 J mol−1 K−1 )
time (s)
half life period (s)
absolute (thermodynamic) temperature (K)
reaction velocity (mol/s)
volume (l)
number of elementary charges transferred in a reaction
wavelength (nm)
relaxation time (s)
69
5 Appendix
5.2 Periodic System of the Elements
Figure 5.1: The periodic system of the elements. Metals grey, non-metals green.
Brown denotes elements with both metallic and non-metallic properties.
70
5.3 Acronyms
5.3 Acronyms
ABS acryl butadien styrole copolymer
ADP adenosine diphosphate
AIDS acquired immuno-deficiency syndrome, disease cause by infection with a
retro-virus, HIV
AMP adenosine monophosphate
ATP adenosine triphosphate
dATP deoxy-ATP, used for making DNA
DNA desoxyribonucleic acid
ER endoplasmic reticulum, intracellular membrane system
FAD flavin adenine dinucleotide, oxidized
FADH2 flavin adenine dinucleotide, reduced, cofactor of flavo-proteins
FMN flavin mononucleotide, cofactor in some oxidoreductases
Glc glucose, a hexose, essential component of food
HIV Human immuno-deficiency virus, causative agent for AIDS
IR infrared, light with wavelength 1 mm > λ > 780 nm
IUBMB International Union for Biochemistry and Molecular Biology
IUPAC International Union of Pure and Applied Chemistry
NAD+ nicotinamide adenine dinucleotide, oxidized
NADH + H+ nicotinamide adenine dinucleotide, reduced, soluble carrier of activated hydrogen in catabolic reactions
NADP+ nicotinamide adenine dinucleotide phosphate, oxidized
NADPH + H+ nicotinamide adenine dinucleotide phosphate, reduced, soluble carrier of activated hydrogen in anabolic reactions
NAD(P)H either NADH or NADPH
PET polyethylene terephtalate, polyester used for bottles
RNA ribonucleic acid
71
5 Appendix
US United States (of America)
UV ultraviolet, light with wavelength 380 nm > λ > 1 nm
WHO World Health Organization
72
Index
acetal, 55
acetyl salicylate, 55
acid, 20
corresponding, 22
adiabatic, 7
alcohol, 48
aldehyde, 52
alkane, 44–47
branched, 44
cyclic, 45
straight, 44
alkene, 47, 53
alkyne, 48
amide, 59
amine, 58
amino acid, 60
amphipatic, 41
anion, 32
antibiotic, 45
R
aspirin
, see acetyl salicylate
atom, 29
base, 20
corresponding, 22
battery, 18
primary, 18
secondary, 18
Bayer, F.-tension, 45
Benedict, S.R., 52
bond
π, 36
σ, 36
covalent, 31
dative, 35
double, 36, 67
conjugated, 38, 48
hydrogen, 39
hydrophobic, 41
ionic, 31
polarised, 39
tripple, 36, 48, 68
van der Waals, 40
Brown, R., 13
buffer, 24
carbohydrate, 55
carbolic acid, see phenol
carboxylic acid, 52
cation, 32
cell, electrochemical, 13
Le Chatelier, H.L., 11
chirality, 60, 63
codeine, 54
composition, 43
configuration, 43
conformation, 43
constitution, 43
delocalisation, of electrons, 38
Demokrit, 29
detergent, 41
Dewar, Sir James, 7
dexter, 65
electrode, 13
electrolysis, 19
electron, 29
electronegativity, 32
element, 29
emf, 15
enantiomer, 63
73
Index
energy, 5, 12
activation, 10
free, 8, 12
of reaction, 5
enthalpy, 7, 12–13
entropy, 12–13, 38, 41
equilibrium, 5, 10
ESR, 42
ester, 53
ethanol, 51
ether, 53
ketone, 52
kinetics, 5
FAD, 17
Faraday, M., 15
fat, 55
Fehling, H., 52
fetal alcohol syndrome, 51
Fischer, E., 65
fuel cell, 12
mass action, 9–13
membrane, 55
mesomery, 38, 48
methyl salicylate, 55
mitochondrium, 16
morphine, 54
Gibbs, J.W., 12
Gibbs-Helmholtz-equation, 8
glycerol, 51
Hasselbalch, K.A., 25
heat, 5
capacity, 5
of evaporation, 8
Heisenberg, W.K., 31
hemiacetal, 55
hemiketal, 55
Henderson, L.J., 25
heroin, 54
Hund, F.H., 32
hydrophilic, 41
hydrophobic, 41
indicator, 27
isomery
cis-trans, 47, 67
Kekulé, A., 38
Kelvin, Sir W. Thompson Lord,
12
ketal, 55
74
laevo, 65
light
polarised, 63
Linde, C. von, 40
lipid, 55
lipophilic, 41
Lister, J., 51
lyase, 53
Na/K-ATPase, 16
narcosis, 53
Nernst, W., 16
neutron, 29
nitroglycerol, 55
noble
gas, 31
metal, 15
nucleus, 29
nylon, 60
orbital, 29, 31
hybridisation, 33
oxidation, 15, 48
passivation, 19
Pasteur, L., 63
penicillin, 45
Pfitzer-tension, 47
pH, 21
phase, 27
phenol, 51
phospholipid, 55
plastid, 16
polarimeter, 64
polarography, 19
Index
polyaddition, 48
polyamide, 59
polycondensation, 48
polyester, 55
polymerisation, 48
proton, 29
ion product of, 20
willow, 57
wintergreen, oil of, 55
work, 5
quantum, 29
Raman, 43
reaction
addition, 48
condensation, 59
polymerisation, 48
rectus, 67
redox-reaction, 15
reduction, 15
salicylic acid, 57
Schrödiger, E., 29
selenol, 51
sinister, 67
soap, 41, 53
spectroscopy
electron spin resonance, 42
spin, 31
stereoisomery, 47
substituent, 44
sugar, 55
superconductivity, 47
system, 7
temperature, 5
thermodynamics, 5–13
first law of, 5, 12
second law of, 8, 13
third law of, 8
thiol, 51
Tollens, 52
vaseline, 44
van der Waals, J.D., 40
water
anomaly of, 40
auto-dissociation of, 20
75

Basic chemistry - Ross University

Transcription

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