Document 6595753

Compounds
Problems: Lab Manual Appendix III: Part H
I. Classify different forms of matter.
A. Classification based on purity
(of sorts)
B. Element: Cannot be separated into simpler
substances by chemical means.
1. Composed of one type of atom
2. Examples:
copper (Cu)
magnesium (Mg)
All those substances on the periodic table
3. Review: What determines the identity of an element?
Note: With respect to elemental identity, an atom with 6 p+ and 6 n in its nucleus is considered to be the
same type of atom as one with 6 p+ and 8 n in its nucleus. These two atoms are said to both be
of carbon.
4. Being “in its elemental state.”
Elemental state means the state in which that element exists in nature.
The elemental state of:
a) He, helium, is as a gas,
b) Fe, iron, is as a solid
c) Hg, mercury, is as a liquid (Not many elements exist as liquids in nature.)
Note: Seven elements are “diatomic” (two atoms of an elements bonded covalently) in their elemental
state.
“I never order broiled clams from Hardees.”
C. Compounds: Pure substances formed when atoms (of _____________________________________)
combine in specific ratios.
The combining is when the atoms form some kind of________________________.
1. Compounds can be separated into simpler substances by chemical means
2. Examples:
______________________ _______________________
D. Mixtures: Two or more types of substances mixed together in varying ratios.
They can be separated by “physical means”.
1. Homogeneous: Uniform composition (also called a _______________)
Examples:
2. Heterogeneous: Non-uniform composition.
Examples:
E. Most matter on earth & in your body isn’t in its elemental form.
1. It is present as compounds or mixtures
2. If we consider the lecture room (or lab?), can you identify any matter that
is in its pure elemental form?
3. Looking outside the classroom, is the blue stuff in the picture a pure
substance or a mixture?
www.all-creatures.org/hope/index.htm
II. Formulas and Models of Compounds
A. Chemical formula, like C3H8O or H2O
1.
Tells what elements are present in a substance.
2.
The subscripts indicate the number of each type of atom.
B. Structural formula
1.
Tells how the atoms are connected
2.
These connections are called bonds
C. Condensed structural formula
CH3-CH2-CH2OH
D. Line structures
1.
Carbons are located at the ends and junctions of lines.
2.
Hydrogens bonded to carbons are not indicated.
III. Formation of two types of bonds
covalent bonds & ionic bonds
A. Energy & Stability
Things at higher energy are less stable!!
1. Living things are high energy & unstable.
2. They must obtain energy from other unstable & high energy things….ie. food!
(The compounds in the food you eat must be relatively unstable for you to get useful energy from the food.)
3. Generally speaking,
calorically rich food is made up of reduced (hydrogen containing) carbon compounds.
4. The higher the hydrogen content (and lower oxygen content), the more energy the food contains.
(Example: Fat is more caloric than carbohydrate or protein.
5. One way to consider stability:
Stable - if something exists for a relatively long time.
Unstable – if something exists for a brief time and changes.
SO
What makes atoms etc. be stable?
It depends on e−.
6. Focus: Lewis Octet Principle.
For representative elements, a full valence shell is considered stable. Why?
Noble or (inert gases) are inert (non-reactive, stable), and they have full valence shells.
“Unstable atoms gain, lose, or share electrons (forming ions or compounds) to end up with eight valence
electrons (and become stable).”
Remember: The valence electrons (e−) are those in the s & p orbitals of the outermost occupied shell.
Which elements have full valence shells?
How many electrons do they each have?
B. Gain, Lose or Share Electrons
1. A covalent bond is a shared e- pair.
(tug of war)
2. The atoms in molecules are held together by covalent bonds.
Water
glucose
adenine
3. An ionic bond is the attraction between oppositely charged ions.
a. When atoms gain or lose electrons, they become ions with a charge (like +1, +2, +3, or -1, -2, -3, etc).
b. Metals tend to _____to become _____ ions.
c. Non-metals tend to ___ to become ___ ions.
C. As stated above, covalent bonding is between two non-metals.
Ionic bonds are formed between a metal and a non-metal.
Based on these two statements,
are the compounds below are ionic or covalent.
NaCl
N2
H2O
K3N
IV. Covalent Bonding
A. This type of compound results from combining ______________________________
B. Hydrogen gas (H2(g)) reacting with oxygen (O2(g)) to form water (H2O(g)):
2 H2(g) + O2(g)  2 H2O(g)
In the rxn. shown above, two non-metals combine to form a covalent compound. The product is a
molecule.
The Hindenberg was a “air” ship (specifically, a dirigible) whose buoyancy was provided by H2(g).
A known risk associated with this technology was hydrogen reacting violently with oxygen.
Is there a readily available source of the reactant O2 on earth?
Viewer warning, people died.
http://www.youtube.com/watch?v=F54rqDh2mWA
2 H2(g) + O2(g)  2 H2O(g)
V. Lewis Structures of Atoms (e- accounting.)
A. Count valence e- (the e- in the outermost shell).
B. Arrange the e- (dots) on the 4 sides of the elemental symbol.
C. Don’t worry transition metals in CHM 109.
D. Examples:
Xe
Which are stable?
C
F
S
He
VI. Seeing patterns in molecules:
A. Atoms share electrons to fill their valence e- shells.
B. For valence e- purposes, we count the shared e- pair as belonging to both atoms.
What would be the stable form of hydrogen gas? Is the stable form of “H”? Why/why not?
Does H have a full valence shell?
Here we show e- sharing in H2 with pictures.
Lewis Dot
“Flat” (no 3-D info)
H H
HH
Ball & Stick
Space filling (most real?)
VII. Electronegativity and Bond Polarity
(view nuclei as fixed reference points)
A. Covalent bonds involve e- sharing, but how evenly is the negative charge arranged around the bonded
atoms?
How you decide? Electronegativity chart!
Electronegativity: the relative of ability an element to attract electrons in a bond.
(“Sharing a blanket” analogy)
Electronegativity Table (Pauling scale) (p. 106 in text)
1A
1
H
2.1
3
Li
1.0
8A
2
He
2A
4
Be
1.5
11
Na
0.9
12
Mg
1.2
3B
4B
5B
6B
7B
1B
19
K
0.8
37
Rb
0.8
55
Cs
0.7
20
Ca
1.0
38
Sr
21
Sc
1.3
39
Y
22
Ti
1.5
40
Zr
23
V
1.6
41
Nb
24
Cr
1.6
42
Mo
25
Mn
1.5
43
Tc
26
Fe
1.8
44
Ru
27
Co
1.9
45
Rh
28
Ni
1.9
46
Pd
29
Cu
1.9
47
Ag
56
Ba
57
La
72
Hf
73
Ta
74
W
75
Re
76
Os
77
Ir
78
Pt
79
Au
3A
5
B
2.0
4A
6
C
2.5
5A
7
N
3.0
6A
8
O
3.5
7A
9
F
4.0
14
Si
1.8
15
P
2.1
16
S
2.5
17
Cl
3.0
18
Ar
2B
13
Al
1.5
30
Zn
1.6
48
Cd
31
Ga
1.6
49
In
32
Ge
1.8
50
Sn
33
As
2.0
51
Sb
34
Se
2.4
52
Te
36
Kr
80
Hg
81
Tl
82
Pb
83
Bi
84
Po
35
Br
2.8
53
I
2.5
85
At
2.1
<–––– Atomic number
–– Elemental symbol
<––– Electronegativity
<---------8B -------->
10
Ne
54
Xe
86
Rn
B. Non-polar covalent bonds have completely even sharing. Examples:
H 2.1
H 2.1
O 3.5
H2
O 3.5
O2
C. Polar covalent bonds: when bonding atoms have different electronegativities.
H 2.1
I 2.5
δ+
H 2.1
δ-
F 4.0
δ+
HI slightly polar bond
δ-
HF exceedingly polar bond
Continuum: Non-polar covalent bonds
Polar covalent bonds
Formation of ions and ionic bonding
Representing molecules & compound ions w/ Lewis Structures
Lewis Structures: visual bookkeeping device for valence e.
Steps to Writing Lewis Structures:
1. Add up number of valence e- (group #) for each atom
2. Make adjustments for non-zero net charge
a) Add one valence e- for each negative charge
b) Subtract one valence e- for each positive charge
3. Write elemental symbols (which atom links to which?)
4. Form single bonds to connect bonded atoms
5. Fill in non-bonding e- pairs
6. Only if not enough e- to go around, form double or triple bonds
Write Lewis structures for:
Cl atom
Cl1- ion
NO31-
HCN
CH4
H2O
III. Using Lewis dot structures to determine Electron Pair Geometry, Molecular Shape &
Polarity
A. VSEPR stands for Valence Shell Electron Pair Repulsion.
We will use this model to determine 3-D shapes of molecules.
1. VSEPR is based on idea that electron pairs are arranged to be as far apart as possible.
2. Three electron pair geometries (discussed in this class)
a) linear (2 electron pairs)
b) trigonal planar (3 electron pairs)
c) tetrahedral (4 electron pairs)
3. Molecular Shapes (determined by location of atom centers)
a) linear electron pair geometry  one possible molecular shape
linear
b) trigonal planar electron pair geometry  two possible molecular shapes
trigonal planar
bent
c) tetrahedral electron pair geometry three possible molecular shapes
tetrahedral
trigonal pyramidal
bent
4. Try some examples
a) CH4
b) CH3Cl
c) H2O
d) CO2
B. Polarity of molecules
1. Polarity: uneven distribution of charge.
2. Since p+ must stay in nucleus, uneven charge distribution refers to e- distribution around specific nuclei
or the molecule as a whole.
3. You must look at e- pair geometry (molecular symmetry) to determine molecular polarity.
4. Symmetry, vectors or a tractor pull(?)
IX. Ionic Bonds
Elemental sodium (Na(s)) reacting with chlorine (Cl2(g)) to form table salt (NaCl(s)), an ______________
2Na(s) + Cl2(g)  2NaCl(s)
Not safe to do, but see: http://www.youtube.com/watch?v=VBReOjo3ri8
http://www.youtube.com/watch?v=Ftw7a5ccubs&feature=related
&
In the rxns shown above, a metal combined with a non-metal to form an ionic compound as the product
a.k.a. a salt. The ions in the salt are held together by electrostatic attraction of oppositely
charged ions (+ and -).
X. Ionic bonds, ionic compounds, ionic compound formulas
A. In an ion: # of p+  # of e-.
B. Ions can be simple or polyatomic.
1. A simple ion consists of one charged atom (for example Cl-).
2. A polyatomic ion has multiple atoms held together by covalent bonds and carries a charge. (for
example NO3-)
3. positively charged ions are called
negatively charged ions are called
Remember: Ionic bond refers to the electrostatic attraction between oppositely charged ions.
C. Ionic compound formulas must have zero net charge.
(No shock when you touch an ionic compound.)
1. Determine charge of simple ions in ionic components from Periodic Table. How?
2. A few “polyatomic” ions like sulfate ion, SO42-, are listed in the lab manual appendix and on p.78 of
textbook.
3. An ionic compound formula consists of the minimum number of each ionic component that results in
equal amounts of (+) and (-) charge.
Problem. Write ionic compound formulas for:
1. potassium chloride
2. calcium iodide
3. sodium phosphide
4. sodium phosphate
5. magnesium oxide
6. iron (II) nitrate
XI. Ionic Bond Formation
A. Example of ionic compound formation: 2Na(s) + Cl2(g)  2NaCl(s)
Important: 2 distinct things occur in these rxns. You must be able to recognize them as distinct.
B.1. First, e- are being transferred to convert atoms (or molecules) into ions:
Na  Na+ + e-
(ox)
Cl2 + 2 e-  2 Cl-
(red)
metals tend to lose e- (give e- away), non-metals tend to gain e-.
2. Second, positively charged ions bind to negatively charged ions to form ionic compounds:
Na+ + Cl-  NaCl
C. Another example:
Elemental magnesium (Mg(s)) reacting with oxygen (O2(g)) to form magnesium oxide (MgO(s)):
2 Mg(s) + O2(g)  2 MgO(s)
Two steps:
1.
Mg(s)  Mg2+ + 2e-,
2.
Mg2+ + O2-  MgO
O2 + 4 e-  2 O2-
D. Remember: Ionic compounds are not molecules!
Although the ionic compound formula, NaCl, shows a 1 to 1 ratio of the
ions, it is incorrect to think of a specific, directionally-oriented interaction
between one Na+ and one Cl-.
You can see from the figure that each Na+ is actually interacting equally
with 6 Cl- ions (and vice versa).
E. An interesting thing happens to ionic compounds that dissolve in water.
The lowest energy state occurs when the ions separate and each is surrounded by water.
The H2O molecules reduce the strength of attractive or repulsion between ions significantly..
Appendix A: Orbital hybridization
A. Introduction
1. The s and p orbitals we used for isolated atoms does not work for molecules
(ie.when atoms have formed covalent bonds).
2. A model that is useful for understanding these systems is called orbital hybridization.
B. Orbital hybridization  orbital food processor
1. Put 2s orbital w/ appropriate # of 2p orbitals
2. Blend to obtain a set of hybrid orbitals.
3. Number of hybrid orbitals obtained equals number of s + number of p orbitals
4. Unlike p orbitals that have two equal lobes per orbital, sp, sp2, and sp3 hybrid orbitals have one main
lobe (but see Dr. Winter’s Orbitron at http://winter.group.shef.ac.uk/orbitron/) per orbital.
5. The geometries of the hybrid orbitals match the patterns predicted from VSEPR theory.
6. The 2 lobed pi bonds that are part of double & triple bonds accurately predict the limited rotation
observed w/ these types of bonds.
sp3
hybridization
Formation of 4 sp3 hybrid orbitals from 1 s and 3 p orbitals
orbital
hybridization
2p
energy
sp3
2s
s
px
py
pz
orbital
hybridization
geometry
Note: Each p orbital
has two lobes.
Note: Each sp3 hybrid orbital
(in blue) has 1 main lobe.
Molecules that contain sp3 hybridized Period 2 atoms, C and N
CH4, methane
NH3, ammonia
H
N
C
H
H
H
H
H
H
sp2 hybridization
Atoms involved in forming one double bond exhibit sp2 orbital hybridization patterns:
Formation of 3 sp2 hybrid orbitals from 1 s and 2 p orbitals
orbital
hybridization
2p
energy
2pz
2
sp
2s
s
px
py
pz
orbital
hybridization
geometry
Note: Combine one s and two p orbitals.
Note: Each sp2 hybrid orbital
(shown in green) has one main
lobe. pz orbital is unchanged.
Molecules that contain sp2 hybridized Period 2 atoms
CH2O, formaldehyde (a.k.a., ethanal)
H
O
O2, elemental oxygen
O
O
C
H
Although you might conclude from the structures shown above that the two bonds that make up a
double bond are identical, that is not the case. (See below.)
Orbital hybridization: Building CH2O from sp2 hybridized CH2 and O
px
s
py
pz
Note: Combine one 2s with two 2p orbitals
to obtain three sp2 hybrid orbitals.
Note: Each sp2 hybrid orbital
(shown in green) has one main lobe.
Trigonal planar geometry
The pz orbital is unchanged.
gives maximum sp2 orbital
separation.
O
C
H
Formaldehyde, CH2O.
H
Imagine pushing an sp2 hybridized CH2
group and an sp2 hybridized O atom together.
Hybrid orbitals are used for
sigma () bonds (black) & non-bonding
e- pairs (green). Unhybridized orbitals
are used for pi () bonds (red). Note
that the  bond has two lobes.
Rotation about O-C axis???
p
p
O
C
p
p


H
H
H
C
O

H
One bond is formed from the direct overlap of sp2 orbitals, a sigma bond. The other is formed from sideways
overlap of p orbitals, a pi bond.
A good on-line review for this is located:
http://www.chem.uiuc.edu/CLCtutorials/104/Hybridization/SeeIt.html
When a double bond is formed, rotation about that bond is eliminated.
It is the pi bond components of double (and triple) bonds that limit rotation about these bonds.
Contrast this with the free rotation about the C-C single bond in ethane.
The limited rotation in double bonds is responsible for another group of cis/trans isomers. The structures of the
cis and trans isomers of 2-butene are shown below.
(The “-ene” ending indicates that there is a double bond, and the “2" indicates the position of the double bond.)
H
H
H
H
H
C
C
C
H
H
H
C
H
H
C
C
C
H
H
H
H
cis-2-butene
C
H
C
H
C
C
cis
C
trans
trans-2-butene
In general terms, a compound is cis if its higher ranking groups are on the same side of a line running through
the C atoms involved in the double bond, and trans if they are on opposite sides of the line.
Biological relevance of cis & trans
Cis- to trans-retinal in Rhodopsin,
The 1st biochemical step in the visual process
cis- double bond
11
12
light
Opsin (big protein)
trans- double bond
11
12
Opsin (big protein)