Periodic Properties of the Elements

1/18/13 Periodic Proper.es of the Elements How Orbitals Fill The Basis of the Periodic Table Trends in Atomic Proper.es of Elements Mendeleev is given credit for the Periodic Table The following slides we shall skip through fast because... •  ...they were stolen from the internet and •  I cannot find their source. •  But the informa.on is useful! 1 1/18/13 The History of
the Modern
Periodic Table During the nineteenth century,
chemists began to categorize the
elements according to similarities
in their physical and chemical
properties. The end result of
these studies was our modern
periodic table. Johann Dobereiner In 1829, he classified some elements into
groups of three, which he called triads.
The elements in a triad had similar
chemical properties and orderly physical
properties. (ex. Cl, Br, I and
Ca, Sr, Ba) Model of triads 1780 - 1849 2 1/18/13 John Newlands In 1863, he suggested that elements be
arranged in “octaves” because he noticed
(after arranging the elements in order of
increasing atomic mass) that certain
properties repeated every 8th element. Law of Octaves 1838 - 1898 John Newlands Newlands' claim to see a repeating pattern was met
with savage ridicule on its announcement. His
classification of the elements, he was told, was as
arbitrary as putting them in alphabetical order and
his paper was rejected for publication by the
Chemical Society.
1838 - 1898 Law of Octaves John Newlands His law of octaves failed beyond the
WHY?
element calcium. Would his law of octaves work today with
the first 20 elements? 1838 - 1898 Law of Octaves 3 1/18/13 Dmitri Mendeleev In 1869 he published a table of
the elements organized by
increasing atomic mass. 1834 - 1907 Lothar Meyer At the same time, he published his own
table of the elements organized by
increasing atomic mass. 1830 - 1895 •  Both Mendeleev and Meyer arranged
the elements in order of increasing
atomic mass.
•  Both left vacant spaces where unknown
elements should fit.
So why is Mendeleev called the “father
of the modern periodic table” and not
Meyer, or both? 4 1/18/13 Mendeleev... •  stated that if the atomic weight of an
element caused it to be placed in the
wrong group, then the weight must be
wrong. (He corrected the atomic
masses of Be, In, and U)
•  was so confident in his table that he
used it to predict the physical
properties of three elements that were
yet unknown.
After the discovery of these unknown
elements between 1874 and 1885, and the
fact that Mendeleev’s predictions for
Sc, Ga, and Ge were amazingly close to
the actual values, his table was generally
accepted. However, in spite of Mendeleev’s great
achievement, problems arose when new
elements were discovered and more
accurate atomic weights determined. By
looking at our modern periodic table, can
you identify what problems might have
caused chemists a headache? Ar and K Co and Ni Te and I Th and Pa 5 1/18/13 Henry Moseley In 1913, through his work with X-rays, he
determined the actual nuclear charge
(atomic number) of the elements*. He
rearranged the elements in order of
increasing atomic number.
*“There is in the atom a fundamental quantity which
increases by regular steps as we pass from each
element to the next. This quantity can only be the
charge on the central positive nucleus.”
1887 - 1915 Henry Moseley His research was halted when the British
government sent him to serve as a foot
soldier in WWI.
He was killed in the
fighting in Gallipoli by a sniper’s bullet, at
the age of 28. Because of this loss, the
British government later restricted its
scientists to noncombatant duties during
WWII. Glenn T. Seaborg After co-discovering 10 new elements, in
1944 he moved 14 elements out of the
main body of the periodic table to their
current location below the Lanthanide
series. These became known
as the Actinide series.
1912 - 1999 6 1/18/13 Glenn T. Seaborg He is the only person to have an element
named after him while still alive.
"This is the greatest honor ever bestowed upon me even better, I think, than
winning the Nobel Prize."
1912 - 1999 Here is a simple periodic table... Comments... •  This is an older style table. •  The rows are called “periods.” •  The columns are “families” or “groups.” –  Type A: representa.ve –  Type B: transi.on –  The guys at the boSom are inner transi)on elements. 7 1/18/13 Elements in Columns Have Similar Proper.es •  This is especially true of type A (representa.ve) elements. •  But, there are changes in reac.vity, etc., as one goes down a column. •  This link gives group IA as an example... hSp://www.youtube.com/watch?v=Ft4E1eCUItI •  More about ver.cal trends later! Looking at the rows... •  The rows are called periods. •  How they fill gives insight into the atomic structures. •  In par.cular, we shall examine the idea of electron configura)ons. Some Points... •  As just stated, rows of the periodic table are called periods. •  These have the following lengths: o 2 o 8 o 18 o 32 8 1/18/13 Look at these numbers in detail... • 
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2 = 2 x 12 8 = 2 x 22 18 = 2 x 32 32 = 2 x 42 •  This gives some insight into quantum numbers! In the last chapter we saw... •  3 quantum numbers, n, l, and ml. •  For each value of n we have n2 orbitals. •  We play some number games here: v n = 1 → 1 orbital v n = 2 → 4 orbitals (4 = 1 + 3) v n = 3 → 9 orbitals (9 = 1 + 3 + 5) v n = 4 → 16 orbitals (16 = 1 + 3 + 5 + 7) •  These are all ½ the lengths of the periods! What’s wrong? Solu.on to this problem... •  H has just one electron. •  The Schrödinger equa.on seen earlier needed another quantum number to handle mul.-­‐
electron atoms. (Technical detail: This comes naturally if rela.vity included.) •  This new quantum number is ms, the electron spin. This has values of = +1/2 and -­‐1/2. •  The physical explana.on is lee for lecture! 9 1/18/13 Some Pictures of Spin Schema'c Picture Language o2en used: Spin up, Spin down Spin +1/2, Spin -­‐1/2 α β |+ > |-­‐ > We observe spin states with a magne.c field quite oeen Spins precess about the magne.c field... 10 1/18/13 Other depic.ons of spin Electron spin discovered in the Stern-­‐
Gerlach Experiment The Guys... O7o Stern Walter Gerlach 11 1/18/13 Effect on energy levels... Single Electron (Schrödinger Produces Degeneracies) Mul'ple Electrons (Degeneracies Li2ed!) Atomic Orbitals & the Pauli Exclusion Principle •  Pauli principle: No two electrons in the same atom can have the same four quantum numbers. •  This means, only two atoms/atomic orbital. •  Orbitals fill in order of increasing energy (part of the AuEau principle—more later!). •  E(s orbital) < E(p) < E(d) < E(f), for a given value of n. Some more guys... Wolfgang Pauli Pauli & Bohr Studying Spin 12 1/18/13 Ordering Changes Caused by Shielding/Penetra.on Shielding Penentra'on An alternate view via radial distribu.on func.ons... Orbital Filling: The AuEau Principle Niels Bohr Friedrich Hund 13 1/18/13 The AuEau Principle •  Electrons fill atoms in order of increasing orbital energy. q 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s, etc. •  Degenerate orbitals fill so as to maximize the number of parallel electron spins. q This is Hund’s rule! Order of filling some.mes visualized as follows: Filling Atomic Orbitals Fairly Simple •  We look at the first few elements. •  Note that we are comparing a)  Electron configura.ons & b)  Orbital diagrams. 14 1/18/13 Elements 3 – 10... Nota.onal Shortcut •  We oeen write in CORE electrons as their INERT GAS SHELLS. •  Examples: –  Na as [Ne]3s1 instead of 1s22s22p63s1 –  Rb as [Kr]5s1 instead of a really long mess! •  We see the first 18 elements to the right... The periodic table shows how we can think in terms of s, p, d, and f blocks... 15 1/18/13 One thing should be obvious... You can construct an electron configura.on directly from the periodic table (usually)! Atoms take on electrons in an orderly manner—
usually but irregulari.es can occur. Do you see any here? A corollary that follows from Hund’s rule & AuEau... •  Half-­‐filled subshells are par.cularly stable. •  Fully-­‐filled subshells are par.cular stable. •  What appears to be an anomaly really is not! 16 1/18/13 Atoms-­‐first Viewpoint Here... •  Electronic configura.ons build atoms. •  Atoms with similar configura.ons are similar in proper.es. •  The periodic table and configura.ons are in.mately related. •  In the next few slides, we shall show a few of these happy families of elements. The most metallic of the metals... Alkali Metals Alkaline Earth Metals The most noxious and snobbiest elements... The Halogens The Noble Metals 17 1/18/13 PreSy Pictures of the Halogens Fluorine Cl, Br, and I One last family •  The last family we list is the chalcogens. •  Some.mes, oxygen is lee out of the group. •  The group portrait is to the right... The black sheep of the families •  The elements in the 2nd row of the periodic table oeen have quite different proper.es than the ones below them. •  The elements to watch here are Li, Be, B, C, N, O, and F. •  Why this happens will be explained later. But, for now: Beware of the 2nd row! 18 1/18/13 Notable things about the excep.ons... •  Li reacts directly with N2. No other element does. •  Be is a deadly poison. Mg and Ca are nontoxic and the others are mildly toxic in this family. •  B, N, C, and O can form double and triple bonds. Elements below them cannot. •  F does not form any ions other than F-­‐. (Cl, for instance, forms ClO-­‐, ClO2-­‐, ClO3-­‐, and ClO4-­‐.) A Useful Aside: Elements that have predictable ion charges... Enough about groups! Let us now discuss PERIODS!!! 19 1/18/13 Trends across periods... •  We go let to right. •  Here are some things we shall examine: –  Atomic size (radii) –  Ion size –  Ioniza.on energies –  Electron affini.es –  Metallic Character •  Trends are generally smooth but there are “bumps” to discuss! Atom Size (Atomic radius) Defini.ons of atomic size: Nonbonding atomic radius (van der Waals) Bonding atomic radius Nonmetals: One-­‐half the distance for two iden.cal atoms bonded together Metals: One-­‐half the distance between nuclei of two touching atoms in the metal crystal Compare Two of These van der Waals radius (Kr) Bromine radius in Br2 20 1/18/13 Trends in Atomic Radius •  A slight surprise: Atoms get smaller from lee to right in a period! •  But, there are a few bumps. •  There is NO surprise with columns. Things get bigger as you go down! (Example: Li < Na < K < Rb < Cs) •  We shall discuss these things verbally as the next two slide are presented. Atom Size (vs. Z) Shown Graphically Atom Size: Periodic Table View 21 1/18/13 Causes of these trends... •  As we move down a column in the periodic table, the principal quantum number of the electrons increases resul.ng in larger orbitals and, hence, larger atomic radii. •  As we move to the right across a row, the effec.ve nuclear charge experienced by the outermost electrons increases (for a given orbital) and, hence, the radius shrinks. (Zeff = Z – s, where s is the screening constant.) Let’s Look at Monatomic Ions •  Ca.ons are posi.ve (atoms lose electrons) •  Anions are nega.ve (atoms gain electrons) •  Most ions form for one of the following reasons: –  Ge|ng an inert gas core –  ASaining a half-­‐filled d-­‐ or f-­‐subshell –  ASaining a fully-­‐filled d-­‐ or f-­‐subshell Examples of Inert Core Ions... •  Each set has the same electronic structure (hence, we call these isoelectronic) •  First set: N3-­‐, O2-­‐, F-­‐, Ne, Na+, Mg2+, Al3+ •  Second set: P3-­‐, S2-­‐, Cl-­‐, Ar, K+, Ca2+, Sc3+, Ti4+ 22 1/18/13 Here are some transi.on metals... Ag loses just 1 electron Zn needs to lose 2 Some At the Board Examples... Diamagne'sm •  Electrons all paired. •  Weakly repelled by magnet. •  Do Al3+ and S2-­‐ as examples. Paramagne'sm •  One or more electrons unpaired (either an odd number or caused by Hund’s rule). •  We look at Fe2+ and Fe3+ as examples. Ionic Radii of Ca.ons (pm) •  Removing an electron obviously makes an ion smaller. •  Each row of atoms to the right consists of isoelectronic sets. •  Ions get smaller as we go to the right. •  Ions get larger as we go down. 23 1/18/13 Ionic Radii of Anions (pm) •  Here the trends are different in that ions get larger as we go to the right. •  They s.ll get larger as we go down. •  All the pairs are isoelectronic. Ioniza.on Energies •  Ca.on forma.on is generally endothermic. •  It takes energy to remove an electron. •  Only valence electrons get removed in ordinary chemical reac.ons. •  Core electrons cannot be removed in ordinary reac.ons! •  The next charts show trends in first ioniza)on energies. First Ioniza.on Energy Trends 24 1/18/13 Second and Higher Ioniza.on Energies •  Only valence electrons can be removed. •  Each successive electron harder to remove since it must now be separated from a ca.on as opposed to a neutral atom. •  Core electrons stay put! •  We show these trends in the next two slides. Comparing Na and Mg Table for Na – Ar (kJ/mol) 25 1/18/13 Electron Affini.es •  This is the energy (usually) given off when a neutral atom receives an electron to become an anion. •  Note that this is also (sign change!) the energy needed to convert an electron back into a neutral atom. •  The next chart shows some of these for X + e-­‐ → X-­‐. A Short Table of E.A.’s (kJ/mol) Three types of elements • 
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Metals Nonmetals Metalloids We discuss all these while looking at the next slide. 26 1/18/13 Trends in Metallic Character A pre|er selec.on! At last... 27