March 24:
Ch. 8.6 Lewis structures - RULES!!!!!!! Know
them. Formal charges -
know how to calculate them. Following the rules for writing (drawing)
Lewis structures may give us a correct (valid) Lewis structure, which may
not necessarily
be the BEST possible Lewis structure. Formal charges allow to find the best
Lewis structure among several valid ones. Lewis structures for H2O, NH4
+ and C2H2. More examples to come after
the Spring Break.
March 22: EXAM 2 INFORMATION HW3 was collected. Ch. 8.4-8.5 More about bond polarity
- we looked at electronegativity differences as a rough estimate of bond
polarity. A true measure of a bond polarity is a DIPOLE MOMENT,
m, which is a vector (has both length and direction). At this point,
we will deal with dipole moments qualitatively only (no calculations). If you have done any vector algebra
before, review it. We will need it in the near future to determine molecular
polarities.
March 20: Ch.8.4 - Covalent bonds - typical
among NOMETALS Electrons are shared between atoms,
and sharing can be: 1) EQUAL - if atoms in a bond are
the same. The bond is pure covalent (or nonpolar) 2) UNEQUAL - of atoms in a bond are
DIFFERENT (most common scenario). The bond is polar covalent. New term: ELECTRONEGATIVITY The degree to which electrons are
shifted towards towards the more electronegative atom (ie. bond polarity)
depends on the electronegativity difference between atoms in a bond.
Know the electronegativity trends in the periodic table. Be able to look
at a couple of atoms in a bond and judge whether the bond is going to be
pure covalent or polar covalent. If the electronegativity difference is too
large (> 1.8 or so), the bond may be ionic, esp. if one of the atoms was
a metal and the other a nonmetal. The "rule of thumb" is that nonmetals
form covalent bonds, while metals and nonmetals, when bonded to each other,
form ionic bonds. There are always some exceptions
to any rule, even the "rule of thumb" . Blindly following rules of any
kind can get you in trouble!
March 18: Quiz
4! Ch. 8.2: Electron configurations
of transition metal (TM) cations and post-TM cations. Ch. 8.3: Sizes of ions. You can
compare isoelectronic ions in terms of their sizes:
the larger the positive charge,
the smaller the cation
the larger the negative charge,
the larger the anion
Otherwise, we can compare ions with
the same charges for atoms from the same group in the periodic table and
use common sense (that is F- is expected do be smaller than Cl
-, or Mg2+ is expected to be smaller than Ca2+
, just like RA(F) < RA(Cl) and RA(Mg)
< RA(Ca)) NOTE: IONIC COMPOUNDS OCCUR AS IONIC
LATTICES - there are no individual molecules of ionic compounds. Instead,
in order to MAXIMIZE ATTRACTIONS between oppositely charged ions, and MINIMIZE
REPULSIONS between same charge ions, ions are arranged in 3D lattices (Fig.
8.3)
March 15: HW3 was
handed out. You can also download it here
. Ch. 8.1 OCTET RULE revisited, Lewis
electron dot symbols. Ch. 8.2 Ionic Bonds: which atoms
form which kind of ions? Lewis electron dot symbols were used to show the
formation of cations and anions ions, and subsequently the formation of an
ionic salt. Lattice energy (see Table 8.2)
What solubility (high or low) would you expect from an ionic solid with a
really high lattice energy? Analyze Table 4.1 and see if there is any correlation
between low solubility of ionic compounds and the large magnitude of the charge
on at least one of the ions.
March 13: Electron affinity -
know the basic facts (trends are less clear) Be able to explain, using previously
acquired knowledge, why certain trends, especially those in RA
and I1
are seen. A very good place to start would be theelectron configurations
of atoms. Keep asking yourself questions about what you have already learnt
in this course that would allow you to explain the trends. Ch. 7.5 Comparison of major physical
and chemical properties of metals and nonmetals.
Associate metals with bases and nonmetals in acids (this will come handy
in CHEM 1202). Ch. 7.6 Group trends - group 1A
(1) We have analyzed, starting with the
electron configurations) trends in atomic sizes and first ionization energy. Reactivity (the ease with which +1
cations were formed when elements in the group 1A react) was related to the
trends in the first ionization energy (the lower the I1
, the more reactive the element towards species which gladly accept an electron,
eg. halogens) Implications the high reactivity
of elements in group 1A in the natural occurrence (we always find them in
salts or minerals, and never as pure metals). Read Ch. 7.6 to the end and Ch.
7.7 on your own and try to analyze information in a similar way we did in
class for group 1A metals. March 11:
Introduction to Ch. 7. Please read Ch. 7.1 Ch. 7.2 Sizes of Atoms (Bonding Atomic
Radius, RA); trends: increases from top to bottom within a group,
and decreases from left to right within a row (period).
Can you explain these trends using electron configurations, concept of screening
and Coulomb's law (electrostatics) in your arguments? Ch. 7.3 First Ionization Energy (
I1); trends are OPPOSITE to those for RA.
Why? Ionization energy increases from
bottom to top within a group and from left to right within a period. Within periods, the largest ionization
energies can be found for noble gases. Why? Within periods, the lowest ionization
energies can be found for the alkali metals.
Why? Second ionization energy is always
higher than the first ionization energy, etc.: I1 <
I2 < I3 < ... Be able to to analyze and draw conclusions
from the following: Figure 7.5, Table 7.2 and Figures 7.6 and 7.7 If ever in doubt, write anelectron
configuration, focus on the valence electrons, and judge what is to be accomplished
by an atom losing, or gaining (next class) an electron. OCTET RULE - satisfied if an element
has ns2 np6
(or ns2
(n-1)d10 np6
) valence electron
configuration. Satisfying the octed rule is VERY
OFTEN a driving force behind atoms either losing or gaining electrons. March 8:
HW2 was collected. Electron configurations for elements
in the first two rows in the periodic table, noble gases up to Xe, and selected
other elements. New concept:
VALENCE electrons -
know the rules of which electrons in the entire electron configuration belong
to the valence! Writing and drawing full electron
configurations, eg. for potassium, K: 1s2 2s2 2p
6 3s2 3p6 4s
1and short-hand electron configurations,
eg.: [Ar] 4s1 Remember - you can only use
the preceding noble gas when writing short-hand electron configurations. In the above example, K is in the
4th period (row), and the preceding noble gas closed the 3
rd period and is Ar. Valence electrons determine the
chemistry of the element, that is the types of reactions elements enter, the
kind of ions they form or oxidation numbers they adopt.
March 6: Quiz 3! Comparison of energy diagrams for
H atom and many-electron atoms: For the H atom - orbitals are DEGENERATE
as long the principle quantum number, n, is the same. eg. E(3s) = E(3p) = E(3d) In many electron atoms, orbitals
are only DEGENERATE if the azimuthal quantum number, l, is the same. eg. E(3s) < E(3p) < E(3d) Why? Screening (aka shielding) -
"inner electrons" screen outer electrons from the full charge (and attraction)
of the nucleus. The less attration, the higher the energy. Eg. 3p electrons
are screened better than 3s electrons, and hence are less attracted to the
nucleus (see less of the nuclear charge), and have higher energy than the
3s electrons which tend to be closer to the nucleus and experience its larger
charge. Mnemonic for relative sublevel energies
(energies of orbitals within a subshell are the same as energy of the subshell). Hund's Rule. Electron configurations for the
elements in the first two rows (periods). Can you write the electron configuration
for nitrogen? Can you draw it using arrows in boxes? March 4:
Quantum numbers and their meaning: n - energy; l - shape; m
l - orientation is space All three are needed to describe
an orbital. For instance, (1, 0, 0) signifies a 1s orbital. Ways of representing orbitals: charge
clouds, plots of y2, balloon pictures
- know how to draw 2-dimentional renderings
of the balloon pictures for s, p and d orbitals. Allowed (permissible) combinations
of n, l, and ml Electron spin, ms (spin quantum number)
- fourth quantum number with describes an electron in an orbital. Pauli's Exclusion Principle. March 1:
HOMEWORK #2 was ANNOUNCED (Due on March 8). You can still pick it up from me
or
ACCESS IT ON-LINE It was a HEAVY CLASS - plenty of
new and abstract terms:
Line spectrum of the hydrogen atom (Balmer series observable in the visible
part of EMR corresponding to the electronic transitions between n = 3, 4,
5, 6 and n = 2)
Matter Waves (qualitatively
only)
Heisenberg's Uncertainty
Principle
Quantum Mechanics - Schroedinger
Equation and WAVE FUNCTION aka ORBITAL
Quantum numbers: n, l
, and ml - know the rules
for which values thay can adopt.
