Study guide: Exam 3 covers Chapter 8, 9 and 10 (10.3 will not be tested)
You’ll be given a periodic table + a VSEPR chart (see P.3 for chart).
You should be able to:
• Define basic terms: subshell, shielding, penetration, degenerate, isoelectronic, lone
pair of electrons, bonding pair of electrons, valence electron, inner electrons, outer
electrons
• State the difference between the electron-spin quantum number vs. the 3 other
quantum numbers (ie. what does ms describe?)
• Assign 4 quantum numbers to a given electron in an atom
• Explain the difference in energies of atomic orbitals for single and multiple electron
systems
• Explain the causes that leads to the difference in energies of the subshells of atomic
orbitals in multiple electron systems
o Specifically be able to explain why on a given energy level, s is lower in E than
p, which is lower than d, which is lower than f
• List out the order of the atomic orbitals in increasing energy (ie. 1s < 2s < 2p < 3s < 3p
< 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p, etc.)
• Explain the 3 rules that governs how electrons fill up atomic orbitals:
o Aufbau principle, Pauli exclusion principle, Hund’s rule
• Write out the electronic configuration (full or condensed) for any element or ion
o Know the exceptions (Cr, Cu, Mo, Ag, Au – see figure 8.10 on P333)!
o Know that for ions of elements with filled d and f orbitals, you lose the electrons
in the higher energy s & p orbitals first before the d/f, and you would lose the p
before s
• Draw orbital diagrams (full or partial) for any element or ion
o Know the exceptions (Cr, Cu, Mo, Ag, Au – see figure 8.10 on P333)!
o Know that for ions of elements with filled d and f orbitals, you lose the electrons
in the higher energy s & p orbitals first before the d/f, and you would lose the p
before s
• Identify the number of inner/core electrons, outer electrons and valence electrons for
each element or ion
• Predict the charges that ions of main group elements will form based on the electronic
configuration and the octet rule
• Explain how atomic radius are determined (metallic radius and covalent radius)
• Predict the following periodic trend (ie. rank different elements/ions) and explain what
they are and why they occur (+ explain the exceptions and why they occur): Atomic
size, ionic size, ionization energy, electron affinity, metallic behavior, acidity/basicity of
oxides, lattice energy (proportional to coulomb’s law), electronegativity
• Explain why for isoelectronic species, cations are smaller than anions
• Explain the difference between first ionization energy vs. second ionization energy vs.
third… etc. and when an element would expect a big jump in its Xth ionization energy.
• Given a set of ionization energy (1st, 2nd, 3rd, 4th, etc.), be able to predict which period
the element resides
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Explain why the 1st electron affinity is usually negative, why group 2A and 5A have
unusually high 1st EA (high = more positive), and why group 8A has all positive 1st EA.
Explain why the 2nd EA is always positive
Write the equation for metal oxides reacting with water and non-metal oxides reacting
with water. Know what an amphoteric compound is (see P. 346).
Predict whether an element is paramagnetic or diamagnetic.
Explain the ionic bonding model, and how it explains the properties of ionic
compounds
Draw Lewis structures (of reactant and product) representing the formation of an ionic
bond
Give the electronic configuration/ draw the orbital diagram of the elements in an ionic
bond
Identify the steps of a Born-Haber cycle (ionization energy, dissociation/bond energy,
electron affinity, lattice energy).
Explain the covalent bonding model and how that explains the properties of covalent
compound
o Explain Figure 9.12 and identify what the different regions of the graph
represent
Explain the relationship between bond order, bond energy and bond length (and how
one increases/decrease with respect to another)
Identify the bond order of a given bond in a structure
Calculate enthalpy of reaction given bond energy data
Identify whether a compound is an ionic or covalent compound given electronegativity
values of the atoms
Explain the metallic bond model and how it explains the properties of metals
Draw Lewis structure of covalent compounds and polyatomic ions
Decide which resonance structure is the major resonance structure using formal
charge
Determine molecular shape given VSEPR theory
For example test questions, see Dr. Woodbury’s Practice Exam 3 Chapters 8-10
Available at: https://sites.google.com/site/woodburychem1a/practice-exams-1
FYI- The key is available on the same site under: Practice Exam 3 Ch 8-10 Key
# of lone pairs
+ bonds
no lone pair
1 lone pair
2 lone pairs
3 lone pairs
4 lone pairs
AX2
2
linear
180 ˚C
AX3
3
trigonal planar
120 ˚C
AX4
AX2E1
bent
<120 ˚C
AX3E1
AX2E2
4
tetrahedral
109.5 ˚C
trigonal pyramidal
<109.5 ˚C
AX5
AX4E1
trigonal bipyramidal
seesaw
<90 ˚C/<120 ˚C
5
90 ˚C/120 ˚C
AX6
AX5E1
bent
<109.5 ˚C
AX3E2
t-shaped
<90 ˚C
AX4E2
AX2E3
linear
180 ˚C
AX3E3
AX2E4
6
octahedral
90 ˚C
square pyramidal
<90 ˚C
square planar
90 ˚C
t-shaped
<90 ˚C
linear
180 ˚C