AGK’s Fall 2006
Chem 111 Exam 2 Review Sheet
*NOTE: This list is fairly comprehensive but designed only as a study aid. You are responsible for all
material covered in class and in assigned readings. While all exams are comprehensive, the emphasis
will be on the material presented since Exam 1. We have covered so much material since Exam 1 that
there is no way I can ask a question about every topic on the midterm. Don’t be disappointed if you
favorite topic isn’t on the test-----there’s always the final exam!
Chapter 5:
Zeff and electron shielding---understand what they mean and effects they have on electrons.
Know the relative E of orbitals in 1-electron atoms and multi-electron atoms and why they are not the
same. E=-2.18 x 10-18J(Zeff2/n2)
Pauli exclusion principle- why it exists and what it means (only two electrons/orbital).
Aufbau Principle-order of filling. Be able to express electron config. in atoms and ions (know exceptions
covered in class) by both 1s22s22p6...and diagrams showing electrons in the same orbital have
opposite spin. For both ground and excited states.
Understand difference between valence and core electrons.
Chapter 7:
Covalent Bonds-why they form
Lewis structures, resonance
Be able to use formal charge to evaluate resonance contributors.
Understand why/how some compounds either do not fill their octets OR expand their octets.
Estimate bond strength/length based on Lewis structures.
VSEPR: understand the basis (electrons repel each other and want to maximize distance in between
Understand difference between electron pair geometry and molecular geometry
Be able to determine arrangement of both electron pairs (including multiple bonds) and atoms around
central atom and give both electron pair geometry name and name of molecular geometry.
Polar covalent bonds, dipole moments (def., arrows, how to predict if a MOLECULE is polar (sect 10.1)
using Lewis structures & VSEPR)
*Hybridization-be able to draw sp, sp2, and sp3 hybrid orbitals and show overlap in bond formation.
*MO Theory- how it works, what it predicts, diagrams, difference in bonding and antibonding orbitals.
Be able to determine and use bond order for diatomic molecules. Para-diamagnetic, relative
stability, bond strength, bond length….
Chapter 8:
PV work
State functions-concept employed in Hess’ Law and Born-Haber calculations.
*Calorimetry-coffee cup and bomb
*Hess’ Law (Be able to add up and cancel out reactions)
*Heats of formation (standard thermodynamic state)
*Calculations involving bond strength.
Chapter 9:
PV=nRT problems
Definition of absolute zero
Kinetic Molecular Theory (main points as discussed in class)
Chapter 10
Identify types of intermolecular forces and be able to draw H-bonds.
*Understand the effect intermolecular forces have on many physical properties and phases:
ex. bp, mp, Pvap, viscosity, surface tension.
*Heating curves and changes of state calculations. Be able to calculate the energy needed/released for
changing physical state or temperature.
*Understand the effect of temperature on vapor pressure and how to quantitatively describe it using the
Clausius -Claperon equation. This also allows you to find ΔH for a substance.
Know the definition of boiling point, melting point, enthalpy of vaporization, enthalpy of fusion, and
enthalpy of sublimation.
Chapter 11:
Be able to interconvert between different units of concentration.
Raoult’s Law/vapor pressure of solutions problems
Fractional distillation.
1. Which of these molecules contains polar bonds and which of these molecules are polar? Which
CCl4
would have the lowest boiling point?
water
PH3
2. A sample of a hydrocarbon, a compound that contains only carbon and hydrogen, was combusted
and produced 21.83 g CO2 and 4.47 g water and 311 kJ of heat.
(a) Is the reaction endo or exothermic?
(b) What was the mass of the combusted sample?
(c) Assuming the molecular formula of the hydrocarbon is the same as the empirical formula,
determine the heat of combustion for this compound in kJ/mol.
(d) What would the temperature change be if the heat produced in combusting 5.5 g of this
hydrocarbon was used to heat a beaker containing 666 g water originally at 3oC?
3. Given the following connectivity for the acetate anion, CH3CO2-, draw the most stable resonance
structure for it, show the formal charge at all atoms, and give the e- and molecular geometry at both
carbon atoms. What types of orbitals are used in bonding by the carbon atom? How many pi and
sigma bonds does this molecule contain? Draw the orbital overlap causing bond formation in the
acetate anion.
H3C
C
O
O
4. Without using any tables, which has a shorter bond length, F2 or O2?
5. Rank by increasing Pvap: BaSO4
SO2
N2O
CH3CH3
HF
6. A bottle of wine contains 12.5% ethanol (CH3CH2OH) by volume. If the density of ethanol is 0.789
g/mL, give the concentration of ethanol in wine in terms of mass percent and molarity.
7. Nitric oxide has an unpaired electron is on the nitrogen atom and a double bond between N and O. If
it requires 151 kJ/mol to break the N-O bond in nitric oxide, determine if the formation of nitric oxide
from nitrogen gas and oxygen gas would require or release energy. .N=O
8. Which is more likely to be an oxidizing agent, S or Na?
10. Predict ΔHrx for the reaction of ammonia, methane (CH4) and oxygen to give hydrocyanic acid and
water in kJ/mol CH4. Use any data tables you want to.
11. Determine ΔHvap for methanol, using the following vapor pressure data.
o
C
Pvap (mm Hg)
-6.0
20.0
5.0
40.0
12.1
21.2
60.0
100.0
12. Derive the equation for calculating the work done by an expanding gas.
13. How many grams of methane (ΔHcomb = -880 kJ/mol) would you have to burn to produce enough
energy to convert 55 kg ice at -15 oC to steam at 250 oC?
14. Given Pvap of water is 23.76 mm Hg at 25 o C, find the vapor pressure of a solution prepared by
dissolving 35.0g sodium sulfate in 175 g water at 25 oC.
15. a. Draw a Lewis structure for SO2, with formal charges, and give electron pair/molecular geometry.
b. Is the above molecule polar?
c. Draw the orbitals overlapping to form these bonds.
16. Give the electron configuration for Cr, N3-, and Ni2+.
17. Does the strength of imf’s change with temperature? Based on your answer, explain why viscosity
changes with temperature.?
18. In Denver, water boils at 95 oC. Why is this different than the standard bp for water. What is the
atmospheric pressure in Denver given ΔHvap = 40.67 kJ/mol for water.
19. 1,3-propanediol (HOCH2CH2CH2OH) can form intramolecular AND intermolecular hydrogen bonds.
Draw an example of both. Could dimethyl ether (CH3OCH3) form both kinds? Which substance
would have a lower freezing point?
20. What is the atomic number and expected electron configuration of the yet undiscovered element
directly below radon (Rn) in the periodic table.
c. 118, [Rn]7s25f146d107p6 d. none of these
a. 104, [Rn]7s27p6 b. 118, [Rn]7s27f147d107p6
21. One of the following orbital diagrams represents the ground state of Cr and another represents the
same atom in an excited state. The corresponding diagrams (ground state, excited state) are:
a. ↑↓ ↑ ↑ ↑
b. ↑
↑↓ ↑ ↑ ↑
c. ↑
↑ ↑ ↑ ↑ ↑ d. ↑↓
↑ ↑↑ ↑
4s
3d
4s
3d
4s
3d
4s
3d
22. a.Order the following molecules by increasing bond angles and briefly explain your reasoning:
water
PH3
tetrachloromethane (CCl4)
N2
b. Which of these molecules contains polar bonds?
c. Which of these molecules are polar?
d. Which of these compounds would have the lowest freezing point? The highest?
23. Why does dew form at night?
24. 300 mL O2 (measured at 25oC and 740 torr) and 400 mL H2(measured at 45 oC and 1250 torr) were
mixed in a 500 mL flask. The mixture was ignited and the gases reacted to form water. If the final
temperature of the flask was 120 oC, determine the pressure inside the flask.
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