This document provides a review for a Chem 105 final exam covering various topics in chemistry. It begins with a disclaimer about the exam problems coming from the textbook and suggestions for additional practice. It then reviews measurements, nomenclature, stoichiometry, reactions such as redox and acid-base, thermochemistry, electrochemistry, quantum mechanics, periodic trends, and various states of matter. Practice problems are provided throughout covering calculations and identifying concepts in several chemistry areas in preparation for the final exam.

1. Chem 105 Final Review
DR. SHIRTS-WINTER 2013
JOANNA WILLIAMS

2. Disclaimer
 All the problems on here come from your textbook! If you need more, I
suggest working through micro-exams and practice sheets/parallel
example problems, seeing your Learning Community Mentor to run
through previous exams, and looking in the book for more problems to
work through over specific things your struggling with.

3. Measurements
 Accurate: numbers close to the actual value
 Precise: numbers close to each other
 Significant Figures:
 All non-zeros are sig. figs.
 Zeros between two non-zeros are sig. figs.
 Zeros left of first non-zero are NOT sig. figs.
 If #>or=1, all zeros right of decimal are sig. figs.
 If #<1, all zeros at end of # and between non-zeros are sig. figs.
 Trailing zeros may or may not be sig. figs. (That’s why we use scientific notation)

4. Nomenclature
 Metal + Nonmetal = Ionic compound
 Charges designate formula, name the elements and add –ide to the end
 Nonmetal + Nonmetal = Covalent molecule
 Use prefixes and add –ide to the end
 Polyatomic Ions
 Organic Functional Groups
 Organic Prefixes
 Acids
 If you “–ate” too much you feel “–ic”ky
 “-ite”s like Nephites and Lamanites are people like “-ous”
 Hypo-ous, ous, ic, per-ic  increasing O
 Hydro-ic

5. Dimensional Analysis
 Use the Mole to Mole Ratio from stoichiometric coefficients in balanced
chemical equation
 Find limiting reactant

6. Practice Problems
 3.51) Determine the empirical and molecular formulas of each of the
following substances:
 Styrene, a compound used to make Styrofoam cups and insulation, contains
92.3% C and 7.7% H by mass and has a molar mass of 104 g/mol
 Caffeine, a stimulant found in coffee, contains 49.5% C, 5.15% H, 28.9% N, and
16.5% O by mass and has a molar mass of 195 g/mol
 Monosodium glutamate (MSG), a flavor enhancer in certain foods, contains
35.51% C, 4.77% H, 37.85% O, 8,29% N, and 13.60% Na, and has a molar mass of
169 g/mol

7. Practice Problems
 3.69) A piece of aluminum foil 1.00 cm square and 0.550 mm thick is
allowed to react with bromine to form aluminum bromide.
 How many moles of aluminum were used? (density of aluminum=2.699 g/mL)
 How many grams of aluminum bromide form, assuming the aluminum reacts
completely?

8. Practice Problems
 3.76) Aluminum hydroxide reacts with sulfuric acid as follows:
2Al(OH)3(s) + 3H2SO4(aq)  Al2(SO4)3(aq) + 6H2O(l)
Which is the limiting reactant when 0.500 mol Al(OH)3 and 0.500 mol
H2SO4 are allowed to react? How many moles of Al2(SO4)3 can form
under these conditions? How many moles of the excess reactant
remain after the completion of the reaction?

9. Reactions
 Know the solubility rules and the exceptions for precipitation reactions
 Oxidation-Reduction (RedOx)
 123FHO7654
 LEO goes GER / OIL RIG
 Oxidizing agents/Reducing agents

10. Reactions
 Acid/Base/Neutralization
 Titrations and Dilutions: M1V1=M2V2 (molarity= mol/L)(volume=L)
 Net Ionic Equations
 Strong Acids (ionize completely)
 H2SO4 , HNO3, HCl, HBr, HI, HClO4
 Strong Bases (dissociate completely)
 Group 1-OH, Group 2-OH from Ca down

11. pH
 pH=-log[H+]
 pOH=-log[OH-]
 pH+pOH=14
 [H+]+[OH-]=1x1014
 pH<7 acidic
 pH=7 neutral
 ph>7 basic

12. Practice Problems
 4.83) Some sulfuric acid is spilled on a lab bench. You can neutralize the
acid by sprinkling sodium bicarbonate on it and then mopping up the
resultant solution. The sodium bicarbonate reacts with sulfuric acid as
follows:
2NaHCO3(s) + H2SO4(aq)  Na2SO4(aq) + 2H2O(l) + 2 CO2(g)
Sodium bicarbonate is added until the fizzing due to the formation of CO2(g) stops.
If 27 mL of 6.0 M H2SO4 was spilled, what is the minimum mass of NaHCO3 that
must be added to the spill to neutralize the acid?

13. Practice Problems
 4.40) Write the balanced molecular and net ionic equations for each of
the following neutralization reactions:
 Aqueous acetic acid in neutralized by aqueous barium hydroxide
 Solid chromium(III) hydroxide reacts with nitrous acid
 Aqueous nitric acid and aqueous ammonia react

14. Practice Problems
 4.51) Which element is oxidized and which is reduced in the following
reactions?
 N2(g) + 3H2(g)  2NH3(g)
 3Fe(NO3)2(aq) + 2Al(s)  3Fe(s) + 2 Al(NO3)3(aq)
 Cl2(aq) + 2NaI(aq)  I2(aq) + 2NaCl(aq)
 PbS(s) + 4H2O2(aq)  PbSO4(s) + 4H2O(l)

15. Thermochemistry
 Ek = ½(mv2)
 1 Cal = 1000 cal
 1 cal = 4.184 J (specific heat of water)
 ∆E=q+w
 H=E+PV
 Bond enthalpies: reactants – products aka bonds broken – bonds formed
 ∆Hf
o: products – reactants (diatomics in natural state = 0)
 q=mCs∆t
 Hess’s Law

16. Practice Problems
 5.43) Consider the following reaction:
2Mg(s) + O2(g)  2MgO(s) ∆H = -1204 kJ
 Is this reaction exothermic or endothermic?
 Calculate the amount of heat transferred when 3.55g of Mg(s) reacts at
constant pressure
 How many grams of MgO are produced during an enthalpy change of -234 kJ?
 How many kilojoules of heat are absorbed when 40.3g of MgO(s) is
decomposed into Mg(s) and O2(g) at constant pressure?

17. Practice Problems
 5.56) When a 4.25g sample of solid ammonium nitrate dissolves in 60.0g of
water in a coffee-cup calorimeter, the temperature drops from 22.0 C to
16.9 C. Calculate ∆H (in kJ/mol NH4NO3) for the solution process
NH4NO3(s)  NH4+(aq) + NO3-(aq)
Assume that the specific heat of the solution is the same as that of pure water.
Is this process endothermic or exothermic?

18. Practice Problems
 5.65) From the enthalpies of reaction
H2(g) + F2(g)  2HF(g) ∆H = -537 kJ
C(s) + 2F2(g)  CF4(g) ∆H = -680 kJ
2C(s) + 2H2(g)  C2H4(g) ∆H = +52.3 kJ
Calculate ∆H for the reaction of ethylene with F2:
C2H4(g) + 6 F2(g)  2CF4(g) + 4HF(g)

19. Electrochemistry
 E=hv=hc/λ
 1/λ=R(1/n1
2-1/n2
2)
 E=-Rhc(1/n2)
 λ=h/mv
 Bohr’s Model

20. Practice Problem
 6.37) Calculate the energy of an electron in the hydrogen atom when n=2
and when n=6. Calculate the wavelength of the radiation released when
an electron moves from n=6 to n=2.
 Is this line in the visible region of the electromagnetic spectrum? If so, what
color is it?

21. Orbitals & Nodes
 s=spherical
 Radial nodes starting at 2s
 p=peanut
 1 planar node and radial nodes starting at 3p
 d=dlover leaf?
 2 planar nodes and radial nodes starting at 4d
 f

22. Quantum Numbers
 Pauli Exclusion Principle
 n=shell (1, 2, 3….)
 l=subshell (n-1 to 0) (0=s, 1=p, 2=d, 3=f…)
 ml=orientation (-l to l)
 ms=spin (-1/2 or +1/2)

23. Practice Problem
 6.56) Which orbital goes with the following quantum numbers? Which are not
allowed?
 2, 1, -1
 1, 0, 0
 3, -3, 2
 3, 2, -2
 2, 0, -1
 0, 0, 0
 4, 2, 1
 5, 3, 0

24. Electron Configuration
 Expanded
 Condensed using Noble Gas configuration
 Cu and Cr exceptions
 Why? Hund’s Rule

25. Practice Problem
 6.61) For a given value of the principal quantum number, n, how do the
energies of the s, p, d, and f subshells vary for
 Hydrogen?
 A many-electron atom?

26. Periodic Trends
 Electronegativity
 Size
 Ionization Energy
 Electron Affinity
 Effective Nuclear Charge
 Family Names
 Cation and Anion Size

27. Lewis Dot Structures
 Count total valence electrons
 Least electronegative atom in the middle
 Fill octet, create multiple bonds if too many electrons
 Resonance structures: none actually what the molecule looks like, it’s a
hybrid of all of them
 Formal charges
 Bond strengths

28. Molecular Orbitals
 Bond Order
 MO Diagrams
 Paramagnetic vs. Diamagnetic
 VSPER
 Molecular shapes and Angles

29. Gases
 PV=nRT
 22.4 L = 1 mole gas @ STP
 Ideal gas characteristics
 Small, high temp, low pressure
 Partial pressures
 Pa=XaPt
 Pt=P1+P2+P3…
 Effusion and Diffusion Rates
 Urms=√(3RT/M)
 r1/r2= √(M2/M1)
 J=kgm2/s2

30. Practice Problems
 10.54) Calculate the density of sulfur hexfluoride gas at 707 torr and 21 C
 Calculate the molar mass of a vapor that has a density of 7.135 g/L at 12 C
and 743 torr

31. Practice Problem
 10.69) A piece of dry ice (solid carbon dioxide) with a mass of 5.50 g is
placed in a 10.0 L vessel that already contains air at 705 torr and 24 C.
After the carbon dioxide has totally vaporized, what is the partial pressure
of carbon dioxide and the total pressure in the container at 24 C?

32. Intermolecular Forces
 London Dispersion
 Induced Dipole (Polarizability)
 Dipole-Dipole
 Hydrogen Bonding
 Ion-Dipole

33. Liquids
 Intermolecular force effect on
 Viscosity
 Surface Tension

34. Phase Changes
 Phase Diagrams
 Specific heats to heat the substance to melting or boiling point
 Heats of vaporization or fusion to melt or evaporate substance

35. Practice Problems
 11.43) For many years drinking water has been cooled in hot climates by
evaporating it from the surfaces of canvas bags or porous clay pots. How
many grams of water can be cooled from 35 C to 20 C by the
evaporation of 60 g of water?
(The heat of vaporization of water in this temperature range is 2.4 kJ/g.
The specific heat of water is 4.18 J/gK.)

36. Colligative Properties
 Depends on number of solute particles present, not identity
 Don’t forget ions dissociate!
 Vapor Pressure ↓
 Pa=XaP°
 Boiling Point ↑
 Freezing Point ↓
 Osmotic Pressure ↑
 π=iMRT

37. Solids
 Unit Cells
 Simple/Primitive Cubic
 1 atom
 Face-Centered Cubic
 4 atoms
 Body-Centered Cubic
 2 atoms

38. Concentrations
 [ ]=M=moles solute/L solution=molarity
 m=moles solute/kg solvent=molality
 X=moles substance/total moles=mole fraction
 ppm=(mass substance/total mass)x106
 mass %=(mass substance/total mass)x100

39. Equilibrium
 kc=[products]/[reactants]
 kp=Pproducts/Preactants
 aA + bB  cC + dD
 kc=[C]c[D]d/[A]a[B]b
 Le Chatelier’s Principle
 Noble gases being pumped in to increase the pressure have no effect (don’t
change individual partial pressures)
 Catalysts have no effect
 Solids and liquids have no effect