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1. Review all vocabulary from all chapters.
2. Explain the origin of the Celsius, Fahrenheit, and Kelvin temperature scales and make
conversions between scales.
Ex. 20.0 o F = ____-6.67________ o C = ____266.33_______ K
Ex. 35.0 o C = ____95________o F = ____308_______K
(5/9) * (fahrenheit - 32)
(9/5) * celsius + 32
3. Make conversions in the SI system and between the SI system and the English system.
Ex. 356 m = _35600______cm = __356000____mm = _3.56 x 10-10_____Tm =
_3.56 x 1014_____pm
yotta- (Y-)
1024 1 septillion
zetta- (Z-)
1021 1 sextillion
exa- (E-)
1018 1 quintillion
peta- (P-)
1015 1 quadrillion
tera- (T-)
1012 1 trillion
giga- (G-)
1 billion
mega- (M-)
1 million
kilo- (k-)
1 thousand
hecto- (h-)
1 hundred
deka- (da-)** 10
1 ten
deci- (d-)
1 tenth
centi- (c-)
1 hundredth
milli- (m-)
1 thousandth
micro- (µ-)
1 millionth
nano- (n-)
1 billionth
pico- (p-)
10-12 1 trillionth
femto- (f-)
10-15 1 quadrillionth
atto- (a-)
10-18 1 quintillionth
zepto- (z-)
10-21 1 sextillionth
yocto- (y-)
10-24 1 septillionth
4. Determine the number of significant figures a number has and use the correct number
of significant figures when determining answers to calculations.
Ex. .00908 has ___3__sig figures
23000 has ___2___ sig. figures
2.02 x .0023 = _______
1.001 + 6.1 + 2.5678 = ______
5. Express numbers in scientific notation, and convert numbers from scientific notation to
standard nonexponential form.
Ex. Express the following in scientific notation.
5.08 x 10-7
1.23 x 108
Ex. Express the following in standard nonexponential form.
1.23 x 105 123000
4.56 x 10-4 .000456
6. Make calculations (addition, subtraction, multiplication, division) when numbers are
in scientific notation.
Ex. (2.0 x 104) (3.0 x 10-2) 6.0 x 102
7. Perform calculations involving density.
Ex. A solid has a mass of 54.0 g and volume of 35.0 g/cm3. Calculate its density . (1.54
g/cm3; 1.54)
8. Calculate the molar mass of any compound using the periodic table.
Ex. HCl, NH4OH, CS2
9. Perform conversions similar to the following.
Ex. 354 g of CH4 = _______moles
Ex. 6.54 moles of NaOH = ____________ grams
Ex. 4.00 molecules of CCl4 = ____________ atoms of Cl
Ex. 6.54 moles of NH3 = ___________molecules
Ex. 5.00 molecules of P = _______________grams of P
(3.94 x 1024)
(2.57 x 10-22)
10. Use data obtained from quantitative analysis for the composition of a compound or
from percentage composition to determine the compound’s empirical formula.
Ex. A compound with 0.90 g Ca and 1.6 g Cl has what empirical formula?
Ex. A white powder used in paints, enamels and ceramics has the following percentage
composition: Ba 69.6 %; C 6.09%; O 24.3 %. What is its empirical formula? (BaCO3)
11. Write and balance chemical equations (other than redox) and identify the reaction
type as composition, decomposition, single replacement, or double replacement.
Ex. potassium hydroxide + sulfuric acid ------> potassium sulfate + water
2KOH + H2SO4 K2SO4 + 2H2O Double replacement
12. Identify the limiting reagent in a reaction.
Ex. How many grams of sodium nitrate can be formed when 10.0 g of sodium hydroxide
is reacted with 20.0 g of nitric acid? Which reactant limits the reaction. ( NaOH, 15.8 g
of nitric acid is needed) (21.3 g of sodium nitrate produced; 4.2 g of nitric acid in
13. Define molarity and work problems involving the molarity of a solution.
Ex. What is the molarity of a solution formed by dissolving 5.00 g of NaOH in 75.0 mL
of water? (1.67 M)
14. Perform stoichiometric calculations with balanced equations when the reactants are
solutions of know molarity.
Ex. How many mL of 0.112 M HCl will react with the sodium carbonate in 21.2 mL of
0.150 M Na2CO3?
(56.8 mL)
15. Define and be able to identify acids and bases from Arrhenius’, Bronsted - Lowrey’s
and Lewis’ point of view.
Ex. HC2H3O2 + H2O  H3O+ + C2H3O2-1
Identify the Bronsted acids and bases and the conjugate acid - base pairs in the
above equation
B. acid = HC2H3O2 conj. Base= C2H3O2-1 B Base = H2O conj. Acid = H3O+
17. Predict if metathesis occurs, explain why it occurs, and write ionic and net ionic
equations for the reaction.
Ex. Predict if the following reaction occurs. If so, explain why and write ionic
and net ionic equations for the reaction. H2SO4 + 2 KOH  K2SO4 + 2 H2O
Yes because water is a product Net Ionic 2H+(aq) + 2OH-(aq)  2H2O(l)
18. Work stoichiometric problems using ionic equations.
Ex. How many mL of 0.100 M AgNO3 solution are needed to react completely
with 25.0 mL of 0.400 M CaCl2 solution? ( Ag+ + Cl-  AgCl)
25.0mL x 0.400MCaCl2 x 2mol Clx 1mol Ag+ x 1mol AgNO3 x 1000ml
1 mol CaCl2 1 mol Cl- 1mol Ag+
0.100M AgNO3
= 200mL AgNO3
19. Work problems involving titration.
Ex. 25.00 mL of HCl are titrated with 37.46 mL of 0.0775 M NaOH. What is the
molarity of the HCl solution?
20. Be able to balance redox equations using the half reaction method.
Ex. Cu + HNO3  Cu(NO3)2 + NO + H2O
21. Be able to balance redox equations for basic or acidic reactions using the ion electron method.
Ex. Cr2O7-2 + Fe+2  Cr+3 + Fe+3
(acid solution)
Ex. SO3-2 + MnO4-1  SO4-2 + MnO2
(basic solution)
22. Using specific heat and heat capacity data as well as temperature changes that occur
in a calorimeter, calculate the heat of a reaction.
Ex. In a calorimeter containing 100 g of water, a reaction caused the temperature to rise
15.0 oC. How many Joules were given off? Convert the value to calories. Was the
reaction endothermic or exothermic? (6270 J; 1499 cal)
Ex. See sample problem on p. 279
23. Manipulate thermochemical equations so as to use Hess’s Law to find the value of
Ho for a reaction for which  H o might be hard or impossible to measure directly.
Ex. See sample problems on p. 286
24. Use standard heats of formation to find the  Ho of a reaction.
Ex. See sample problems on p. 290