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Chem 11
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Block:
FINAL EXAM REVIEW SHEET
CHAPTER 1
1. Name the following units:
a. mm
c. cg
millimetres
centigrams
b. ug micrograms
d. dm decimetres
2. Complete the following conversions
a. 4700 g = 4.7 kg
d. 1000 mm = 1 m
b. 0.000000146 km = 146um
e. 0.1 cm = 1 mm
g. 5.280 m = 5280 mm
h. 4530 mm = 453 cm
3. What is the difference between precision and accuracy?
Precision relates to the number of decimal places, which accuracy is how close
you are to the accepted value.
CHAPTER 2
1. Measuring mass or volume of an object will not reveal what the object is. How can
you use mass and volume to produce a quantity that would describe a characteristic
property? mass divided by volume gives density. Density is a quantitative
property and is always the same for a particular pure substance.
2. Give an example of the Law of Multiple Proportions. Carbon and oxygen can
combine in different ways. CO, CO2,CO3
3. Is the burning of charcoal a physical or chemical change? Chemical change;
anytime you burn anything it is a chemical change
4. What are three different ways to separate mixtures or compounds? Filtration can
separate undissolved particles from a suspension. Distillation/Evaporation can
separate dissolved particle from a solution. Hand separation, or a sieve can
separate large particles from small particles.
5. Name these compounds
a. Cu3(PO4)2
b. H2SO3
c. FeSO3
d. KHCO3
copper II phosphate
sulphurous acid
iron II sulphite
potassium bicarbonate
e. Ca(NO3)2
f. N2O3
g. CCl4
h. Ni3P2
calcium nitrate
dinitrogen trioxide
carbon tetrachloride
nickel II phosphide
6. Write a formula for the following compounds.
a. lead (II) iodide
b. sodium hydroxide
c. ammonium carbonate
d. phosphorus pentachloride
PbI2
NaOH
(NH4)2CO3
PCl5
e. copper (II) chromate
f. tin (IV) nitride
g. magnesium nitrate
h. calcium hydroxide
CuCrO4
Sn3N4
Mg(NO3)2
Ca(OH)2
CHAPTER 3
1. How many sig figs?
a. 200
b. 1.28467
c. 1.500
1
6
4
d. 1500
e. 0.00015
f. 15.00
2
2
4
2. Round each answer to the correct number of sig figs.
a. 2.86 x 1.824 = 5.22
b. 21/ 8.0 = 2.6
c. 3.7 + 1.86 + 0.0024 = 5.6
d. 98.9 x 1.22 = 121
e. 2.1 + 1.0 = 3.1
f. 65 – 3.27 = 62
3. A millimetre rulers measures to what number of decimal places if your answer is in
centimetres? 2 decimal places: centimetres is whole number, millimetre line is first
decimal, second decimal is uncertain digit between millimetres lines.
CHAPTER 4
1. How many atoms of each element in:
a. 5 molecules of C5H12Cl4 C = 25, H = 60, Cl = 20
b. 10 molecules of Co(ClO4)2 ● 6H2O
Co = 10, Cl = 20, O = 140, H = 120
2. What is mass of:
a. 1.800 moles of Ag
194.2 grams
b. 1.5 x 1015 molecules of NaOH
1.0 x 10-8 grams
3. How many moles are in:
a. 5.00 g of methane
b. 4.55 atoms of Na
c. 527 mg of KNO3
0.312 moles
7.56 x 10-24 moles
0.00521 moles
4. What is the molar mass of:
a. sodium phosphate Na3PO4 = 164.0 grams/mol
b. copper (III) sulphate
Cu(SO4)3 = 351.8 grams/mol
5. How many oxygen atoms are in 2.75 moles of oxygen gas?
3.31 x 1024
6. A compound of chlorine and oxygen is found to contain 2.5 moles of chlorine and 4.9
moles of oxygen. What is the empirical formula of this compound? ClO2
7. A compound is found to contain 168g of iron, 128g of oxygen and 62g of phosphorus.
What is the empirical formula of the compound? Fe3(PO4)2
8. A compound has the empirical formula of CH2 and a molecular mass of 42g/mol.
Determine its molecular formula.
C3H6
9. A 9.2g sample of a compound is 2.8g nitrogen and 6.4g oxygen. Find the empirical
formula of the compound. NO2
10. A particle of freon contains 9.93% carbon, 56% chlorine, and 31.4% fluorine by
mass. What is the empirical formula? CCl2F2
Extra: what would you call the compound from question 10? Dichloro-difluoromethane
11. A compound consists of 40% calcium, 12% carbon, and 48% oxygen by mass, what
is its empirical formula?
CaCO3
12. You are asked to make a solution of H2SO4.
a. What mass of H2SO4 would you need to prepare 250.0 mL of 13.50 M H2SO4?
331.1 grams
b. What are two methods to make the solution more concentrated?
Add more solute (H2SO4), or evaporate some of the water
c. What volume of your original solution would you need to make 250.0 mL of 2.5M
H2SO4?
0.046 L
13. You have two solutions of NaOH. Solution A is 50 mL of 3.0 M NaOH and Solution
B is 75 mL of 0.50 M NaOH. You mix together solution A and B. What is the molar
concentration of the resulting solution?
1.5 M
14. If you start with 500.0 mL of 0.50 M NaOH, describe how would you increase the
concentration to 3.0M NaOH, without adding more solute? Evaporate the water until
there was only 83mL left.
15 What mass of KBr is contained in 500.0mL of 0.235 M KBr?
14.0 grams
16. What is the concentration of HNO3 when 5.00 mL of 3.5 M HNO3 is mixed with 95.0
mL of 0.20 M HNO3?
0.36 M
CHAPTER 5
1. How can you tell a reaction is exothermic?
is given off, ∆H is Negative
Temp goes up, it feels warm, heat
2. On the following reaction, indicate: coefficients, subscripts, reactants, products.
2HCl + Zn → H2
+ ZnCl2
Big 2 = coefficient, Little 2 = subscript, HCl and Zn = reactants, H2
products
+ ZnCl2 =
3. Balance the following equations.
a. _____Ca(OH)2(s) + __2___HCl(aq) ---> ____CaCl2(aq) + ___2__H2O(l)
b. __2___FeCl3(aq) + ___3__(NH4)2S(aq) ---> _____Fe2S3(s) + __6___NH4Cl(aq)
c. __2___KNO3 (s) ---> __2___KNO2 (s) + _____O2(g)
d. ___2__Ag2O(s) ---> __4___Ag(s) + _____O2(g)
e. __2___C4H10(g) + ___13__O2(g) ---> __8___CO2(g) + ___10__H2O(g)
4. For each equation above, state the type of reaction.
a. neutralization
b. double replacement
c. decomposition
d. decomposition
e. combustion
4. Write out the following word equations as chemical equations.
a. potassium chlorate ---> potassium chloride + oxygen
KClO3 → KCl + O2
b. magnesium + oxygen ---> magnesium oxide
Mg + O2 → MgO
c. magnesium oxide + water ---> magnesium hydroxide
MgO + H2O → Mg(OH)2
d. aluminum + sulphuric acid ---> aluminum sulphate + hydrogen
Al + H2SO4 → Al2(SO4)3 + H2
5. Predict the products and balance the following equations.
a. _____Br2(aq) + ___2__KI(aq) ---> 2KBr
+ I2
b. ___2__C5H12O (l) + __15___O2(g) --> 10CO2 + 12H2O
c. _____Zn(s) + _____H2SO4 (aq) ---> ZnSO4 + H2
d. ___2__Ag(s) + _____Cl2(g) ---> 2AgCl
CHAPTER 6
1. In the reaction 3 Fe + 4H2O -----
Fe2O3 +
4H2
a. How many molecules of Fe2O3 are produced when 12 atoms of Fe react?
4
b. How many H2 molecules are made when 40 molecules of Fe2O3 are produced?
2. Consider the reaction
160
4NH3 + 5O2 ------- 6H2O + 4NO
a. what mass of NO is produced when 2.00 mol of NH3 are reacted with O2? 60g
b. What mass of H2O is produced when 4.00 mol of O2 are reacted with NH3?
86.4g
3. Pentane, C5H12, burns according to the reaction :
C5H12 + 8O2 ------ 5CO2 + 6H2O
a. What mass of CO2 is produced when 100.0g of C5H12 is burned? 305.6g
b. What mass of O2 is required to produce 60.0g of H2O? 142.2g
4. What volume of 0.250 M HCl is required to completely neutralize 25.0mL of 0.318 M
NaOH? (Hint: write a balanced equation for the reaction between NaOH and HCl)
31.8mL
5. What volume of 0.200 M KOH is required to react with 125 mL of 0.250 M H3PO4 in
order to produce a solution of K2HPO4? 0.312 L
6. If 19.8 mL of H3PO4 with an unknown molarity react with 25.0 mL of 0.500 M KOH,
according to the following reaction:
H3PO4 + 2KOH -------- K2HPO4 + 2H2O
a. What is the molarity of the H3PO4? 0.316M
7. 2 C8H18 + 25 O2 ---------------> 16 CO2 + 18 H2O
(a) If 1.50 mol of oxygen gas is combined with 0.50 mole of octane, which reactant is in
excess?
OCTANE
(b) How many moles of excess? 0.38 MOLES
(c) What is the mass of the excess? 43 grams
(d) How many grams of carbon dioxide will be produced? 42 grams
8. A mixture is prepared by adding 20.0 mL of 0.200 M Na3PO4 to 30.0 mL of 0.150 M
Ca(NO3)2. What weight of Ca3(PO4)2 will be formed? 0.465g
9. The average human requires 120.0 grams of glucose (C6H12O6) per day. How many
grams of CO2 (in the photosynthesis reaction) are required for this amount of glucose?
The photosynthetic reaction is: 6 CO2 + 6 H2O ---> C6H12O6 + 6 O2 176.9g
10. According to the following reaction:
4NH3 + 5O2 ----- 6H2O + 4NO
What volume of 0.200 M NH3 is required to produce 45.3 g of H2O if the reaction has a
% yield of 89.0%? 9.45L
CHAPTER 8
1. Complete the following table :
Atom/Ion
39 +1
K
40
Ca
70
Ga+3
35 -1
Cl
16
O
238 +4
U
Atomic #
19
20
31
17
8
92
Mass #
39
40
70
35
16
238
Protons
19
20
31
17
8
92
Neutrons
20
20
39
18
8
146
Electrons
18
20
28
18
8
88
2. Naturally occurring silver consists of 51.8% Ag-107 and 48.2% Ag-109. What is
the expected average molar mass of a sample of natural silver, expressed to 3
decimal places?
107.964g/mol
CHAPTER 10
1. What are the electron configurations for:
a)P b) Cl
c) Ar
d) F-1 e) S-2
2
2
6
2
3
a. 1s 2s 2p 3s 3p
b. 1s2 2s2 2p6 3s2 3p5
c. 1s2 2s2 2p6 3s2 3p6
d. 1s2 2s2 2p6
e. 1s2 2s2 2p6 3s2 3p6
f. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
f) Kr
2. How many valence electrons are there in:
a)Si b) K+1 c)Ne
d)O-2 e)H f)N
a. 4
b. 0
c. 8
d. 8
e. 1
f. 5
CHAPTER 11
1. Which atom is bigger, Bi or P? Why? Bi. Atomic radius increases as you do
down a group because each row down adds another orbital
2. Name at least one trend among each of the following periodic families:
a. Alkali Metals – most reactive metals, react violently with water
b. Alkaline Earth Metals – reactive metals but less so than alkali
c. Halogens – most reactive non-metals, most are poisonous
gases
d. Noble Gases – do not react
3. If a noble gas could form a +1 ion, which of the noble gases would form a +1 ion
most easily? Rn because it is the largest so there is a lot of electron
shielding on the outermost electrons, making it easier to lose one electron
and gain a +1 charge.
CHAPTER 12
1. With respect to electrons, how does an ionic bond differ from a covalent bond?
The pull on the shared electron(s) in an ionic bond is so great that the nonmetal pulls the shared electron all the way over to itself, making a negative
ion and leaving the metal as a positive ion
2. Indicate whether the following compounds are ionic, covalent, or polar covalent.
Explain.
a)NaCl
b)H2O
c)NO2
d)CS2
a. ionic ΔEN = 2.1
b. polar covalent ΔEN = 1.4
c. polar covalent ΔEN = 0.4
d. covalent ΔEN = 0.0
Draw electron dot structures (Lewis Structures) for:
a)C2H4
b)H2
c)PO4-3
d)HCN
e)NH4+1
3. What is the geometric shape of the following molecules?
a. NH3
b. HCN
c. PCl5
d. CH4
a. Trigonal planar
b. Linear
c. Trigonal bipyramidal
d. tetrahedral
CHAPTER 16
1. Which would be miscible with water: ethanol or ethane? Why?
Ethanol because it is polar so the water molecules would be attracted to
the polar end of the ethanol molecule.
Ethane would not mix because it is non-polar so the water molecules are
not attracted to it and they would just stay together.
2. Why are some solvents polar and some non-polar? Which would you use to
dissolve salt? Polarity depends on the electronegativity difference between
the atoms in a bond and the configuration of the molecule. If a molecule
has evenly distributed polar bonds, the polarity cancels and the molecule
itself may be non-polar. To dissolve salt, you would use the polar solvent,
because ionic compounds have charged components that are soluble in
polar solvents
3. What will be the [Cl-1] if 60.0g of BaCl2(s) is dissolved in 600.0mL of H2O? 0.964M
4. If 35.0g of V(NO3)5 is dissolved in 1.0L of H2O, what will be the [NO3-1] be?
0.485M
5. What mass of NaBr must be dissolved in 4.9L to make a 5.2M NaBr solution?
2.6kg
6. If 700.0mL of 1.9M KCl is mixed with 1.05L of 0.750M MgCl2, what will the final
[Cl-1] be? 1.66M
7. 100.0mL of 0.5 M Mg(OH)2 is titrated with 45.0mL of H3PO4. What is the
concentration of the H3PO4? 0.74M
Sorry no CH 23 review. You just did the test! You are experts at organic
chemistry…right?