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AP Topic 1 Review
1. Classify the following as an element, a compound or a mixture, if you classify it as a mixture
state whether it is a homogeneous or heterogeneous mixture.
i)
Helium
ii)
Pure water
iii)
Pure table salt (sodium chloride)
iv)
Air
v)
Fruit cake
2. Classify the following as either chemical or physical changes.
i)
Ice melting
ii)
Gasoline burning
3. Mercury is a liquid metal that has a density of 13.58 g/mL. Calculate the volume of mercury
that must be poured out in order to obtain 0.5g of mercury.
4. Classify the following as either quantitative or qualitative observations.
i)
My eyes are brown
ii)
My neck size is 17 inches
iii)
My average grade last year was 79%
7. Round the following numbers to four figures.
i)
2.16347 x 105
ii)
4.000574 x 106
iii)
3.682417
iv)
7.2518
v)
375.6523
vi)
21.860051
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8. Perform the following conversions.
i)
0.75 kg to milligrams
ii)
1500 millimeters to km
9. Complete the following table of average temperatures on the planets.
Kelvin
Mercury
439
Venus
729
Fahrenheit
Celsius
Earth
15
AP Topic 2 Review
1. Using the periodic table, complete the following table. (13)
Isotope Symbol
Atomic #
Mass #
38
88
# Protons
# Neutrons
35
44
23Na
11
2. What is the charge on a sodium atom?
__________
3. What is the charge on a sodium nucleus?
__________
4. What is the atomic number of potassium?
__________
5. How many protons in the nucleus of a potassium atom?
__________
6. How many electrons in the potassium nucleus?
__________
7. Define the term isotope.
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# Electrons
8. Using the periodic table,
a) Complete the following table.
Isotope Symbol
Atomic #
# Protons
# Neutrons
Mass #
13C
17
18
26
56
17
37
2
52
3
128
50
70
Naturally occurring lead is found to have the following isotopic relative abundance. 204Pb 3%,
206Pb 24%, 207Pb 20% and 208Pb 53%. Calculate the average relative atomic mass of Pb from the
data.
The results taken from a mass spectrum of chlorine gas show peaks at m/e 35.00 and m/e 37.00
(The mass spectrum identifies the different isotopes of an element that are present in a sample).
a) Given that the relative abundances of Cl 35.00 and Cl 37.00 are 77.50% and 22.50%
respectively, calculate the average relative atomic mass of chlorine to four significant
figures.
Give the name of each of the following binary compounds.
a) KCl
b) BaO
c) Rb2S
d) Na3P
e) AlF3
f)
Mg3N2
g) CaI2
h) RaCl2
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Write the name of the following ionic substances using the system that includes a Roman
numeral to specify the charge on the cation.
a) FeI3
b) MnCl2
c) HgO
d) Cu2O
e) CuO
f)
SnBr4
g) MnO2
Write the name of each of the following binary compounds of non-metals.
a) N2Br4
b) P2S5
c) SeO2
d) N2O5
e) SiO2
Write the name of each of the following acids.
a) HCl
b) HNO2
c) H2SO4
d) HF
e) HI
f)
HC2H3O2
AP Topic 3 Review
1. Give full and abbreviated (noble gas core) electronic configurations for the following.
a) Cr
b) Fe
c) S2-
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FULL
______________________________
NOBLE GAS CORE
______________________________
FULL
______________________________
NOBLE GAS CORE
______________________________
FULL
______________________________
NOBLE GAS CORE
______________________________
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2. For each of the following sets of orbitals, indicate which orbital is higher in energy.
a) 1s, 2s
_________
b) 2p, 3p
_________
c) 4s, 3dyz
_________
d) 3px, 3py, 3pz
_________
3. Indicate the block (s, p or d) in which each of the following elements found.
BLOCK
a) Sc
______
b) P
______
c) Fr
______
d) Ni
______
e) As
______
4. Identify the element from the electron configurations of atoms shown below.
a) [Ne] 3s2 3p2
_________
b) [Ar] 4s2 3d7
_________
c) [Xe] 6s2
_________
5. State which atom or ion is represented by the following sets of atomic numbers and
electronic configurations.
Atomic #
Electronic Configuration
a)
8
1s2 2s2 2p4
________
b)
11
1s2 2s2 2p6
________
c)
14
1s2 2s2 2p6 3s2 3p2
________
d)
22
1s2 2s2 2p6 3s2 3p6 3d2
________
6. Give the electron configurations for the following transition metal ions.
a)
Sc3+
__________________________________
b)
Cr2+
__________________________________
c)
Ni3+
__________________________________
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7. Identify two positive and two negative ions that are isoelectronic with argon.
Two Positive ions
________
________
Two Negative ions
________
________
8. Using the electrons in boxes notation complete the electronic configurations of the
following elements.
1s
2s
2p
3s
3p
3d
4s
4p
Element
 
 
 V
1s
3s
4s
2s
2p
3p
3d
4p
Element
 
 
 Ar
1s
3s
4s
2s
2p
 
3p
3d
4p
Element
 Zn
 
9. State the number of unpaired electrons in each of the electronic configurations in
question 9
# of unpaired electrons
V
________
Ar
________
Zn
________
10. Write three possible sets of quantum numbers for the highest energy electrons in the
aluminum atom.
n
l
ml
ms
Electron #
11
Electron # 12
Electron # 13
11. Which atomic theory is violated by the following sets of quantum numbers
representing beryllium’s outer shell electrons? Explain your answer.
n
2
2
L
0
0
ml
0
0
ms
+½
+½
_____________________________________________________________________________
_____________________________________________________________________________
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AP Topic 4 Review

TYPE 1: Those involving Avogadro’s number.
Question 1
A sample of Ag is found to contain 9.7 x 1023 atoms Ag. How many moles of Ag atoms are in
the sample?
Question 2
How many Sb atoms are found in 0.43 moles of pure Sb?

TYPE 2: Those involving the relationship between mass, moles and molar mass.
Question 3
What is the mass in grams of 2.53 moles Al?
Question 4
How many moles of Na in 20g of Na?
Question 5
If 50 moles of a simple, binary, group I chloride have a mass of 3725g identify the group I
metal.

TYPE 3: Those combining types 1 & 2.
Question 6
How many Zr atoms are found in a 1.23g sample of Zr?
Question 7
What is the mass of 5.14 x 1023 atoms of uranium?
Question 8
What mass of C atoms have the same number of atoms as are in a 11.2g sample of Si?
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
TYPE 4: % by mass Composition.
Question 9
Calculate the percent by mass composition of ethanol, C2H6O.
Question 10
What is the percent by mass composition of N2O5?
Question 11
A compound has the formula Al4[Fe(CN)6]3. What is the percent by mass composition of this
compound?

TYPE 5: Empirical formulae.
Question 12
A compound contains 26.9% N and 73.1% F. What is the empirical formula of the compound?
Question 13
2.3g of magnesium is completely reacted with 6.75g of chlorine. What is the empirical formula
of the compound formed?

TYPE 6: Molecular formulae from empirical formulae.
Question 14
What is its molecular formula of hydrocarbon that has an empirical formula of C 2H5 and a
molecular mass of 58.
Question 15
A compound contains 68.54% carbon, 8.63% hydrogen, and the remainder oxygen.
The molecular weight of this compound is approximately 140g/mol. What is the
empirical formula? What is the molecular formula?
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 Type 7: Combustion Analysis.
Question 16
The combustion of 2.95 grams of a compound that contains only C, H and S yields 5.48
grams of CO2 and 1.13 grams of H2O. What is the empirical formula of the compound?
 Type 8: % Yield.
In questions 17 assume that the reactant that data is supplied for, is limiting.
Question 17
If 31 grams of C4H10 produces 41 grams of CO2 what is the % yield?
 Type 9: Limiting Reactant.
Question 18
Consider the reaction between Aluminum and Iron (III) oxide to produce Aluminum
oxide and Iron metal.
a)
Write an equation for the reaction.
b)
If 1240g of Al are reacted with 6010g of Iron (III) oxide, identify the limiting
reagent. Which reagent is in excess?
c)
Calculate the mass of Iron formed.
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d)
How much of the excess reagent is left over at the end of the reaction?
 Type 10: Analysis of hydrated salts.
Question 19
Copper (II) sulfate is found as a hydrated salt, CuSO 4.xH2O. A technician carefully heats
2.50g of the salt to a constant mass of 1.60g.
a) What is meant by constant mass?
b) How many moles of copper sulfate are there in 1.60g of anhydrous copper (II) sulfate?
c) How many moles of water were lost?
d) What is the value of x in the formula?
 Type 11: Moles and reacting ratios (including solutions).
Question 20
Sodium hydrogen carbonate, NaHCO3, combines with HCl as indicated below.
NaHCO3(aq) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l)
a) What volume of 1.5 M HCl solution should be present to combine totally with 0.14 moles
of NaHCO3?
b) How many moles of CO2 are produced when 0.49 g of NaHCO3 combines with excess
HCl?
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c) Calculate the mass of NaCl that results when 1.48 moles of HCl combines with excess
NaHCO3.
d) What mass of NaHCO3 is required to produce 6.1 x 103 moles of H2O?
 Type 12: Dilution.
Question 24
Calculate the volume of 2M sulfuric acid that must be diluted with water to produce
2dm3 of 0.5M sulfuric acid.
AP Topic 5 Review
In all questions show all relevant working and balance any equations you write.
1. Electrolytes & non-electrolytes
Indicate if you would expect the following compounds to be electrolytes or non-electrolytes when
in aqueous solution. In each case very briefly explain your answer. Use equations if
appropriate. (6)
a) Sodium chloride
b) Methanol (CH3OH, an alcohol similar to ethanol – NOT a hydroxide)
c) Magnesium nitrate
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3. Precipitation reactions and net ionic equations
Predict if a precipitation reaction will occur in each of the following cases. If it does, write the full,
balanced equation AND the net ionic equation (including state symbols) to show the formation of
the precipitate. If there is no reaction, say so, and indicate why. (9)
a) CuSO4(aq) + Na2CO3(aq) 
b) HI(aq) + Zn(NO3)2(aq) 
c) AgNO3(aq) + NaBr(aq) 
4. Acids and bases & neutralization
a) Write a full, balanced equation for the reaction of aqueous sulfuric acid and aqueous
potassium hydroxide to produce an aqueous salt and water.
b) Rewrite the equation in b) removing any spectator ions (i.e. write the net ionic equation)
with state symbols.
c) If I used 34.89 mL of a 0.345 M solution of sulfuric acid to neutralize a 20.00 mL sample
of potassium hydroxide, what is the molarity of the KOH solution?
5. Oxidation number concept
What is the oxidation number of each of the underlined atoms in each of the following species?
Think carefully about the rules that are being applied and write a very brief, simple explanation
of your answer in each case.
a) CaI2
b) GeO2
c) KO2
d) NH3
6. REDOX - Balance the following equations
Zn +
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HCl  Zn2+
+
H2
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(acidic)
Cr2O7
+
Cl- 
MnO4- + S2-

Cr3+
+
MnS
+
Cl2
S
(acidic)
(basic)
8. REDOX Titration calculations
a) Consider the following half equations.
SO32- + H2O  SO42- + 2H+ + 2eMnO4- + 8H+ + 5e-  Mn2+ + 4H2O
Combine these two equations to obtain the overall reaction of sulfite ions with manganate (VII)
ions.
b) Use the equation you have written in a) to calculate the volume of 0.222 M manganate
(VII) ions that are required to react completely with 25.0 cm 3 of 0.456 M sulfite ions.
9. Classification of chemical reactions
By choosing two of the following reaction types from the list below, classify each of the following
reactions in two ways.
Reaction types: precipitation, acid-base, REDOX (oxidation and reduction), single displacement,
double displacement, combination, decomposition, combustion.
a) 2NaOH(aq) + CuSO4(aq)  Na2SO4(aq) + Cu(OH)2(s)
AND
b) HI(aq) + KOH(aq)  KI(aq) + H2O(l)
AND
c) C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(l)
AND
d) 2HgO(s)  2Hg(l) + O2(g)
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AND
e) 4B(s) + 3O2(g)  2B2O3(s)
AND
f)
2Al(s) + 6HNO3(aq)  2Al(NO3)3(aq) + 3H2(g)
AND
AP Topic 6 Review
1. Small quantities of hydrogen gas can be prepared in the laboratory by the
following reaction:
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Assume you carried out this experiment and collected 653 mL of hydrogen gas over
water. The gas mixture collected includes hydrogen and water vapor. The temperature of
this gas mixture was 23.0 °C and the total pressure was 892 mm Hg. How many moles of
hydrogen did you prepare? The vapor pressure of water at 23.0°C is 19.8 mm Hg.
2. A mixture of H2(g), O2(g), and 2 milliliters of H2O(l) is present in 0.500-liter rigid
container at 25 °C. The number of moles of H2 and the number of moles of O2 are
equal. The total pressure is 1,146 millimeters of mercury. (The equilibrium vapor
pressure of pure water is 24 millimeters mercury.)
The mixture is sparked, and H2 and O2 react until one reactant is completely
consumed.
a. Identify the reactant remaining and calculate the number of moles of the
reactant remaining.
b. Calculate the total pressure in the container at the conclusion of the
reaction if the final temperature is 90 °C. (The equilibrium vapor pressure
of water at 90 °C is 526 millimeters mercury.)
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c. Calculate the number of moles of water present as vapor in the container
at 90 °C.
AP TOPIC 7 Review
2. Define ionization energy. (2)
_____________________________________________________________________________
_____________________________________________________________________________
3. Using the metal magnesium, write an equation to summarize the process of first ionization
energy. (Remember state symbols are important as they from part of the definition). (2)
_____________________________________________________________________________
4. Which of the following pairs has the largest size? (3)
PAIR 1
Ca, Ca2+
____________________________________
PAIR 2
Na, Mg
____________________________________
PAIR 3
S, S2-
____________________________________
5. Put the following species into order of increasing ionization energy (smallest first). (2)
Ca, Ca+, Ca2+, Sr
_____________________________________________________________________________
6. Given that the first ionization energy of sodium is 496 kJmol -1. Calculate (showing your working
and logic) the energy required to turn 2.3g of sodium from gaseous atoms to gaseous +1 ions. (2)
7. Explain carefully why rubidium tends only to form a +1 ion? (2)
_____________________________________________________________________________
_____________________________________________________________________________
8. Explain carefully why elements in the same group react in similar ways? (1)
_____________________________________________________________________________
_____________________________________________________________________________
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9. Explain carefully the factor, when moving up and down groups I & II, that determines the
pattern of reactivity that is observed? (2)
_____________________________________________________________________________
_____________________________________________________________________________
10. Write equations to illustrate the following.
(i) The basic nature of potassium oxide. (2)
_____________________________________________________________________________
(ii) The acidic nature of both the common oxides of sulfur. (4)
(SO2)_________________________________________________________________________
(SO3)_________________________________________________________________________
(iii) The amphoteric nature of aluminum oxide. (4)
(ACIDIC CHARACTER)__________________________________________________________
(BASIC CHARACTER)__________________________________________________________
11. Using the following data estimate the boiling point of bromine. (1)
ELEMENT
Boiling Point/K
Fluorine
Chlorine
Iodine
85
239
458
_____________________________________________________________________________
12. Define electron affinity. (2)
_____________________________________________________________________________
_____________________________________________________________________________
13. Write an equation to summarize the process of second electron affinity of oxygen. (2)
_____________________________________________________________________________
14. Consider the table of ionization energies for element X shown below.
Ionization
Energy in
kJ/mol
1st
2nd
3rd
4th
5th
6th
737
1450
7732
10540
13360
17995
(i) In which group will X be found? (1) ___________________________________________
(ii) Explain your answer to (i). (2)
_____________________________________________________________________________
_____________________________________________________________________________
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(iii) Predict the formula of X’s bromide. (1) __________________________________________
Topic 8 Review
Complete the Table
AROUND CENTRAL ATOM
Molecule
or ion
# of
Bonding
Areas
# of Lone
pairs
Atom geometry
SF6
PCl5
SO3-2
CO32-
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Lewis dot structure and
hybridization