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NAME: _________________________________________ DATE: ____________
1. What gas is formed when Zn metal is mixed with hydrochloric acid, HCl?
a) CO2
c) O2
b) He
d) H2
2. Which property is always conserved during a chemical reaction?
a) mass
c) pressure
b) volume
d) solubility
3. A cylinder is weighed empty and with a liquid.
Cylinder with liquid
51.85 g
Cylinder, empty
40.11 g
Volume of liquid in
7.0 mL
cylinder
What is the density of the liquid?
a) 13 g/mL
c) 5.7 g/mL
b) 7.4 g/mL
d) 1.7 g/mL
4. Which one of the following is the correct formula for aluminum oxide?
a) AlO
c) Al2O3
b) Al6O6
d) Al3O2
5. What is the name of the compound CF4?
6.
7.
8.
a) fluorocarbonate
b) carbon tetrafluoride
c) tricarbo fluoride
d) carbon difluorate
Sodium nitride has the formula Na3N. What is the formula for magnesium nitride?
a) Mg2N
c) Mg3N2
b) Mg3N
d) Mg2N3
What is the mass of one mole of aluminum sulfate, Al2(SO4)3?
a) 630 g
c) 273 g
b) 342 g
d) 123 g
Which set of coefficients balances the equation for the complete combustion of ethane, C2H6?
__C2H6 + __O2  __CO2 + __H2O
a) 1,3,2,3
c) 2,6,4,5
b) 1,6,2,6
d) 2,7,4,6
9. When this expression is balanced,
2C3H6 + O2 
CO2 + 6H2O
what is the coefficient of oxygen, O2?
a) 6
c) 12
b) 9
d) 18
10. During a “titration lab,” an acid was neutralized by the following reaction:
NaOH + HCl  NaCl + H2O
This reaction would be classified as…
a) synthesis
b) decomposition
c) double replacement
d) single replacement
11. Which reaction below would be classified as a decomposition reaction?
a) NaHCO3  NaOH + CO2
b) 2 H2 + O2  2 H2O
c) 2 AgNO3 + Cu°  Cu(NO3)2 + 2 Ag°
d) Ba(OH)2 + H2SO4  BaSO4 + 2 H2O
12. The complete combustion of ethane, C2H6, produces
a) C2H5OH
c) CO2 and H2
b) CH3COOH
d) CO2 and H2O
13. What quantity of sulfur dioxide, SO2 (64.0 g/mole), is produced when 245 g of sulfuric acid,
H2SO4 (98.0 g/mole) reacts completely with zinc metal?
Zn° + 2 H2SO4  ZnSO4 + SO2 + 2 H2O
a) 64.0 g
c) 128 g
b) 80.0 g
d) 160 g
14. How many moles of FeS2 are required to produce 64 grams of SO2 according to the equation
4FeS2(s) + 11O2(g)  2Fe2O3(s) + 8SO2(g)
a) 0.40
c) 3.2
b) 0.50
d) 4.5
15. Glass, SiO2, reacts with hydrofluoric acid, HF, according to this equation
SiO2 + 4 HF  2 H2O + SiF4
Which reagent is completely consumed when 2 moles of SiO2 is added to 6 moles of HF?
a) SiF4
c) HF
b) H2O
d) SiO2
16. Which of the following is an acid?
a) NaOH
c) HCl
b) NH3
d) KOH
17. The acid, H2S, is correctly named as:
a) hydrosulfuric acid
b) sulfuric acid
c) dihydrogen sulfide
d) hydrogen(I) sulfide
18. Hydrogen gas was produced according to the following equation:
Zn° + 2 HCl  ZnCl2 + H2
Which chemical is oxidized?
a) Zn°
c) ZnCl2
b) HCl
d) H2
Table of common polyatomic ions, arranged by charge.
2+
Hg22+ mercury(I)
2-
1CH3COO- acetate
CO32-
ClO3-
chlorate
CrO42- chromate
ClO2-
chlorite
Cr2O72- dichromate
CN-
cyanide
NH4+ ammonium
H2PO4-
dihydrogen phosphate
HPO42- hydrogen
phosphate
H3O+ hydronium
HCO3-
hydrogen carbonate
or
mercurous
1+
or
bicarbonate
HSO4-
hydrogen sulfate
OH-
hydroxide
ClO-
hypochlorite
NO3-
nitrate
NO2-
nitrite
ClO4-
perchlorate
MnO4-
permanganate
SCN-
thiocyanate
or bisulfate
carbonate
O22-
peroxide
SO42-
sulfate
SO32-
sulfite
S2O32-
thiosulfate
3PO43-
phosphate
SOLUBILITY RULES:
Salts containing the following ions are normally soluble:
 All salts of group IA (Li+, Na+, etc) and the ammonium ion (NH4+) are soluble.
 All salts containing nitrate (NO3-)acetate ( CH3COO-), and perchlorates are
soluble.
 All chlorides (Cl-), Bromides (Br-), and iodides(I-) are soluble except those of
Cu+, Ag+, Pb2+, and Hg22+
 All salts containing sulfate (SO42-) are soluble except those of Pb2+, Sr2+, and
Ba2+.
Salts containing the following ions are normally insoluble:
 Most carbonates (CO32-) and phosphates (PO43-) are insoluble except those of
group IA and the ammonium ion.
 Most sulfides (S2-) are insoluble except those of group IA and IIA and the
ammonium ion.
 Most hydroxides (OH-) are insoluble except those of group IA, calcium,
strontium, and barium.
 Most oxides (O2-) are insoluble except those of group IA, and calcium,
strontium, and barium which react with water.
Exercise #3
Type of
Compound
How To
Recognize
How To Name
Types of Compounds
Ionic
Acids
metal + non-metal
starts with H + anion
names of + ion then ion
“ides”  hydro---ic acid
“ates”  ----ic acid
S (add “ur”) P (add “or”)
N A M I N G
Molecular
two non-metals
mono, di, tri, tetra, penta,
hexa, hepta, octa, nona ,deca
names ends with “ide”
pentaoxide  pentoxide, etc.
Indicate the Type of Compound and then name the compound using the appropriate rules:
1.
NaF
___ __________________
21. CuCl2
___ __________________
2.
FeCl3
___ __________________
22. AgNO3
___ __________________
3.
CO2
___ __________________
23. CO
___ __________________
4.
MgCl2
___ __________________
24. H3PO4
___ __________________
5.
HF
___ __________________
25. NaCl
___ __________________
6.
SF4
___ __________________
26. N2O5
___ __________________
7.
HC2H3O2
___ __________________
27. NO2
___ __________________
8.
H2O
___ __________________
28. HNO3
___ __________________
9.
NH3
___ __________________
29. NaOH
___ __________________
10. CaO
___ __________________
30. SnCl2
___ __________________
11. NH4NO3
___ __________________
31. CaSO4
___ __________________
12. NaI
___ __________________
32. HBr
___ __________________
13. PbCO3
___ __________________
33. Cu(OH)2
___ __________________
14. Na2O
___ __________________
34. Zn(OH)2
___ __________________
15. Ba(NO3)2 ___ __________________
35. BaCl2
___ __________________
16. K2CrO4
___ __________________
36. PCl5
___ __________________
17. NO
___ __________________
37. PCl3
___ __________________
18. HCl
___ __________________
38. AsF5
___ __________________
19. MnO2
___ __________________
39. H2CO3
___ __________________
20. H2S
___ __________________
40.
___ __________________
OF2
Exercise #4
Chemical Equations and Stoichiometry
WRITING IONIC COMPOUNDS
Ionic compounds are formed from a positive ion (cation) and a negative ion (anion).
The positive ion is always written first.
The resulting compound must be electrically neutral.
Use parentheses when you need two or more polyatomic ions in a formula.
Cl
OH
S2
CO32
PO43
Na+
NH4+
Ca2+
Al3+
Sn4+
Hg22+
Naming ionic compounds is easy.
The name is simply the name of the cation, followed by the name of the anion.
#
1.
Name
ammonium phosphate
Cation
NH4+
Anion
PO43-
2.
barium nitrate
Ba2+
NO3-
3.
cuprous sulfide or copper (I)
sulfide
Cu+
S2-
4.
aluminum carbonate
Al3+
CO32-
5.
strontium hydroxide
Sr2+
OH-
Formula
Exercise #6 2f: Inorganic Nomenclature II
Add either a name or a formula to complete each table.
1. Potassium dichromate
2. Lithium sulfide
3. Potassium bromide
4. Cesium iodide
5. Calcium phosphide
6. Sodium fluoride
7. Strontium oxide
8. Beryllium sulfide
9. Magnesium bromide
10. Lithium oxide
11. Strontium chloride
12. Barium bromide
13. Magnesium sulfide
14. Magnesium iodide
15. Hydrogen fluoride (Hydrogen monofluoride)
16. Barium phosphide
17. Sodium hydrogen phosphate
18. Potassium chloride
19. Lithium nitride
20. Calcium sulfide
21. Rubidium oxide
22. Strontium nitride
23. Cesium phosphide
24. Magnesium carbonate
25. Beryllium sulfate
26. ScCl3
27. HCl
28. PtO2
29. Sb(ClO3)5
30. GeS2
31. ZnO
32. VSO4
33. CuCl2
34. TiO2
35. NiN
36. Ni3(PO4)2
37. CoF3
38. Au2O3
39. Zn3P2
40. Cr(NO3)6
41. NaIO2
42. NaIO3
43. NaI
44. H2SO3
45. H2CO3
46. AlN
47. AlH3
48. Li3AsO4
49. NaCN
50. Na2O2
Exercise #7: Balancing Equations I
Balance the following equations.
O2

H2O
H3PO4 +
NaOH

Na3PO4
+
H2O
3.
Na
+
B2O3

Na2O
+
B
4.
HCl
+
KOH

KCl
+
H2O
5.
K
+
KNO3

K2O
+
N2
6.
C
+
S

CS2
7.
Na
+
O2

Na2O2
8.
N2
+
O2

N2O4
9.
H3PO4 +
Ca(OH)2

Ca3(PO4)2
+
H2O
10.
KOH
H2CO3

K2CO3
+
H2O
11.
NaOH +
HBr

NaBr
+
H2O
12.
H2
+
O2

H2O2
13.
K
+
O2

K2O
14.
Al(OH)3 +
H2SO3

Al2(SO3)3
+
H2O
15.
Al
+
S8

Al2S3
16.
Li
+
N2

Li3N
17.
Ca
+
Cl2

CaCl2
18.
Rb
+
RbNO3

Rb2O
+
N2
19.
C6H12
+
O2

CO2
+
H2O
20.
N2
+
H2

NH3
1.
H2
2.
+
+
Exercise #8: Stoichiometry Summary

TYPE 1: Those involving Avogadro’s number.
Question 1
A sample of Ag is found to contain 9.7 x 10 23 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?

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?

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 CO 2 and 1.13
grams of H2O. What is the empirical formula of the compound?
Question 17
If, in the reaction below, of 31 grams of C4H10 produces 41 grams of CO2 what is the % yield?
2C4H10 + 13O2  8CO2 + 10H2O
Question 18
If, in the reaction below, 80.1 grams of Cl2 produces 33.12 grams of CCl4 what is the % yield?
CS2 + 3Cl2  CCl4 + S2Cl2

Type 9: Limiting Reactant.
Question 19
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.
d) How much of the excess reagent is left over at the end of the reaction?

Type 10: Analysis of hydrated salts.
Question 20
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 21
Sodium hydrogen carbonate, NaHCO3, combines with HCl as indicated below.
NaHCO3(aq) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l)
a) What volume of 1.5M 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?
c) Calculate the mass of NaCl that results when 1.48 moles of HCl combines with excess NaHCO 3.
d) What mass of NaHCO3 is required to produce 6.1 x 103 moles of H2O?
Question 22
Carbon tetrachloride, CCl4, can be produced in the reaction below.
CH4 + 4Cl2  CCl4 + 4HCl
a) What mass of CH4 is needed to exactly combine with 3.4 g Cl2?
b) How many grams of Cl2 are required to produce 91 g CCl4, assuming excess CH4?
c) What mass of CH4 must have reacted, if 2 mg HCl is liberated?
d) Calculate the mass of both CH4 and Cl2 required to produce exactly 0.761 kg CCl4?
Exercise #10: Writing chemical equations
Write balanced equations for the following reactions. Where possible include state(g, l, s) symbols.
1. Pure, molten iron forms when solid iron (III) oxide reacts with carbon monoxide gas. Carbon dioxide gas
is also a product.
2. Potassium oxide reacts with water to produce potassium hydroxide.
3. During photosynthesis glucose (C6H12O6) forms from carbon dioxide and water. Oxygen is also a
product.
4. Sodium phosphate and barium sulfate are made during a reaction between sodium sulfate and barium
phosphate.
5. Ammonium nitrate can decompose explosively to form nitrogen, water and oxygen.
6. The combustion (combination with oxygen) of liquid octane (C 3H8) produces gaseous carbon dioxide and
steam.
7. The combination of sodium metal and chlorine gas yields solid sodium chloride.
8. Hydrogen gas forms when magnesium metal comes in contact with an aqueous solution of ethanoic acid
(CH3COOH). An aqueous solution of magnesium ethanoate Mg (CH3COO)2 is the other product.
9. The decomposition of solid copper (II) nitrate yields solid copper (II) oxide and nitrogen dioxide, and
oxygen gases.
10. Solid mercury (II) oxide forms from the uncombined elements. Mercury is a liquid at room temperature.
11. Magnesium metal and steam combine to form solid magnesium hydroxide and hydrogen gas.
12. An aqueous solution of hydrogen peroxide (H2O2) and solid lead (II) sulfide combine to form solid lead
(II) sulfate and water.
13. Solid Sodium reacts with liquid water to produce aqueous sodium hydroxide and hydrogen gas.
14. Zinc reacts with silver nitrate to produce silver and zinc nitrate.
15. Aluminum sulfate reacts with calcium chloride to produce aluminum chloride and calcium sulfate.
16. Solid potassium hydroxide pellets decompose on heating to form solid potassium oxide and water.
Name_____________________
Period __ Date __/__/__
Exercise #9_Chemical
Equations and Stoichiometry
COMBUSTION EQUATIONS
For burning to occur, you need a fuel, an oxidizer, and heat. When hydrocarbons are the fuel and
O2 in the air is the oxidizer, then CO2 and H2O are the products.
Example:
Write the balanced equation for the complete combustion of propane, C3H8, in air.
Solution:
First, set up the basic equation. You memorize the “+ O2  CO2 + H2O” part.
C3H8 + O2  CO2 + H2O
Next, balance. 3 C’s in C3H8 result in 3CO2’s; 8 H’s in C3H8 result in 4 H2O’s;
C3H8 + __ O2  3 CO2 + 4 H2O
Total O’s on the product side = 10 [(3 x 2) + (4 x 1)] = total O’s on the
reactant side.
This would mean that 5 O2’s were involved.
Tip: If an UNEVEN number of O’s need to be represented, a fraction should be used. 7 O’s = 7/2
O2
Tip: Take into account fuels that contain oxygen. Subtract the O’s from that represented as O2’s
Practice: Write the balanced combustion equations for the following substances.
1.
CH4
2.
C5H12
3.
C9H20
4.
C2H6
5.
C8H18
6.
C4H10
7.
C2H5OH
8.
C3H7OH
9.
HC2H3O2
10.
CH3COCH3
Name___________________________
______
Period ___ Date ___/___/___
Exercise #11 Acids & Bases
TITRATION PRACTICE
An acid and a base neutralize each other when the moles of H+ = the moles of OH. A formula
similar to the dilution formula can be used to determine the concentration of an unknown acid or
base.
VaMa=VbMb
Example:
where a = H+ b = OH-
What is the concentration of a 10.0 mL sample of HCl if 35.5 mL of 0.150 M NaOH
is needed to titrate it to a pink endpoint?
(10.0 mL) (x) = (35.5 mL) (0.150 M)
x=
(35.5 mL)(0.150 M)
= 0.5325 M = 0.533 M
10.0 mL
1. What is the concentration of a 15.0 mL sample of HCl if 28.2 mL of 0.150 M NaOH is needed
to titrate it?
2. A 10.0 mL sample of a monoprotic acid is titrated with 45.5 mL of 0.200 M NaOH. What is
the concentration of the acid?
3. A 5.00 mL sample of vinegar has a concentration of 0.800 M. What volume of 0.150 M NaOH
is required to complete the titration?
4. A 10.0 sample of household ammonia, NH3(aq), is titrated with 0.500 M HCl. If 25.7 mL of
acid is required, what is the concentration of the household ammonia?
5. A 5.00 mL sample of H2SO4 is titrated with 0.150 M NaOH. If 20.0 mL of the base is required
to titrate the acid sample, what is the [H+] of the acid? ________ What is [H2SO4]?
_________
12  The Gas Laws
THE IDEAL
GAS LAW
PV = nRT where
P = pressure in atmosphere
V = volume in liters
n = number of moles of gas
R = Universal Gas Constant = 0.0821 Latm/molK
T = Kelvin temperature
1.
What is the pressure of 1.20 moles of SO2 gas in a 4.00 L container at 30°C?
2.
How many moles of oxygen will occupy a volume of 2.50 liters at 1.20 atm and 25 C?
3.
What is the volume of 0.60 moles of helium gas at 50°C if the pressure is 600 torr?
4.
At what temperature will 1.80 moles of gas occupy 4.00 L if the pressure is 350 mmHg?
5.
A balloon filled with helium has a volume of 1.30 L at 15°C when the atmospheric pressure
is 700 torr. How many molecules of helium is in the balloon?
6.
What is the mass of a 300 mL sample of gaseous hydrogen chloride at 2.0 atm and 30°C?
. m = µPV/RT where m is mass in grams and µ is molar mass
What is the density of this sample? D = m/V where D is density in grams per Liter
7.
A 30.0 g sample of CH4 occupies 150 mL at 0°C. What is the pressure of this sample of gas?
P = mRT/Vµ
8.
What volume (in liters) does a 85.0-g sample of CO2 gas occupy at 1.40 atm and 80°C?
V = mRT/µP
What is the density of this gas? D = m/V
9.
How many moles of an unknown diatomic gas are contained in a 500 mL container at 1.20
atm and 25°C?
If this unknown diatomic gas has a mass of 0.93 g, what is the molar mass of the gas? What
is the gas?
µ = g/mol
12  The Gas Laws
BOYLE’S LAW
Boyle’s Law states that the volume of a gas varies inversely with its pressure if temperature is held
constant.
(If one goes up, the other goes down.) We use the formula:
P1  V1 = P2  V2
Solve the following problems (assuming constant temperature). Assume all number are 3
significant figures.
1.
A sample of oxygen gas occupies a volume of 250 mL at 740 torr pressure. What volume will
it occupy at 800 torr pressure?
2.
A sample of carbon dioxide occupies a volume of 3.50 Liters at 125 kPa pressure. What
pressure would the gas exert if the volume was decreased to 2.00 liters?
3.
A 2.00-Liter container of nitrogen had a pressure of 3.20 atm. What volume would be
necessary to decrease the pressure to 1.00 atm?
4.
Ammonia gas occupies a volume of 450 mL as a pressure of 720 mmHg. What volume will
it occupy at standard pressure (760 mmHg)?
5.
A 175 mL sample of neon had its pressure changed from 75.0 kPa to 150 kPa. What is its
new volume?
6.
A sample of hydrogen at 1.50 atm had its pressure decreased to 0.50 atm producing a new
volume of 750 mL. What was the sample’s original volume?
7.
Chlorine gas occupies a volume of 1.20 liters at 720 torr pressure. What volume will it
occupy at 1 atm pressure?
8.
Fluorine gas exerts a pressure of 900 torr. When the pressure is changed to 1.50 atm, its
volume is 250 mL. What was the original volume?
12  The Gas Laws
CHARLES’S LAW
Charles’ Law states the volume of a gas varies directly with the Kelvin temperature, assuming the
pressure is constant. We use the following formulas:
V1 V2

T1 T2
or
V1  T2 = V2  T1
K = C + 273
Solve the following problems assuming a constant pressure. Assume all numbers are 3
significant figures.
1.
A sample of nitrogen occupies a volume of 250 mL at 25 C. What volume will it occupy at
95 C?
2.
Oxygen gas is at a temperature of 40 C when it occupies a volume of 2.30 Liters. To what
temperature should it be raised to occupy a volume of 6.50 Liters?
3.
Hydrogen gas was cooled from 150 C to 50 C. Its new volume is 75.0 mL. What was its
original volume?
4.
Chlorine gas occupies a volume of 25.0 mL at 300 K. What volume will it occupy at 600 K?
5.
A sample of neon gas at 50 C and a volume of 2.50 Liters is cooled to 25 C. What is the
new volume?
6.
Fluorine gas at 300 K occupies a volume of 500 mL. To what temperature should it be
lowered to bring the volume to 300 mL?
7.
Helium occupies a volume of 3.80 Liters at –45 C. What volume will it occupy at 45 C?
8.
A sample of argon gas is cooled and its volume went from 380 mL to 250 mL. If its final
temperature was –55 C, what was its original temperature?
12  The Gas Laws
THE COMBINED GAS LAW
In practical terms, it is often difficult to hold any of the variables constant. When there is a change
in pressure, volume and temperature, the combined gas law is used.
P1V1 P2 V2

T1
T2
or
P1  V1  T2 = P2  V2  T1
K = C + 273
Complete the following chart.
1
P1
V1
T1
P2
1.50 atm
3.00 L
20.0 C
2.50 atm
V2
T2
30.0 C
2
3
25.0 C
250. mL
50.0 C
720. torr
256. mL
600. mmHg
2.50 L
22.0 C
760. mmHg
1.80 L
750. mL
0.00 C
2.00 atm
500. mL
25.0 C
101. kPa
6.00 L
471. K or
198. C
900. torr
225. mL
150. C
2.50 L
30.0 C
100 mL
75.0 C
4
5
95.0 kPa
4.00 L
6
650. torr
7
850. mmHg
1.50 L
8
125. kPa
125. mL
100. C
15.0 C
100. kPa
Oxidation Numbers:
1) In which reaction does the oxidation number of oxygen increase?
A) Ba(NO3 )2 (aq) + K 2SO4 (aq)  BaSO4 (s) + 2 KNO3 (aq)
B) HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l )
C) MgO (s) + H2 O (l )  Mg(OH)2 (s)
D) 2 SO2 (g) + O2 (g)  2 SO3 (g)
E) 2 H2 O (l )  2 H2 (g) + O2 (g)
2) In which reaction does the oxidation number of hydrogen change?
A) HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l )
B) 2 Na (s) + 2 H2 O (l )  2 NaOH (aq) + H2 (g)
C) CaO (s) + H2 O (l )  Ca(OH)2 (s)
D) 2 HClO4 (aq) + CaCO3 (s)  Ca(ClO4 )2 (aq) + H2 O (l ) + CO2 (g)
E) SO2 (g) + H2 O (l )  H2 SO3 (aq)
3) In which species does sulfur have the highest oxidation number?
A) S8 (elemental form of sulfur)
B) H 2S
C) SO 2
D) H 2SO3
E) K 2SO4