Possible Practice Final Exam

NAME: ______DATE: ______

1.What gas is formed when Zn metal is mixed with hydrochloric acid, HCl?

a)CO2c)O2

b)Hed)H2

2.Which property is always conserved during a chemical reaction?

a)massc)pressure

b)volumed)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 cylinder / 7.0 mL

What is the density of the liquid?

a)13 g/mLc)5.7 g/mL

b)7.4 g/mLd)1.7 g/mL

4.Which one of the following is the correct formula for aluminum oxide?

a)AlOc)Al2O3

b)Al6O6d)Al3O2

5.What is the name of the compound CF4?

a)fluorocarbonate

b)carbon tetrafluoride

c)tricarbo fluoride

d)carbon difluorate

6.Sodium nitride has the formula Na3N. What is the formula for magnesium nitride?

a)Mg2Nc)Mg3N2

b)Mg3Nd)Mg2N3

7.What is the mass of one mole of aluminum sulfate, Al2(SO4)3?

a)630 gc)273 g

b)342 gd)123 g

8.Which set of coefficients balances the equation for the complete combustion of ethane, C2H6?

__C2H6 + __O2 __CO2 + __H2O

a)1,3,2,3c)2,6,4,5

b)1,6,2,6d)2,7,4,6

9.When this expression is balanced,

2C3H6 + O2 CO2 + 6H2O

what is the coefficient of oxygen, O2?

a)6c)12

b)9d)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)C2H5OHc)CO2 and H2

b)CH3COOHd)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 gc)128 g

b)80.0 gd)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)0.40c)3.2

b)0.50d)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)SiF4c)HF

b)H2Od)SiO2

16.Which of the following is an acid?

a)NaOHc)HCl

b)NH3d)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)HCld)H2

Table of common polyatomic ions, arranged by charge.

2+
Hg22+ / mercury(I) or mercurous

1+
NH4+ / ammonium
H3O+ / hydronium
/ 1-
CH3COO- / acetate
ClO3- / chlorate
ClO2- / chlorite
CN- / cyanide
H2PO4- / dihydrogen phosphate
HCO3- / hydrogen carbonate or bicarbonate
HSO4- / hydrogen sulfate or bisulfate
OH- / hydroxide
ClO- / hypochlorite
NO3- / nitrate
NO2- / nitrite
ClO4- / perchlorate
MnO4- / permanganate
SCN- / thiocyanate
/ 2-
CO32- / carbonate
CrO42- / chromate
Cr2O72- / dichromate
HPO42- / hydrogen phosphate
O22- / peroxide
SO42- / sulfate
SO32- / sulfite
S2O32- / thiosulfate

3-
PO43- / 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 solubleexcept those of Cu+, Ag+, Pb2+, and Hg22+
All salts containing sulfate (SO42-) are solubleexcept those of Pb2+, Sr2+, and Ba2+.

Salts containing the following ions are normally insoluble:

Most carbonates (CO32-) and phosphates (PO43-) are insolubleexcept those of group IA and the ammonium ion.
Most sulfides (S2-) are insolubleexcept those of group IA and IIA and the ammonium ion.
Most hydroxides (OH-) are insolubleexcept those of group IA, calcium, strontium, and barium.
Most oxides (O2-) are insolubleexcept those of group IA, and calcium, strontium, and barium which react with water.

Exercise #3 Types of Compounds

NAMING

Type of Compound / Ionic / Acids / Molecular
How To Recognize / metal + non-metal / starts with H + anion / two non-metals
How To Name / names of + ion then - ion / “ides”  hydro---ic acid
“ates”  ----ic acid
S (add “ur”) P (add “or”) / 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______
2.FeCl3______
3.CO2______
4.MgCl2______
5.HF______
6.SF4______
7.HC2H3O2______
8.H2O______
9.NH3______
10.CaO______
11.NH4NO3______
12.NaI______
13.PbCO3______
14.Na2O______
15.Ba(NO3)2______
16.K2CrO4______
17.NO______
18.HCl______
19.MnO2______
20.H2S______/ 21.CuCl2______
22.AgNO3______
23.CO______
24.H3PO4______
25.NaCl______
26.N2O5______
27.NO2______
28.HNO3______
29.NaOH______
30.SnCl2______
31.CaSO4______
32.HBr______
33.Cu(OH)2______
34.Zn(OH)2______
35.BaCl2______
36.PCl5______
37.PCl3______
38.AsF5______
39.H2CO3______
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.

# / Name / Cation / Anion / Formula
1. / ammonium phosphate / NH4+ / 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-

Exercise #6 2f: Inorganic Nomenclature II

Add either a name or a formula to complete each table.

Potassium dichromate

Lithium sulfide

Potassium bromide

Cesium iodide

Calcium phosphide

Sodium fluoride

Strontium oxide

Beryllium sulfide

Magnesium bromide

Lithium oxide

Strontium chloride

Barium bromide

Magnesium sulfide

Magnesium iodide

Hydrogen fluoride(Hydrogen monofluoride)

Barium phosphide

Sodium hydrogen phosphate

Potassium chloride

Lithium nitride

Calcium sulfide

Rubidium oxide

Strontium nitride

Cesium phosphide

Magnesium carbonate

Beryllium sulfate

ScCl3

PtO2

Sb(ClO3)5

GeS2

VSO4

CuCl2

TiO2

Ni3(PO4)2

CoF3

Au2O3

Zn3P2

Cr(NO3)6

NaIO2

NaIO3

H2SO3

H2CO3

AlH3

Li3AsO4

NaCN

Na2O2

Exercise #7: Balancing Equations I

Balance the following equations.

1. H2+ O2H2O

2. 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

Exercise #8: Stoichiometry Summary

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?

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 C2H5 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 CO2 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, CuSO4.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 NaHCO3.

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 (C3H8) 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 = 7/2O2

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 where a = H+ b = OH-

Example: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 = = 0.5325 M = 0.533 M

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:

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.

or P1 V1 T2 = P2 V2 T1
K = C + 273

Complete the following chart.

P1 / V1 / T1 / P2 / V2 / T2
1 / 1.50 atm / 3.00 L / 20.0 C / 2.50 atm / 30.0 C
2 / 720. torr / 256. mL / 25.0 C / 250. mL / 50.0 C
3 / 600. mmHg / 2.50 L / 22.0 C / 760. mmHg / 1.80 L
4 / 750. mL / 0.00 C / 2.00 atm / 500. mL / 25.0 C
5 / 95.0 kPa / 4.00 L / 101. kPa / 6.00 L / 471. K or 198. C
6 / 650. torr / 100. C / 900. torr / 225. mL / 150. C
7 / 850. mmHg / 1.50 L / 15.0 C / 2.50 L / 30.0 C
8 / 125. kPa / 125. mL / 100. kPa / 100 mL / 75.0 C

Oxidation Numbers:

1) In which reaction does the oxidation number of oxygen increase?

B) HCl (aq) + NaOH (aq) → NaCl (aq) +

2) In which reaction does the oxidation number of hydrogen change?

A) HCl (aq) + NaOH (aq) → NaCl (aq) +

3) In which species does sulfur have the highest oxidation number?

A) (elemental form of sulfur)