CHEMISTRY 104 Rogue Community College HOMEWORK...

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CHEMISTRY 104

Rogue Community College

HOMEWORK SHEETS

to be used with General, Organic and Biochemistry, 2nd edition, by Janice Gorzynski Smith

Classification of Matter (CH. 1) ….………………………………………………..... 3

Significant Figures (CH. 1) ……..….…………………..……………………………... 5

Unit Conversions (CH. 1) …………………..…..…..………………………………... 7

Temperature Scales (CH. 1) ……………………….………………………………….. 9

Atoms and Formulas (CH. 2) ………………….……...………………………………. 11

Atomic Structure (CH. 2) ………..…………………...………………………………. 13

Periodic Table (CH. 2) …………………………………….………………………….. 15

Chemical Bonding (CH. 3&4) ………...………………………..………...…………... 17

Naming and Formula Writing (CH. 3&4) ……………………………..……………... 19

Molecular Geometry (CH. 4) …………………………………….………..…………… 21

Balancing Equations (CH. 5) ………………..……………..………………………….. 23

Moles (CH. 5) ………………………………………………………………….……… 25

Stoichiometry (CH. 5) …….…………………………………………………………... 27

Energy (CH. 6) ………………….………………………………………...…………… 29

Reaction Rate (CH. 6) ………………………………………………………..……….. 31

Gas Laws (CH. 7) …………………………………………………………………….. 33

Intermolecular Forces (CH. 7) ……………………………………………...…….….. 35

Solutions (CH. 8) ….………………………………………………………………..... 37

Acids and Bases (CH. 9) ….…………………………………………………….…..... 39

Neutralization and Buffers (CH. 9) ….……………………...…………………..…..... 41

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Chemistry 104 – Rogue Community College

Classification of Matter Homework

Name ___________________________

Distinguish between each of the following pairs of terms. Do not use examples!

1. mass and weight

2. element and compound

3. substance and mixture

4. physical change and chemical change

5. liquid and gas

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6. Classify each of the following as an element, a compound or a mixture:

water _______________ cheese _______________

mercury _______________ air _______________

sugar _______________ salt _______________

oxygen _______________ aluminum _______________

helium _______________ brass _______________

vinegar _______________ coffee _______________

steel _______________ iron _______________

steam _______________ ice _______________

ice cream _______________ seawater _______________

methane _______________ copper _______________

granite _______________ blood _______________

arsenic _______________ krypton _______________

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Chemistry 104, Rogue Community College

Significant Figures Homework

Name ___________________________

Part One. Record each of the following three measurements to the appropriate number of

significant figures. Include the appropriate units.

Centimeter Ruler:

4 5 6321

5.0 5.5

3.96 4.00

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Part Two.

Rewrite the following measurements using scientific notation. Do not change the units.

1. 0.00573 g -------->

2. 0.45400 kg -------->

3. 12055 years -------->

4. 93,000,000 miles -------->

Part Three. Write down the number of significant figures in each of the following

measurements.

1. 0.00503 g -------->

2. 500 miles -------->

3. 1.75 yards -------->

4. 0.75 yards -------->

5. 1.00 x 10-10 m -------->

6. 93,000,000 miles -------->

7. 6.022 x 1023 atoms -------->

Part Four. Perform each of the following calculations. Record each answer to the correct

number of significant figures and in the correct units.

mL

g

24.2

75.3 = _______________ cmxcm 74.275.2 = _______________

g

g

2900.0

05.2 = _______________ gg 4.129.11 = _______________

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3.12 g = _______________ cgg 4.129.11 = _______________

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Unit Conversions Homework

Name ___________________________

Set up and solve the problem in the space provided. Show all of your work!

1. Convert the measurement 92.0 g into pounds. (1 lb = 453.6 g)

2. Convert the measurement 92.0 cm into inches. (2.54 cm = 1 in)

3. Convert the measurement 2.15 mi into kilometers. (1 mi = 1.609 km)

4. Convert the measurement 0.001065 metric tons into grams. (1 metric ton = 1 x 106 g)

5. Convert the measurement 0.0039 L into cm3. (1 cm3 = 1 mL)

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6. Convert the measurement 45.0 mg into grams.

7. Convert the measurement 45.0 g into milligrams.

8. Convert the measurement 1.25 L into milliliters.

9. Convert the measurement 15.0 cm into mm.

10. Convert the measurement 125 g into g.

11. Convert the measurement 0.0058 m into cm.

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Temperature Scales Homework

Name ___________________________

There are three common temperature scales: Fahrenheit (F), Celsius (C) and kelvins (K).

It is easiest to convert between kelvins and Celsius: TK = Tc + 273.15 or Tc = TK 273.15.

It is also straightforward to convert between Celsius and Fahrenheit: TF = 1.8(Tc) + 32

or Tc = (TF 32)(5/9).

To convert between kelvins and Fahrenheit, it is easiest to do it in two steps: first convert from

kelvins to Celsius and then from Celsius to Fahrenheit, or first convert from Fahrenheit to

Celsius and then from Celsius to kelvins.

Practice using the above equations on the following temperature conversions:

1. Convert normal body temperature (98.6F) into Celsius.

2. Convert the freezing point of water (0C exactly) into kelvins.

3. Convert room temperature (25C) into Fahrenheit.

4. Convert absolute zero (0 K exactly) into Celsius.

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5. Convert the temperature of liquid nitrogen (77 K) into Celsius and Fahrenheit.

6. Convert the temperature of dry ice (78.5C) into Fahrenheit and kelvins.

7. Convert the temperature of the surface of Venus encountered by Soviet spacecraft (about

465C) into Fahrenheit and kelvins.

8. Choose a temperature that is of interest to you for some reason. What makes this

temperature interesting? Express this temperature in C, F and K.

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Atoms and Formulas Homework

Name ___________________________

1. Classify each of the following elements as a metal, nonmetal or metalloid.

(a) lithium (Li) ____________________________

(b) beryllium (Be) ____________________________

(c) boron (B) ____________________________

(d) carbon (C) ____________________________

(e) nitrogen (N) ____________________________

(f) oxygen (O) ____________________________

(g) fluorine (F) ____________________________

2. Fill in the blanks to accurately complete the following statements about the structure of

the atom.

(a) All atoms of the same element have the same number of ________________ in

the ________________.

(b) The subatomic particle with no electrical charge is called the ________________.

(c) When an atom has no overall electrical charge, it has the same number of

________________ and ________________.

(d) Atoms of the same element that have different weights are called

________________ of that element.

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3. Distinguish between the terms atomic number and mass number. Do not use examples!

4. Write the chemical formula for a compound that contains:

(a) two atoms of hydrogen and one atom of oxygen

(b) one atom each of carbon and oxygen

(c) two atoms of hydrogen, one atom of carbon and three atoms of oxygen

(d) two atoms of potassium, one atom of sulfur and four atoms of oxygen

(e) three atoms of sodium, one atom of phosphorus and four atoms of oxygen

5. How many of each kind of atom are present in each of the following formulas?

(a) KOH

K = O = H =

(b) AlCl3

Al = Cl =

(c) Mg3(PO4)2

Mg = P = O =

(d) (NH4)2SO3

N = H = S = O =

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Atomic Structure Homework

Name ___________________________

1. How many protons, neutrons and electrons are there in each of the following species?

(a) 612C p+ (b) 6

14C+

p+ (c) 11H-

p+

n0 n0 n0

e- e- e-

(d) 92238U p+ (e) 11

23Na

+ p+ (f) 16

32S2-

p+

n0 n0 n0

e- e- e-

2. Give one example of (a) an isotope of carbon-12, (b) an ion that is isoelectronic with

neutral carbon-12, (c) an atom that has the same number of protons as carbon-12, (d) an

atom that has the same number of neutrons as carbon-12, and (e) an atom that has the

same number of nucleons as carbon-12.

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3. Write the complete symbol ( chgAZ E ) for an atom that contains 15 protons, 17 neutrons and

15 electrons. (Write the correct atomic symbol in your answer, not E!!)

4. Write the complete symbol ( chgAZ E ) for an atom that contains 1 proton, no neutrons and

no electrons. (Write the correct atomic symbol in your answer, not E!!)

5. Write the complete symbol ( chgAZ E ) for an atom that contains 82 protons, 125 neutrons

and 80 electrons. (Write the correct atomic symbol in your answer, not E!!)

6. Write the electronic configuration for each element.

(a) lithium _______________________________________________

(b) carbon _______________________________________________

7. Name the element with each of the following electronic configurations.

(a) [He] 2s2 ______________________________

(b) [Ne] 3s23p1 ______________________________

(c) [Ar] 4s1 ______________________________

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Periodic Table Homework

Name ___________________________

1. Distinguish between a group and a period on the periodic table.

2. Give one example of each of the following:

alkali metal =

halogen =

noble gas =

metalloid =

alkaline earth metal =

transition metal =

lanthanide =

an element used by the thyroid gland to make thyroxine =

an element whose forms include diamond, graphite and buckminsterfullerene =

an element in the p-block of the periodic table that is not a halogen, noble gas or

metalloid =

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3. How many electrons can the first energy level (or shell) in an atom hold?

4. How many electrons can the fifth energy level in an atom hold? (hint: use 2n2)

5. How many valence electrons are there in a neutral oxygen atom?

6. How many valence electrons are there in a neutral sodium atom?

7. How many valence electrons are there in a neutral carbon atom?

8. How many valence electrons are there in a potassium ion, K+?

9. How many valence electrons are there in a calcium ion, Ca2+?

10. Which group of atoms has a stable octet of valence electrons?

11. Distinguish between an s-orbital and a p-orbital.

12. Comparing C and F, which atom has the larger radius? Which has the higher ionization

energy?

13. Comparing N and P, which atom has the larger radius? Which has the higher ionization

energy?

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Chemical Bonding Homework

Name ___________________________

1. Draw the electron-dot symbol for each of the following atoms:

arsenic, As iodine, I

rubidium, Rb beryllium, Be

gallium, Ga silicon, Si

helium, He radon, Rn

oxygen, O nitrogen, N

2. Draw electron-dot symbols for the ions in each of the following binary ionic compounds:

sodium chloride, NaCl magnesium oxide, MgO

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3. Explain the octet rule in your own words.

4. Briefly explain why hydrogen does not follow the octet rule.

5. Distinguish between ionic and covalent bonds.

6. What is a polyatomic ion? What makes it different from a molecule? What makes it

different from a monatomic ion? (Include an example of each in your answer.)

7. Draw a valid Lewis structure for each of the following molecules:

oxygen, O2 iodine, I2

water, H2O methane, CH4

phosphorus trichloride, PCl3 oxygen dichloride, OCl2

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Naming and Formula Writing Homework

Name ___________________________

Name each of the following compounds:

H2CO3 ______________________________________________

NaHCO3 ______________________________________________

CO2 ______________________________________________

SCl2 ______________________________________________

Zn3(PO4)2 ______________________________________________

AgI ______________________________________________

TiF4 ______________________________________________

Cr(ClO3)3 ______________________________________________

Al2S3 ______________________________________________

H2O ______________________________________________

NBr3 ______________________________________________

Fe(OH)2 ______________________________________________

CBr4 ______________________________________________

H2SO4 ______________________________________________

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Write the correct formula for each of the following compounds:

hydrochloric acid ______________________________

magnesium chloride ______________________________

aluminum phosphate ______________________________

ammonium hydroxide ______________________________

ammonium carbonate ______________________________

dinitrogen trioxide ______________________________

dinitrogen tetroxide ______________________________

dinitrogen pentoxide ______________________________

sulfur dioxide ______________________________

sulfuric acid ______________________________

potassium bicarbonate ______________________________

ammonium cyanide ______________________________

calcium hydrogen phosphate ______________________________

iron (III) cyanide ______________________________

silver oxide ______________________________

nitric acid ______________________________

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Molecular Geometry Homework

Name ___________________________

Molecular

Formula

Number

of v.e.

Lewis Structure

Shape Polar or

Nonpolar?

H2O

NH3

CH4

CH2F2

HCN

CO2

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Molecular

Formula

Number

of v.e.

Lewis Structure

Shape Polar or

Nonpolar?

Cl2O

BF3

H2CO (formaldehyde,

C in the center)

NF3

H2S

C2H4 (ethylene,

H2C=CH2)

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Balancing Equations Homework

Name ___________________________

Balance each of the following equations by placing an integer coefficient on each of the blank

lines:

___ KNO3 ----------> ___ KNO2 + ___ O2

___ NH4NO2 ----------> ___ N2 + ___ H2O

___ MgO + ___ H2O ----------> ___ Mg(OH)2

___ Fe2O3 + ___ CO ----------> ___ Fe + ___ CO2

___ CaO + ___ P2O5 ----------> ___ Ca3(PO4)2

___ MgCl2 + ___ AgNO3 ----------> ___ Mg(NO3)2 + ___ AgCl

___ Cu + ___ H2SO4 ---------> ___ CuSO4 + ___ SO2 + ___ H2O

___ K2CO3 + ___ BaCl2 ----------> ___ KCl + ___ BaCO3

___ Mg(OH)2 + ___ (NH4)3PO4 ----------> ___ Mg3(PO4)2 + ___ NH3 + ___ H2O

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Note: Be very careful about writing chemical formulas! If your formulas are wrong, you may

come up with an equation that is impossible to balance!!

methane + oxygen --------> carbon dioxide + water

magnesium oxide + carbon dioxide --------> magnesium carbonate

magnesium oxide --------> magnesium + oxygen

water + carbon dioxide --------> carbonic acid

water + diphosphorus pentoxide --------> phosphoric acid

water + sodium oxide --------> sodium hydroxide

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Mole Homework

Name ______________________________

1. Calculate the number of moles of O2 in 100.0 g of O2 gas.

2. Calculate the number of O2 molecules in 100.0 g of O2 gas.

3. Calculate the number of moles of Ti in 38.2 g of titanium metal.

4. Calculate the number of Ti atoms in 38.2 g of titanium metal.

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5. Calculate the mass of 2.08 moles of CO2.

6. Calculate the mass of 0.064 moles of C6H12O6.

7. Calculate the mass of 1 hundred trillion water molecules. (By definition, 1 hundred trillion = 1 x 1014.) Do you think you could see a sample of water this small?

8. Proteins can have extremely high molecular weights, up into the thousands and millions of grams/mole. Calculate the mass of 1 hundred trillion protein molecules if the protein

has a molecular weight of 10,000,000 g/mol. Do you think you could see a sample of

protein this small?

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Stoichiometry Homework

Name ___________________________

1. Balance the following chemical equation, and use it to answer the following questions:

_____ Al (s) + _____ O2 (g) -----> _____ Al2O3 (s)

(a) What mass of oxygen will react with 2.75 g of aluminum?

(b) What mass of aluminum oxide will be formed when 2.75 g of aluminum reacts

completely with oxygen?

(c) Show how your results are consistent with the Law of Conservation of Mass.

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2. Hydrogen gas reacts with oxygen gas to form liquid water. How many grams of O2 are

required in order to produce 10.0 g of H2O?

3. Sodium metal reacts with chlorine gas to form sodium chloride. How many grams of

sodium chloride can be produced from 10.0 g of Cl2?

4. Methane reacts with oxygen to form carbon dioxide and water. How many grams of

oxygen are required in order to react completely with 10.0 g methane?

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Energy Homework

Name ___________________________

1. Use Table 6.2 on p. 171 to help you answer these questions.

(a) Which bond is stronger, H-H or F-F?

(b) Which bond is stronger, H-H or H-F?

2. (a) Bond __________________ is always endothermic. (making or breaking)

(b) Bond __________________ is always exothermic. (making or breaking)

3. In the reaction H2 + F2 2HF, one H-H bond is broken, one F-F bond is broken and two

H-F bonds are formed. Is the reaction exothermic or endothermic? Explain your

reasoning. (Use the bond dissociation energies in Table 6.2 to help you answer this

question.)

4. (a) Slow reactions have __________________ activation energies. (high or low)

(b) Fast reactions have __________________ activation energies. (high or low)

5. Catalysts speed up chemical reactions by ________________ the activation energy.

6. When the forward and reverse rates of a chemical reaction are equal, we say that the

reaction is at _________________________.

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7. Distinguish between the terms endothermic reaction and exothermic reaction.

8. Distinguish between the terms potential energy and kinetic energy.

9. Distinguish between the terms calorie and Joule.

10. According to Le Chatelier’s Principle, adding reactants (or removing products) drives the

equilibrium to the __________, adding products (or removing reactants) drives the

equilibrium to the __________, increasing temperature favors the ___________________

reaction, decreasing temperature favors the ____________________ reaction, increasing

pressure shifts the reaction toward _______________ gas molecules, decreasing pressure

shifts the reaction toward ________________ gas molecules, and adding a catalyst has no

effect on equilibrium position whatsoever. (Choices: left/right, endothermic/exothermic,

more/fewer)

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Reaction Rate Homework

Name ___________________________

1. According to molecular collision theory, only collisions that have sufficient energy and

proper orientation lead to a reaction.

(a) In your own words, explain why sufficient energy is necessary for a reaction to occur.

(b) In your own words, explain why proper orientation is necessary for a reaction to occur.

2. The transition state is the point of ________________ energy in a molecular collision.

(maximum/minimum)

3. Sketch a labeled energy diagram for an endothermic reaction with H = 10 kcal/mol and

and Ea = 30 kcal/mol. Be sure to label the axes (energy and reaction coordinate), the

reactants, products, transition state, H and Ea.

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4. For a given chemical reaction, state whether each of the following changes will result in a

faster or a slower reaction (circle your answer). Then, explain each answer in terms of

molecular collision theory (collision rates, collision energies and activation energy).

(a) The concentrations of reactants is decreased. Faster Slower

(b) The temperature is decreased. Faster Slower

(c) A catalyst is added. Faster Slower

(d) The volume of a gas-phase reaction is decreased. Faster Slower

(e) Any solid reactants are crushed. Faster Slower

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Gas Law Homework

Name ___________________________

Ideal Gas Law: PV = nRT, where R = 0.08206 Latm/molK

Boyle's Law: P1V1 = P2V2 (n, T constant);

Charles' Law: V1/T1 = V2/T2 (n, P constant)

Gay-Lussac’s Law: P1/T1 = P2/T2 (n, V constant)

Avogadro’s Law: V1/n1 = V2/n2 (T, P constant)

1. What is the volume of 3.25 moles of ideal gas at 25.0C and 0.759 atm pressure?

2. What pressure is exerted by 0.25 moles of ideal gas at 50.0C and a volume of 1.75 L?

3. How many moles of ideal gas will occupy a volume of 20.0 L at 100.C and 0.275 atm?

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4. At what temperature (in C) will 1.25 moles of ideal gas occupy a volume of 10.0 L with

a pressure of 0.500 atm?

5. A fixed quantity of gas at 25.00C occupies 3.75 L. If the gas is heated at constant

pressure to 50.00C, what will its new volume be?

6. A fixed quantity of gas at 2.75 atm occupies 3.75 L. If the gas is compressed at constant

temperature to 1.75 L, what will its new pressure be?

7. A 1.25-mol sample of gas occupies 25.0 L at a certain temperature and pressure. What

volume will a 2.25-mol sample occupy at the same temperature and pressure?

8. A fixed quantity of gas at 300.0 K exerts a pressure of 1.45 atm. If the gas is heated at

constant volume to 350.0 K, what will its new pressure be?

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Intermolecular Forces Homework

Name ___________________________

______ 1. Which of the following molecules exhibits hydrogen bonding?

(A) CH4 (B) PH3 (C) H2S (D) H2O

______ 2. Which of the following are intermolecular forces?

(A) dipole-dipole interactions (B) London forces

(C) hydrogen bonding (D) all of the above

______ 3. Which of the following will tend to increase intermolecular forces?

(A) increased molecular weight (B) increased polarity

(C) increased hydrogen bonding (D) all of the above

______ 4. Choose the substance with the highest boiling point:

(A) H2O (B) Ne (C) O2 (D) Ar

______ 5. Choose the substance with the highest boiling point:

(A) F2 (B) Cl2 (C) Br2 (D) I2

______ 6. Which state of matter has both definite shape and definite volume?

(A) solid (B) liquid (C) gas (D) none

______ 7. Which state of matter has neither definite shape nor definite volume?


______ 8. Which state of matter has definite shape but not definite volume?


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9. Can nonpolar molecules experience London dispersion forces? Why or why not?

10. Which types of intermolecular forces are present in HCN? Explain your reasoning.

11. (a) Which of the phase changes are endothermic?

(b) Which of the phase changes are exothermic?

12. Higher vapor pressures are found in liquids with __________________ temperatures and

____________________ intermolecular forces. (higher/lower, stronger/weaker)

13. Distinguish between the terms viscosity and surface tension.

14. Distinguish between the bonding in ionic, molecular, network and metallic solids.

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Solutions Homework

Name ___________________________

1. A solution is prepared by dissolving 10.0 g of (solid) naphthalene (C10H10) in 500.0 mL

of (liquid) chloroform (CHCl3). The resulting solution is liquid. Identify the solute and

the solvent.

2. Calculate the molarity of the solution from #1.

3. Calculate the weight/volume percent concentration of C10H10 in the solution from #1.

4. What volume of the solution from #1 contains 2.7 g of naphthalene?

5. How many grams of naphthalene are contained in 25.0 mL of the solution from #1?

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Remember that for dilution calculations c1V1 = c2V2.

6. Calculate the final concentration obtained when 150.0 mL of a 0.757 M solution of CaCl2

is diluted to a final volume of 1.000 L.

7. What volume of 6.00 M HCl should be diluted to a final volume of 500.0 mL in order to

obtain a 1.50 M solution of HCl?

8. Determine the number of “particles” contained in one mole of each solute, and then rank

the solutions in order of increasing vapor pressure.

0.10 M sucrose (C12H22O11) 0.08 M NaCl 0.03 M MgCl2

9. Rank the solutions from #8 in order of increasing boiling point.

10. Rank the solutions from #8 in order of increasing freezing point.

11. Rank the solutions from #8 in order of increasing osmotic pressure.

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Acids & Bases Homework

Name___________________________________

1. Write the formula for the conjugate base of each of the following acids:

HC2H

3O

2 __________ H

2O __________

HCO3− __________ H3PO4 __________

NH4+ __________ H3O+ __________

2. Write the formula for the conjugate acid of each of the following bases:

NH3 __________ H

2O __________

HCO3− __________ F− __________

PO43− __________ OH− __________

3. Give one example each of a strong acid, a strong base, a weak acid and a weak base.

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4. How many grams of H2SO

4 are present in 250.0 mL of 0.175 M solution?

5. What volume of 0.1073 M NaOH contains 0.0250 moles of NaOH?

6. What is the pH of 0.100 M HCl?

7. What is the pH of 0.100 M NaOH?

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Neutralization and Buffer Homework

Name___________________________________

Write the net ionic equation for each of the following neutralization reactions.

1. HCl reacts with NaOH in a neutralization reaction.

2. H2SO4 reacts with NaOH in a neutralization reaction.

3. HCl reacts with NaHCO3 in a neutralization reaction.

4. H2SO4 reacts with NaHCO3 in a neutralization reaction.

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5. State whether each of the following salts is acidic, basic or neutral.

sodium chloride ________________ sodium acetate _________________

ammonium chloride ______________ potassium nitrate _______________

sodium phosphate _______________ lithium fluoride ________________

For questions #6 and 7, use the equation [H3O+] = Ka[HA]/[A−] or pH = pKa – log[HA]/[A−].

6. Find the pH of a buffer that contains 0.15 M acetic acid and 0.10 M sodium acetate.

7. Find the pH of a buffer that contains 0.20 M sodium carbonate and 0.10 M sodium

bicarbonate.

8. What is the property of a buffer that makes it different from a non-buffer solution?

9. Why are buffers important in the environment and in the human body?

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