chemistry 11 final examination review -...

Chemistry 11 Final Examination Review - Answers

Part A - True or False. Indicate whether each of the following statements is true or false. Correct the false statements.

T 1. Dalton's atomic theory includes a statement that says atoms of the same element are identical. F 2. Dalton's theory states that atoms have variable masses.

Dalton’s atomic theory states that atoms of the same element have the same mass and that atoms of different elements have different masses.

F 3. J.J. Thomson is credited with the discovery of protons. J.J. Thomson is credited with the discovery of electrons.

F 4. The mass of an electron is equal to the mass of a proton. The mass of an electron is less than the mass of a proton. T 5. The mass of a proton is approximately equal to the mass of a neutron. T 6. The atomic number represents the number of protons in a nucleus. T 7. The proton has a mass of approximately 1 u. F 8. The difference in mass of isotopes of the same element is due to the different number of protons in the nucleus.

The difference in mass of isotopes of the same element is due to the different number of neutrons in the nucleus.

T 9. The isotope carbon-12 is used as the relative mass standard for the atomic mass scale. F 10. The mass of the most common isotope of each element is listed on the periodic table.

The average mass of the naturally occurring isotopes of each element is listed on the periodic table.

Part B - Multiple Choice

D 1. A(n) __ is used to represent a compound. a) symbol b) equation c) subscript d) formula A 2. ___ atoms or groups of atoms are called ions. a) charged b) diatomic c) neutral d) monatomic A 3. For the formula of a compound to be correct, the algebraic addition of the charges on the atoms or ions in the compound must add up to __. a) zero b) one c) two d) four D 4. Potassium bromide is an example of a(n) __ compound. a) molecular b) organic c) polyatomic d) ionic B 5. The only common polyatomic ion that has a positive charge is the __ ion. a) phosphate b) ammonium c) sulfate d) nitrate A 6. In the formula H2SO4, the number 4 would be called a(n) __. a) subscript b) oxidation number c) coefficient d) charge B 7. Which subatomic particle contributes the least to the mass of an atom? a) nucleon b) electron c) proton d) neutron

Chemistry 11 – Final Examination Review – Answers – of 27

D 8. In the free, or uncombined, state the number of protons in the nucleus of an element must equal the __.

a) mass number c) mass number - atomic number b) number of neutrons in the nucleus d) number of electrons present

C 9. Which of the following ideas of the Bohr model is not retained in the modern theory of atomic structure?

a) Electrons can absorb or emit energy only in whole numbers of photons. b) Atoms have a central positively charged nucleus. c) Electrons move around the nucleus as planets orbit the sun. d) Most of the volume of an atom is empty space.

B 10. Which of the following orbitals is spherical in shape? a) 3p b) 2s c) 4d d) 5f B 11. The third energy level of an atom may have __ electrons. a) 2 b) 18 c) 8 d) 32 C 12. How many sublevels are possible at the fourth energy level? a) 2 b) 3 c) 4 d) 18 A 13. Lustrous, malleable, ductile elements that are good conductors of electricity and heat are classified as __.

a) metals b) nonmetals c) metalloids d) noble gases C 14. The electron configuration of a certain element ends with 3p5. Which of the following

describes its position in the periodic table? a) period 5, group 13 b) period 3, group 15 c) period 3, group 17 d) period 5, group 15 B 15. Which of the following is an example of a metalloid?

a) iodine b) boron c) bromine d) indium B 16. The periodicity of the elements is basically a function of their __. a) nuclear stability b) atomic numbers c) mass numbers d) none of these D 17. An element with seven electrons in the outer level would be a __.

a) metal b) metalloid c) noble gas d) nonmetal A 18. As the atomic number in a period increases, the degree of nonmetallic character __.

a) increases c) increases then decreases b) decreases d) remains the same

B 19. Elements in a group have similar chemical properties because of their similar __. a) nuclear configurations c) mass numbers

b) outer electron configurations d) names D 20. The period number in the periodic table designates the __ for the row.

a) total nuclear charge c) maximum number of outer electrons b) maximum number of nucleons d) highest energy level

B 21. The radii of the atoms become smaller from sodium to chlorine across period 3. This is primarily a result of __.

a) the shielding effect c) the increased number of electrons b) increased nuclear charge d) decreased metallic character

D 22. Which of the following would correctly characterize a nonmetal? a) low first ionization energy, low electron affinity b) high first ionization energy, low electron affinity c) low first ionization energy, high electron affinity d) high first ionization energy, high electron affinity


D 23. Compared to the stability of the original atom, the stability of its ion that resembles a noble gas configuration would be __.

a) identical b) sometimes less c) less d) greater B 24. Which of the following does NOT affect the ionization energy of an electron?

a) the distance of the electron from the nucleus c) the number of energy levels in the atom b) the number of neutrons in the nucleus d) the charge on the nucleus

A 25. For each subsequent electron removed from an atom the ionization energy required __. a) increases b) decreases c) remains the same d) no pattern

C 26. The formation of bonds between atoms depends on __. a) the electron configurations of the atoms involved b) the attraction the atoms have for electrons

c) both of the preceding factors d) neither of the preceding factors

D 27. The particle that results when two or more atoms form covalent bonds is a __. a) single charged atom b) molecule c) polyatomic ion d) b or c

A 28. Compounds that have low melting points, are brittle, and do not conduct electricity are probably __. a) covalent c) ionic b) metallic d) compounds of polyatomic ions

A 29. The most active __ have the highest electronegativities. a) nonmetals b) metalloids c) metals d) noble gases D 30. __ compounds have high melting points, conduct electricity in the molten phase, and tend to be

soluble in water. a) hydrogen b) metallic c) covalent d) ionic A 31. The element in the following group that has the lowest electronegativity is __. a) potassium b) arsenic c) bromine d) chromium D 32. If there are only two electron pairs in the outer energy level of an atom in a molecule, they will be

found __. a) at 90º to one another c) at 120º to one another b) on the same side of the nucleus d) on opposite side of the nucleus

B 33. The __ molecule has two bonding pairs and two unshared pairs of electrons. a) CH4 b) H2O c) NH3 d) HF

D 34. A certain atom contains 34 protons, 34 electrons, and 45 neutrons. This atom has a mass number of __. a) 34 b) 45 c) 68 d) 79

C 35. An example of a compound is a) oxygen b) mercury c) salt d) diamond

A 36. Carbon is classed as an element rather than as a compound because it a) cannot be chemically decomposed into two or more substances b) has been known for many centuries

c) is formed when wood is heated out of contact with air d) combines with oxygen to form a gas

A 37. The positively charged particles in the nucleus of an atom are called a) protons b) neutrons c) electrons d) ions

C 38. The charged particles that are found outside the nucleus of an atom are called a) protons b) ions c) electrons d) mesons


D 39. A nonmetallic atom generally becomes a negative ion by a) losing protons b) losing electrons c) gaining protons d) gaining electrons

A 40. Which of the following subatomic particles has the smallest mass? a) electron b) neutron c) proton d) nucleus A 41. The nuclear atom was first postulated by a) Rutherford b) Bohr c) Dalton d) Thompson B 42. Which of the following symbols represents an atom that contains the largest number of neutrons?

a) 235 92U b) 239

92U c) 239 93Np d) 239

94Pu

C 43. The nuclide symbol 16 8O represents an oxygen atom with a) a mass of 8 u c) a mass of 16 u b) an atomic number of 16 d) 16 neutrons

C 44. If Z represents the atomic number of an element and A represents the mass number, then the number of neutrons in one atom is

a) A b) A + Z c) A - Z d) Z - A D 45. Which of the following statements about the elemental species 24

11X and 2512Z is correct?

a) They are isotopes of the same element. b) They are nonmetals. c) They are members of the same chemical family. d) They have the same number of neutrons per atom.

B 46. An element X has a mass number of 32 and an atomic number of 16. The most common ion of element X is represented by a) X+ b) X2- c) X- d) X2+

Part C - Short Answer 1. What is the maximum number of electrons that may occupy one orbital? two 2. The Lewis electron dot diagram is used to represent only which electrons in the atom? Valence electrons 3. What is the diagonal rule used to predict? The order in which orbitals are filled 4. How many sublevels are possible at the third energy level? three 5. How many orbitals are there in the f sublevel? seven 6. What is the maximum number of electrons that can occupy a d sublevel? ten 7. Which sublevel may contain a maximum of three pairs of electrons? ‘p’ sublevel


8. What must be true about the spins of two electrons occupying the same orbital? They must have opposite spin 9. Write the electron configuration for each of the following elements:

a) lithium 1s22s1

b) radium 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s2

c) sodium 1s22s22p63s1

d) mercury 1s22s22p63s23p64s23d104p65s24d105p66s24f145d10

10. Draw the energy level diagram for each of the following elements: a) tin

6s ↑ ↑ 5p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

↑↓ 4d 5s ↑↓ ↑↓ ↑↓ 4p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

↑↓ 3d 4s ↑↓ ↑↓ ↑↓

↑↓ 3p 3s ↑↓ ↑↓ ↑↓

↑↓ 2p 2s

↑↓ 1s


10. Draw the energy level diagram for each of the following elements: b) krypton

5s ↑↓ ↑↓ ↑↓ 4p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

↑↓ 3d 4s ↑↓ ↑↓ ↑↓

↑↓ 3p 3s ↑↓ ↑↓ ↑↓

↑↓ 2p 2s

↑↓ 1s

c) gold

6p ↑↓ ↑↓ ↑↓ ↑↓ ↑ 5d ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

↑↓ 4f 6s ↑↓ ↑↓ ↑↓ 5p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

↑↓ 4d 5s ↑↓ ↑↓ ↑↓ 4p ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

↑↓ 3d 4s ↑↓ ↑↓ ↑↓

↑↓ 3p 3s ↑↓ ↑↓ ↑↓

↑↓ 2p 2s

↑↓ 1s


10. Draw the energy level diagram for each of the following elements: d) potassium

↑ 3d 4s ↑↓ ↑↓ ↑↓

↑↓ 3p 3s ↑↓ ↑↓ ↑↓

↑↓ 2p 2s

↑↓ 1s

11. Write the energy level population for each of the following elements: a) calcium 2, 8, 8, 2 b) sulfur 2. 8. 6 c) scandium 2, 8, 9, 2 d) tungsten 2, 8, 18, 32, 12, 2 12. Element X has the following configuration: 1s22s22p63s22p64s23d104p65s1

a) What period is this element located in? five

b) What group is this element in? one (alkali metals)

c) Identify the element. Rubidium (Rb)

d) Is the element a metal, nonmetal, or metalloid? Metal 13. List five properties of metals and five properties of nonmetals.

Metals - good conductors of heat and electricity - lustre - malleable and ductile - high densities - high boiling points and melting points - resist stretching and twisting - solids at room temperature

Nonmetals - no lustre - poor conductors - may be solid, liquid or gas - low densities - low melting and boiling points


14. In a group, do the radii of atoms increase or decrease as the atomic number increases? Explain why.

In a group, the radii increases as the atomic number increases because electrons are being added at higher energy levels which are located further from the nucleus producing a larger atom.

15. Does the radii of atoms within a period increase or decrease as the atomic number increases? Explain why.

The radii of atoms in a period decrease as the atomic number increases. As you go across (left to right) a period the electrons are being added to the same energy level while more protons are added to the nucleus. The attractive force is greater and the electrons are pulled closer to the nucleus producing a smaller atom.

16. In each of the following pairs of atoms, pick the one that is larger. a) Mg, Na b) K, Ca c) Al, B d) Br, Cl e) F, N f) Ne, Ar 17. In a group, will the ionization energy tend to increase or decrease with increasing atomic number?

Explain why. In a group the ionization energy decreases as the atomic number increases. As you go down a group the size of the atom increases causing the electrons to be located further away from the nucleus. The attraction between the nucleus (+) and electrons (-) which decreases as the distance between them increases, which makes it easier to remove an electron.

18. In a period, will the ionization energy tend to increase or decrease with increasing atomic number? Explain why. In a period the ionization energy increases as the atomic number increases. As you go across a period (left to right) the size of the atom decreases, the electrons are located closer to the nucleus with a greater attraction between the electrons and the nucleus. Because the attraction is greater, more energy is required to remove an electron.

19. Do metals or nonmetals generally have lower ionization energies? Metals generally have lower ionization energies than nonmetals. 20. Carbon has a first ionization energy of 1086.5 kJ/mol. Predict whether the first ionization energies of the

following elements will be more or less than that of carbon. a) helium higher ionization energy b) lithium lower ionization energy c) fluorine higher ionization energy d) silicon lower ionization energy


21. As the distance between the nucleus and the outer electrons of an atom increases, will the ionization energy increase or decrease? As the distance between the outer electrons and the nucleus increases, the attractive force decreases which causes the ionization energy to decrease.

22. As the positive charge on an ion increases, will the ionization energy increase or decrease? As the positive charge on the ion increases the electrons are held more tightly and the ionization energy increases.

23. How are substances that are gases or soft solids at room temperature classified?

Substances that are gases or soft solids at room temperature are classified as covalent compounds.

24. Would an element with two outer electrons be a metal or a nonmetal?

An element with two outer electrons would be a metal. 25. What Russian scientist designed the first periodic table?

Mendeleev designed the first periodic table. 26. Which group of elements has eight outer electrons?

The noble gases have eight outer electrons. 27. For the transition elements, as the atomic number increases, to which sublevel are the electrons

being added? For the transition metals, the electrons are being added to the “d” sublevel.

28. According to the octet rule, how many pairs of outer electrons do the most stable atoms have?

According to the octet rule, the most stable atoms have four pairs of outer electrons.

29. In the lanthanide series, as the atomic number increases, to which sublevel are electrons being

added? In the lanthanide series, electrons are being added to the “f” sublevel.

30. Which atom in each of the following pairs would have the lower first ionization energy? a) N, O b) Te, Se c) C, Ge d) Br, I e) I, Sb f) Al, N


31. In a group, will the electron affinity tend to increase or decrease with increasing atomic number? Explain why. In a group, electron affinity decreases as atomic number increases. As you go down the group, the size of the atom increases, electrons are added further from the nucleus where the attraction is less, therefore less energy is released when an electron is added to the atom, and electron affinity is smaller.

32. In a period, will the electron affinity tend to increase or decrease with increasing atomic number?

Explain why. In a period, the electron affinity increases as the atomic number increases. As you go across (left to right) the size of the atom decreases and the electrons are added closer to the nucleus where the attraction is greater – causing more energy to be released when the electron is added, increasing the electron affinity. (Noble gases do not follow this trend because of their increased stability, they do not gain electrons easily.)

33. Which atom in each of the following pairs would have the higher electron affinity? a) S, Cl b) At, Rn c) Y, Os d) Ga, In e) Br, I f) O, N 34. Give examples of molecules with the following shapes:

a) linear HCl, CO2, BeCl2 b) trigonal planer BH3, BCl3, AlBr3 c) trigonal pyramidal NH3, NH3 d) tetrahedral CH4, CCl4 e) bent H2O, H2S

35. Use electron dot formulas to show the complete balanced bonding reactions between the following elements. Indicate the type of bond expected. a) sodium and oxygen

b) nitrogen and hydrogen

c) zirconium and sulfur

d) magnesium and chlorine


36. Write the correct names for the following chemical compounds. a) HCl(aq) hydrochloric acid b) KOH potassium hydroxide c) HgOH mercury(I) hydroxide d) FeCl3 iron(III) chloride e) Al2(SO4)3 aluminum sulfate f) N2O5 dinitrogen pentoxide g) HF hydrofluoric acid h) Pb(OH)2 lead(II) hydroxide i) NH4NO3 ammonium nitrate j) NaHCO3 sodium bicarbonate k) Zn(NO2)2 zinc nitrite l) H3PO4(aq) phosphoric acid

m) SO3 sulfur trioxide n) NaC2H3O2 sodium acetate o) HFO2(aq) flourous acid p) Al(BrO4)3 aluminum perbromate q) SbF3 antimony(III) fluoride r) Pd(CN)2 palladium(II) cyanide s) Ca(MnO3)2 calcium manganate t) Be(NO4)2 beryllium pernitrate u) NiSeO4 nickel(II) selenate v) H2SO3(aq) sulfurous acid w) Ba(OH)2 barium hydroxide x) PbS lead(II) sulfide

37. Write the correct chemical formula for each compound and balance the equation.

a) sodium carbonate + calcium hydroxide sodium hydroxide + calcium carbonate Na2CO3 + Ca(OH)2 2NaOH + CaCO3

b) carbon dioxide + water carbonic acid

CO2 + H2O H2CO3

c) phosphorus + oxygen phosphorus pentoxide 2P + 5O2 2PO5

d) sodium + water sodium hydroxide + hydrogen

2Na + 2H2O 2NaOH + H2

e) zinc + sulfuric acid zinc sulfate + hydrogen Zn + H2SO4 ZnSO4 + H2

f) aluminum sulfate + calcium hydroxide aluminum hydroxide + calcium sulfate

Al2(SO4)3 + 3Ca(OH)2 2Al(OH)3 + 3CaSO4 g) calcium oxide + water calcium hydroxide

CaO + H2O Ca(OH)2 h) iron + copper(I) nitrate iron(II) nitrate + copper

Fe + 2CuNO3 Fe(NO3)2 + 2Cu i) iron(II) sulfide + hydrochloric acid hydrogen sulfide + iron(II) chloride

FeS + 2HCl H2S + FeCl2 j) potassium oxide + water potassium hydroxide

K2O + H2O 2KOH



k. carbon + ferric oxide iron + carbon dioxide

3 C + 2 Fe2O3 4 Fe + 3 CO2

l. sulfur tetrafluoride + water sulfur dioxide + hydrofluoric acid

SF4 + 2 H2O SO2 + 4 HF

m. calcium hydroxide + phosphoric acid calcium phosphate + water

3 Ca(OH)2 + 2 H3PO4 Ca3(PO4)2 + 6 H2O

n. ethane + oxygen carbon dioxide + water

2 C2H6 + 7 O2 4 CO2 + 6 H2O

o. aluminum sulfate + ammonia + water aluminum hydroxide + ammonium sulfate

Al2(SO4)3 + 6 NH3 + 6 H2O 2 Al(OH)3 + 3 (NH4)2SO4 38. Identify the type of reaction. Write the correct chemical formulas of the compounds, complete and

balance the equations and name the products formed. a) calcium carbonate + hydrochloric acid calcium chloride + carbonic acid

CaCO3 + 2HCl CaCl2 + H2CO3

double displacement b) iron + sodium bromide

Fe + NaBr no reaction b/c iron is less active than sodium single displacement

c) ammonium acetate + iron(II)chloride ammonium chloride + iron(II) acetate

2NH4C2H3O2 + FeCl2 2NH4Cl + Fe(C2H3O2)2

double displacement d) silver bromide + ammonium sulfate silver sulfate + ammonium bromide

2AgBr + (NH4)2SO4 Ag2SO4 + 2NH4Br double displacement

e) zinc + sulfuric acid hydrogen + zinc sulfate

Zn + H2SO4 H2 + ZnSO4

single displacement f) neon + potassium

Ne + K no reaction b/c noble gases are unreactive combination

g) lead(II )hydroxide + hydrochloric acid lead(II) chloride + water

Pb(OH)2 + 2HCl PbCl2 + 2H2O double displacement

h) iron + sulfur iron(II) sulfide

Fe + S FeS combination


i) potassium chlorate (heated) potassium chloride + oxygen

2KClO3 2KCl + 3O2

decomposition

j) calcium oxide + water calcium hydroxide CaO + H2O Ca(OH)2

combination

11. dinitrogen pentoxide + water nitric acid N2O5 + H2O 2HNO3

combination 12. carbon dioxide + water carbonic acid

CO2 + H2O H2CO3

combination 13. chlorine + chromium(III) bromide bromine + chromium(III) chloride

3 Cl2 + 2 CrBr3 3 Br2 + 2 CrCl3Single displacement

14. sulfur + oxygen sulfur dioxide or sulfur trioxide

S + O2 SO2 or S + O2 SO3

combination

15. zinc + hydrochloric acid zinc chloride + hydrogen

Zn + 2 HCl ZnCl2 + H2

Single displacement 16. sodium + water sodium hydroxide + hydrogen

2 Na + 2 H2O 2 NaOH + H2

single displacement 17. magnesium + water

Mg + H2O no reaction – Mg is not active enough to replace H in water

Single displacement 18. copper + stannic nitrate

Cu + Sn(NO3)4 no reaction – Cu is less active than Sn

Single displacement 19. aluminum + cupric sulfate copper + aluminum sulfate

2 Al + 3 CuSO4 3 Cu + Al2(SO4)3

Double Displacement


Part D – Calculations Show all work. Express your answers using significant digits and include units with your answers. 1. Calculate the average atomic mass of the following elements: a) Mg-24 mass = 23.985 u 78.70% Mg-25 mass = 24.986 u 10.13% Mg-26 mass = 25.983 u 11.17% (0.7870)(23.985) + (0.1013)(24.986) + (0.1117)(25.983) = 24.31 u b) Ir-191 mass = 191.0 u 37.58% Ir-193 mass = 193.0 u 62.42% (0.3758)(191.0) + (0.6242)(193.0) = 192.2 u 2. Calculate the percentage of each of the isotopes of silver if silver-107 has a mass of 106.905 u and silver-109 has a mass of 108.905 u and the average atomic mass of silver is 107.869 u 107Ag = x, 109Ag = y x + y = 1

106.905x + 108.905y = 107.869 -106.905x - 106.905y = -106.905 use elimination or substitution 106.905x + 108.905y = 107.869 2y = 0.964 y = 0.482 x = 0.518 107Ag = 51.8%, 109Ag = 48.2%


*3. Naturally, occurring silicon consists of three isotopes, 28Si, 29Si, and 30Si, whose atomic masses are 27.9769, 28.9765, and 29.9738, respectively. The most abundant isotope is 28Si, which accounts for 92.23 percent of naturally occurring silicon. Given that the observed atomic mass of silicon is 28.0855, calculate the percentages of 29Si and 30Si in nature.

Isotope 28Si 29Si 30Si Average

Mass 27.9769 g 28.9765 g 29.9738 g 28.0855 Percentage 92.23% x y

(eqn 1) (0.9223)(27.9769) + (28.9765)(x) + (29.9738)(y) = 28.0855 (eqn 1) 25.8031 + 28.9765x + 29.9738y = 28.0855

(eqn 1) 28.9765x + 29.9738y = 2.2824 (eqn 2) 0.9223 + x + y = 1 (eqn 2) x + y = 0.0777 use substitution or elimination (eqn 1) 28.9765x + 29.9738y = 2.2824 (-28.9756)(eqn 2) -28.9765x – 28.9765y = -2.2824 y = 0.03093 0.9973y = 0.03093 0.9973 y = 0.0310, x = 0.0467, ∴ 29Si = 3.10% and 30Si = 4.67% 4. Complete the chart below:

Element Atomic # Mass # Protons Neutrons Electrons

calcium-43 20 43 20 23 20

lead-211 82 211 82 129 82

plutonium-242 94 242 94 148 94

chromium-50 24 50 24 26 24

65Cu2+ 29 65 29 36 27

34S2- 16 34 16 18 18

128 53I- 53 128 53 75 54

208 82Pb4+ 82 208 82 126 78

5. Convert each of the following to moles.

a. 8.8 g of potassium carbonate

8.8 g K2CO3 + 138.2052 g/mol = 0.0637 mol b. 0.257 g of arsenic pentachloride

0.257 g AsCl5 + 252.1866 g/mol = 1.02 x 10-3 mol c. 12.5 L of carbon dioxide at STP

12.5 L CO2 + 22.4 L/mol = 0.558 mol d. 5.00 x 1013 atoms of iron

5.00 x 1013 atoms of Fe + 6.022 x 1023 = 0.830 mol e. 8.63 x 1028 molecules of water

8.63 x 1028 molecules of H2O + 6.022 x 1023 = 1.43 x 105 mol f. 450.0 mL helium at STP

0.4500 L He + 22.4 L/mol = 0.0201 mol g. 236.0 g of ammonium phosphate

236.0 g (NH4)3PO4 + 149.086 74 g/mol = 1.58 mol h. 15.0 g of butanoic acid

15.0 g CH3CH2CH2COOH + 88.106 32 g/mol = 0.170 mol

6. Calculate the mass of each of the following. a. 2.60 mol of sodium carbonate

(2.60 mol)(105.988 74 g/mol) = 276 g

b. five million atom of gold

(5 000 000 atoms) + (6.022 x 1023) = 8.30 x 10-18 mol

(8.30 x 10-18 mol)(196.9665 g/mol) = 1.64 x 10-15 g

c. 25.0 mL of carbon dioxide at STP

(0.0250 L) + (22.4 L/mol) = 0.001 12 mol

(0.001 12 mol)(44.0098 g/mol) = 0.0491 g

d. 4.50 x 1021 molecules of decanoic acid

CH3CH2CH2CH2CH2CH2CH2CH2CH2COOH or C10H202

( )( ) gmol

mol

molg 29.12676.17200747.0

00747.010022.61050.4

23

21

=

=××


7. Calculate the molarity of 825 mL of solution, which contains 30.0 g of acetic acid.

L

mol

molg L

gg

COOHCHg606.0

825.04996.0

4996.005256.60

0.30 3 ==

8. What volume of solution can be made from 80.0 g of sodium hydroxide if a 2.00 M solution is required?

L00.100.2

mol00.2Cn

V

mol00.211887.38

NaOHg0.80

Lmol

molg

===

=

9. What mass of calcium chloride is required to produce 750.0 mL of a 0.500 M solution?

( )( )

( )( ) g6.41984.110mol375.0

mol375.0L7500.0500.0VCn

molg

Lmol

=

==×=

10. What volume 0f 14.0 M nitric acid is required to produce 750.0 mL of a 0.250 M solution?

( )( )mL4.13orL0134.0

0.14

L7500.0250.0

CVC

VVCVC

Lmol

Lmol

1

2212211 ====

11. What concentration results when 250.0 mL of a 0.125 M solution of hydrochloric acid is mixed with 125.0

mL of a 1.00 M solution?

( )( ) ( )( )

Lmol

Lmol

Lmol

ba

bbaaab

417.0L375.0

mol125.0mol03125.0

L1250.0L2500.0

L1250.000.1L2500.0125.0

VVVCVC

C

=+

=

+

+=

++

=


12. What volume of 0.225 M sulfuric acid must be mixed with 500.0 mL of a 0.750 M solution in order to obtain a 0.500 M solution?

( ) ( )( )

( )( ) ( )

( )

( )

L455.0275.0

mol125.0V

mol125.0V275.0

mol375.0V225.0mol250.0V500.0

mol375.0V225.0L5000.0V500.0

L5000.0V

L5000.0750.0V225.0500.0

VVVCVC

C

Lmola

aLmol

aLmol

aLmol

aLmol

aLmol

a

Lmol

aLmol

Lmol

ba

bbaaab

==

=

+=+

+=+

+

+=

++

=

E. STOICHIOMETRY - Begin each problem by writing a balanced chemical equation. 1. Carbon dioxide is produced in the reaction between calcium carbonate and hydrochloric acid. How many

grams of calcium carbonate would be needed to react completely with 15.0 g of hydrochloric acid? How many grams of calcium chloride would be formed?

CaCO3 + 2HCl H2O + CO2 + CaCl2

( )

( ) 2molg2

molg

3molg3

molg

CaClg8.22986.110HClmol2

CaClmol1

94460.36

HClg0.15

CaCOg6.200892.100HClmol2

CaCOmol1

46094.36

HClg0.15

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛


2. Sulfur dioxide may be catalytically oxidized to sulfur trioxide. How many grams of sulfur dioxide could be converted by this process if 100.0 g of oxygen are available for the oxidation?

2SO2 + O2 2SO3

( ) 2molg

2

2

molg

2SOg4.4000648.64

Omol1

SOmol2

9988.31

Og0.100=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

3. Phosphoric acid is produced in the reaction between calcium phosphate and sulfuric acid. How much of the phosphoric acid would be produced from 55.0 g of the calcium phosphate?

Ca3(PO4)2 + 3H2SO4 3CaSO4 + 2H3PO4

( ) 43molg

243

43

molg

243POHg8.34995.97

)PO(Camol1

POHmol2

183.310

)PO(Cag0.55=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

4. How much magnesium sulfate is needed to completely react with 145 g of sodium chloride? How much

sodium sulfate could be produced by this reaction?

MgSO4 + 2NaCl Na2SO4 + MgCl2

( )

( ) 42molg42

molg

4molg4

molg

SONag176043.142NaClmol2

SONamol1

443.58

NaClg145

MgSOg1493686.120NaClmol2

MgSOmol1

443.58

NaClg145

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

5. Gold will dissolve in the acid mixture known as aqua regia according to the following reaction:

Au + HNO3 + 3 HCl AuCl3 + NO + 2 H2O How much gold(III)chloride will be produced in this reaction when one starts with 5.0 mg of gold?

( ) 3

37.73255.303

1

1

9665.196

0.5AuClmg

Aumol

AuClmolAumgmmol

mg

mmolmg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

How much hydrochloric acid must be added initially to dissolve the all this gold?

( ) HClmgAumol

HClmolAumgmmol

mg

mmolmg

8.2461.361

3

9665.196

0.5=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛


6. How many grams of sulfuric acid will react with 400.0 g of aluminum metal?

( ) 42

422181079.98

2

3

98154.26

0.400SOHg

Almol

SOHmolAlgmol

g

molg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

7. It is desired to prepare 50.0 g of water by synthesis. How many litres of hydrogen must be used?

( ) 2

2

222.624.22

2

2

01528.18

0.50HL

OHmol

HmolOHgmol

L

molg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

8. An unknown amount of potassium chlorate was heated until no more oxygen was evolved. 15.824 g of potassium chloride remained in the test tube.

2KClO3 2KCl + 3O2

What mass of potassium chlorate had originally been placed in the tube?

( ) 3

30.265495.122

2

2

5513.74

824.15KClOg

KClmol

KClOmolKClgmol

g

molg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

What volume of oxygen gas was evolved in the process?

( ) 2

213.74.22

2

3

5513.74

824.15OL

KClmol

OmolKClgmol

L

molg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

9. How many grams of carbon can be completely burned in 15.0 L of oxygen?

( ) CgOmol

CmolOLmol

g

molL

04.8011.121

1

4.22

0.15

2

2=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

How many litres of carbon dioxide gas are produced in this reaction?

( ) 2

2

220.154.22

1

1

4.22

0.15COL

Omol

COmolOLmol

L

molL

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛


10. How many grams of magnesium metal are required to liberate 250.0 mL of hydrogen gas from hydrochloric acid?

( ) MggHmol

MgmolHLmol

g

molL

271.0305.241

1

4.22

2500.0

2

2=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

Exactly how much acid would be used up in this reaction?

( ) HClgHmol

HCmolHLmol

g

molL

8141.046094.361

lg2

4.22

2500.0

2

2=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

11. Will 30.0 L of fluorine gas completely react with 50.0 L of hydrogen gas? Which gas is in excess and by how much? What volume of hydrogen fluoride is formed in this reaction?

( )

( )

( )

excess in L 20.0 needed L 30.0 - available hydrogen of L 50.0

0.304.221

1

4.22

0.30

excess. in is hydrogen and reactant limiting the is Fluorineproduced. is fluoride hydrogen of L 0.60

.1004.221

2

4.22

0.50

0.604.221

2

4.22

0.30

2

22

2

2

2

2

=

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

neededHFLOmol

HmolFL

HFLHmol

HFmolHL

HFLFmol

HFmolFL

molL

molL

molL

molL

molL

molL


12. A mixture containing 100.0 g of H2 and 100.0 g of O2 is sparked so that water is formed. How much water is formed?

( )

( )

formed. are water of g 113excess. in is hydrogen and reactant limiting the is Oxygen

11301528.181

2

9988.31

0.100

89401528.182

2

01588.2

0.100

2

2

22

2

2

22

OHgomol

OHmolOg

OHgHmol

OHmolHg

molg

molg

molg

molg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

13. When copper is heated with sulfur, Cu2S is formed. How many grams of Cu2S could be produced if 100.0 g of copper is heated with 50.0 g of sulfur?

( )

( )

formed. is sulfide copper(I) of g 125.2excess. in is sulfur and reactant limiting the is Copper

248158.1591

1

066.32

0.50

2.125158.1592

1

546.63

0.100

2

2

2

2

SCugSmol

SCumolSg

SCugCumol

SCumolCug

molg

molg

molg

molg

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

=⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

14. What volume (at STP) of carbon monoxide is required to produce 100.0 g of iron according to the equation: Fe2O3 + CO Fe + CO2 (not balanced)

( ) COLFmol

COmolFegmol

L

molg

2.604.2222

3

847.55

0.100=

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛

⎟⎟⎟

⎠

⎞

⎜⎜⎜

⎝

⎛


*15. Hydrozine is a nitrogen-hydrogen compound hoving the formulo NzH+. rt is on oily, colourless liguid thot

freezesot 1.5oC andloils ot 113.5"c. ihe principol use of hydrozine ond certoin compounds derived from it

is o rocket fuels, but it is olso used in fuei cells, in the treotment of woter in boilers to removed dissolved

oxygengos, ond in the plostics industry. One widely used method for the monufoctureol hydrozine is the

Roschig process. The Roschig process involves three reoction steps. rn the first sfep, sodium hydroxide is

reocted with chlorine to proJuce sodium hypochlorite, sodium chloride ond woter. The sodium hypochlorite

produced in the first step is reocted with ommonio in the second step to produce chloromine (NHzcl) ond

sodium hydroxide. The chloromine ond sodium hydroxide produced in the second step is reocted with

ommonio in the third step to produce hydrazine, sodium chloride and woter. A chemicol plone using the

Roschig process obtoins O.2gg k9 of g8.O% hydrozine for every 1.00 k9 of chlorine. Whot ore the

theoreticol, octuol, ond percent yields of pure hydrozine?

2NaOH * Clz ) NaOCI + NaCl * HzON a O C I * N H g t N H 2 C | + N a O H

NH2CI + NaOH * NHg ) NzHr + NaCl * HzO

ctnol

- ORoANIC CHEAAI5TRY

Drow the structure of eoch of the following compounds:

o) 2,3-dinitrophenolot+

,\ro.LO^ do,v

b) 2,3,4-'trimethyloctone

ct{. -cl+ -cil+ ^ Ct\-CAz-Uza\r- C(ltCHz ctl. ctlt

c) 1,3-cyclohexodiene

d) 4,4-dimethyl-2-Pentene

QiltCIlrcH=cH - q

crL

lfr -"r ruaocr]f

Jt 1-", % JtL mot NH,ct)lt mor n,n.l6r.oorr,

;r)t mot naoct )lL mot NH2ct )

actual yield = (98.Oolo[zSS g) = zel g = dctual yield

percentage Yield =actual yield

x 1OO = 64.80/o yieldtheoretical yield

F1.

1OOO g Ct,

70.906

= 452 g NrHo = theoretical Yield

Chemistry 11 - Finol Exomination Review - Answers - of ?7

cilt

e) phenylethene

Cl+ = CFl.

f) 2-phenylProPone

l) proponone

ot l

cSg- c - ct{3

m) l-proPoxybutone

C |{tcH ,-CA;o -Ct{>- CilL-CLz'C[!3

9) 1,3-dibromonoPhtholene6v

@."n) ethyl chloroethonoote

o,, erlr_ Ctl,l f t re -D-ct

p) 3-methYl-3-PentonolorII

CA;cHv- l- ' 'Hr-CIlscHr

o) iodobenzene

L

,\| / ) t\Z

-methylphenol

ot^.

Yctf g

q ) 4

.,*rd -c*=

h) 1,2-ethonediol

fil'- ts'otl o l+

i) proponol

ol l

cA r-cr|;cr+

i) pentonoic ocid

Ot,

CLl; ctl;cl;il\;C-o t't

k) propyl ethonoote r) 2-omino-3-phenylproponol o

I c*-- cu - ttCllr-C-c: -CHr{A>- CL A

" IU,\?

Chemistry 11 - Finol Exominotion Review - Answers - Poge 24 of ?7

s) 3-methyl-3-pentonol

cr,c l{t cHr- L -c{r- ctl,

r e l

o|+

t) ethylmethyl ether

CHr'tfit o - c{t

u) butyl butanoote

w) nitromethone

iJ 0 z- ct\3

x) 3-ethyl-5-methyl-2-naPhthol

0r l

CAr-1t ; C$ ; Go - cttg ctl r-cl{ ; t! t

o-t-ou

v) 3-ethylhexonoicocid

CtJ;cU;t[r-ctt-crt]cAzICtlr

Z. Write o bolonced eguotion for each of the reoctions below. Identify the type of reoction ond nome the

products.

o) pentone + bromine )

Ctls- CvI; CII;Cil;Ctl. t

g^,t"r1. S"*t a" nx n

b) 2,3-dichloro-?-penlene + woter )

ci lo-9= t -cl lz-Lt l t + t l"-c) ->

cr c l

ad/i lsn- (xn

c) butcnoic ocid + ethonol )o

Ctle-tl;Crl; ' i-oU' ct\-CH>-oH

(sler ili ealrtm rrYr'l {Vlrf arlw,oo'fu t

t'f oAt l

cHr- ?- i- ctt>- cHtC t c l

?,?- d:Nwo 4 "*+"Jo

+ eLr4lll-cil'-t -o -cLt'c4t + HrO

palor

w:L*.cfit

y) 1,3,5-cyclohexonetriol0uIn

L}-

noAlotl

z) propyl 2,3-dimethylPentonoote

f t

Ctlr-c+,- f - fA

-t-o- cL-ut>-ctt3CH, CH'

&,k - -> " * r ' cL r -c l r - c {> -cd ' I H6 .

t- VNno f-r#*-,, + hgd"+ogt'^ brorv'tolt

Chemistry 11 - Finol Exominotion Review - Answers - Poge 25 of 27

d) benzene + iodine )

V t T - - )

*wbg\L(r^ Yxva

e) *t[f';ti oxYsen ), A

/,',Y n-r + lL O LS/-*/

AOvy'hv{firn rKn

f) 2-nophthol + oxygan )( ( - ' r H c o )

z AryoP -, -zz o z- >

46P,l,.sha*- | x r'r

9) 2,5,6-trimethyl-3-heptyne + fluorine )

c1t,-c+-c= c- 7H - Cl ' l -cHt r'

br+t it+t c{t

aold; lsaw,rF*

llro

r daleY

F>-> ,r,- t;,- 7= i

@t + i l - !

I- iodo fun# + hgd-'ro6t'* ;o/i'J"

to QO. + 4 H"o

w bon d,irr; oh r aa/t'r'

ZD eO> + &

eo.rbun dr'orioL

-c t l -cH -c i ls

brt, b't'

3,4 -d:flrrwp -716,L'hl,*d't^"( - S- huflov*

h) 3,4-diethyl-2-methylhexonoic ocid + 1-Penfonol ) o

t$$$$$;tr ::; - eH - E-o u r [X;ru;cftut{t1-+ c t',,-i;'ti$; :'-'r:or:':*e bil" iU

btl. tl+- clt.

Lu, b*, ?nV s,.t -dicl4/-v,*iii n +*oJ" r woln

i) t-^"d6i{,0*Ztfn" * oxvsen )

cp.-ctlr- LH-ctlL- ct! r I 0, t b cDL 't 1 H-bo

, Ir*. ot'v b"-' d'iry;J' r "laletCdn-bu5'1".;' lrx'^

j) a-pentyne + (2 mol) woter )

| Oil

cH3-c=c-cHL-C{1" t z t+>o ) C}+r - f^ f -CHz-cH,n o f l

/)"/d; h "t' ' fxn g.s -

f/.Jo^'r,il:6(

Chemistry 11 - Finol Exominotion Review - Answers - Poge 26 of 27

3. Name the following compounds.

a) 3-methyl-5-nitro-3-hexene

b) 1,3-dichlorocyclohexane

c) 1,3-diamino-5-nitrobenzene

d) 3,6-dibromo-2-naphthol

e) diethyl ether or ethoxyethane

f) 1,4-cyclohexadiene

g) 2,3-dimethylbutanoic acid

h) 2,3,4-pentanetriol

i) 2-methyl-3-pentanone

j) 3-ethyl-2-methylpentanal

k) methyl propanoate

l) 4,5-dimethyl-7-propyl-1-naphthol

m) 2,4-hexadiyne

n) 1,2-butadiene

o) pentyl 3,3-dimethylbutanoate

p) 1,3,5-cyclohexanetriol

q) 1-amino-2-bromo-4-nitrocyclopentane

r) 2,3-dimethylpentanedioic acid

s) butyl 2-methylpropanoate

t) phenyl-3,4-diaminopentanoate


chemistry 11 final examination review -...

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