Free Printable A level chapter 19-24 review

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Worksheets

A level chapter 19-24 review

Total questions: 159

Worksheet time: 2hrs 24mins

Name

Class

Date

1.

Which is an example of an ionic crystal?

a)

Quartz

b)

NaCl

c)

Ammonia

d)

Iron

2.

This type of solid forms a regular repeating

three-dimensional structure called a crystal lattice. What type of solid is this?

a)

amorphous

b)

crystalline

c)

glass

d)

diamond

3.

What is Lattice Energy?

a)

one mole of an ionic compound is formed from its gaseous ions under standard conditions

b)

one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1- ions under standard conditions

c)

one mole of an ionic compound is formed from its atoms under standard conditions

d)

one mole of gaseous atoms is formed from its element under standard conditions

4.

Bond breaking is

a)

exothermic

b)

endothermic

5.

bond making is

a)

exothermic

b)

endothermic

6.

Factors affecting the value of lattice energy

a)

the size of the ions only

b)

the charges of the ions only

c)

BOTH the size of the ions AND the charges of the ions

7.

When comparing different ionic compounds, the one with the greater charges will have a _______ magnitude of lattice energy.

a)

greater

b)

less

c)

same

8.

When comparing different ionic compounds, the one with the ions of greater size or radius will have a _______ magnitude of lattice energy.

a)

greater

b)

smaller

c)

same

9.

Choose the ionic compound with the greater magnitude of lattice energy.

a)

NaCl

b)

KCl

10.

Choose the ionic compound with the greater magnitude of lattice energy.

a)

CaCl₂

b)

KCl

11.

Choose the ionic compound with the greater magnitude of lattice energy.

a)

CaCl₂

b)

CaS

12.

Choose the ionic compound with the greater magnitude of lattice energy.

a)

NaCl

b)

NaBr

13.

Choose the ionic compound with the greater magnitude of lattice energy.

a)

NaCl

b)

CsI

14.

The equation below represents .........

Mg²⁺ (g) + 2Cl^-(g) --> MgCl (g)

a)

lattice energy

b)

enthalpy of formation

c)

enthalpy of atomisation

d)

electron affinity

15.

Lattice energy becomes more exothermic when .......

a)

size of ions are small

b)

charge of ions are small

16.

Lattice energy is always ...........

a)

endothermic

b)

exothermic

c)

can be both endothermic or exothermic

17.

Enthalpy of atomisation is always ...........

a)

endothermic

b)

exothermic

c)

can be both endothermic or exothermic

18.

Ionisation energy is always ...........

a)

endothermic

b)

exothermic

c)

can be both endothermic or exothermic

19.

Second and third electron affinities are always exothermic.

a)

true

b)

false

20.

What is the general trend of first electron affinity going down the group 16 (ignore the first period)?

a)

more negative

b)

less negative

c)

no change

21.

enthalpy of formation is always exothermic

a)

true

b)

false

22.

What is the first electron affinity?

a)

one mole of electrons is added to one mole of gaseous 1- ions to form one mole of gaseous 2- ions under standard conditions

b)

one mole of gaseous atoms is formed from its element under standard conditions

c)

one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1- ions under standard conditions

23.

If we want to calculate lattice energy we use

a)

Hess's law

b)

Born-Haber cycle

c)

Enthalpy change of formation

d)

Electron Affinity

24.

Factors affecting anion polarisation

a)

Charge density of the cation

b)

The polarisability of the anion

25.

The solubility of the group 2 hydroxides ........... down the group

a)

increases

b)

decreases

26.

The solubility of the group 2 sulfates ...... down the group

a)

decreases

b)

increases

27.

Hess’s law can be used to calculate (a) for reactions that do not occur directly or cannot be found by experiment

28.

When an ionic solid dissolves in water, bonds are formed between water molecules and the ions. These bonds are called ion–dipole bonds

a)

dipole –dipole bonds

b)

ion–ion bonds

c)

ion–dipole bonds

29.

According to Hess's law, we can calculate lattice energy with the formula

a)

ΔHlatt = ΔHf – ΔH1

b)

ΔHf = ΔHlatt – ΔH1

c)

ΔHlatt = ΔH1 - ΔHf

30.

Rank the compounds below from highest to lowest lattice energy: NaBr, AlN, BeO, Li₂S

a)

AlN > BeO > Li₂S > NaBr

b)

NaBr > Li₂S > BeO > AlN

c)

BeO > Li₂S > AlN > NaBr

d)

AlN > Li₂S > BeO > NaBr

31.

Using the bank of lettered choices below, choose the correct sequence of choices necessary to complete this descriptions below:

A. Lattice Energy

B. Ionization Energy

C. Binding Energy

_______________ is the energy required to remove any particle from a system of particles, in this case, an electron from any atomic orbital, while _______________ is the energy required to remove the most loosely bound electron from an atom.

On a separate note, _______________ is the energy required to separate an ionic compound into its constituent ions.

a)

A, B, C

b)

C, B, A

c)

B, C, A

d)

C, A, B

32.

Given the following ionic compounds, place them in order of decreasing lattice energy:

Al₂S₃,BaI₂, KBr, NaF, MgO

a)

Al₂S₃, BaI₂, MgO, KBr, NaF

b)

KBr, NaF, BaI₂, MgO, Al₂S₃

c)

Al₂S₃, MgO, BaI₂, NaF, KBr

d)

MgO, NaF, KBr, BaI₂, Al₂S₃

33.

Given the list of ions below without their radii, use what you know about trends in atomic size to rank the ions in order of increasing strength of interactions (i.e. - most exothermic):

N^3-, O^2-, F^-, Na⁺, Mg²⁺, Al³⁺

a)

Na⁺ < F^- < Mg²⁺ < O^2- < N^3- < Al³⁺

b)

F^-< O^2- < N^3- < Na⁺ < Mg²⁺< Al³⁺

c)

Al³⁺< N^3-< O^2-< Mg²⁺< F^- < Na⁺

d)

F^-< Na⁺< O^2- < Mg²⁺ < N^3- < Al³⁺

34.

Which of the equation below refers to the standard enthalpy of formation, ΔH_f^o?

a)

Na(g) ---> Na⁺(g) + e^-ΔH= -364 kJmol^-1

b)

C₂H₅OH (l) + 3O₂(g) ---> 2 CO₂(g) + 3H₂O (l)ΔH= - 1286 kJmol^-1

c)

2C (s) + 2H₂(g) ---> C₂H₄ (g)ΔH= - 52.3 kJmol^-1

d)

Na⁺(g) ---> Na⁺(aq) ΔH= - 364 kJmol^-1

35.

Heat released when 1 mole of aqueous ions is formed from gaseous ions.

a)

Enthalpy of Formation

b)

Enthalpy of Combustion

c)

Enthalpy of Hydration

d)

Enthalpy of Lattice Dissociation

36.

Which equation represents the process that occurs when the standard enthalpy of atomisation of iodine is measured?

a)

½ I₂(s) → I (g)

b)

I₂(s) → 2I (g)

c)

½ I₂(g) → I (g)

d)

I₂(g) → 2I (g)

37.

Lowest lattice enthalpy (∆_LEH)?

a)

NaBr

b)

LiCl

c)

KI

38.

First electron affinity of an element is usually ________ and second electron affinity is ____________.

a)

exothermic, endothermic.

b)

exothermic, exothermic.

c)

endothermic, endothermic.

d)

endothermic, exothermic.

39.

Given the Born-Haber cycle below, which of the following enthalpy change is not represented by steps a-g

a)

Enthalpy of formation

b)

Enthalpy of atomization of calcium

c)

1^st electron affinity of bromine

d)

2^nd electron affinity of bromine

40.

Identify the enthalpy change represented by letter g.

a)

Enthalpy of formation

b)

Enthalpy lattice of formation

c)

Electron affinity of bromine

d)

Enthalpy of atomisation of zinc

41.

What enthalpy change does the number 4 represent?

a)

Formation

b)

1^st ionisation energy of Na

c)

1^st electron affinity of Na

d)

Atomisation of Na

42.

Enthalpy change is represented by step y ?

a)

Enthalpy of formation calcium bromide

b)

Enthalpy of atomisation of calcium

c)

1^st electron affinity of bromine

d)

Enthalpy of atomisation of bromine

43.

What is value of the enthalpy of lattice energy for NaCl (s)?

a)

+766

b)

+230

c)

-230

d)

-766

44.

Calculate value x.

a)

-1970 kJ

b)

-266 kJ

c)

+266 kJ

d)

+1970 kJ

45.

What is the value of enthalpy of formation of NaF (s)?

a)

-928 kJ mol^-1

b)

-574 kJ mol^-1

c)

+107 + 496

d)

+107 + 496 - 928

46.

What is the definition of lattice enthalpy?

a)

Heat released when one mole of solid ionic compound is formed from its gaseous ions.

b)

Heat released when an ionic lattice is formed from one mole of its gaseous ions.

c)

Heat released when one mole of an ionic lattice is formed from ions in their standard states.

d)

Heat released when an ionic lattice is formed from one mole of gaseous atoms.

47.

Which of the equation below refers to the standard enthalpy of formation, ΔH_f^o?

a)

Na(g) ---> Na⁺(g) + e^-ΔH= -364 kJmol^-1

b)

C₂H₅OH (l) + 3O₂(g) ---> 2 CO₂(g) + 3H₂O (l)ΔH= - 1286 kJmol^-1

c)

2C (s) + 2H₂(g) ---> C₂H₄ (g)ΔH= - 52.3 kJmol^-1

d)

Na⁺(g) ---> Na⁺(aq) ΔH= - 364 kJmol^-1

48.

The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions. This is the enthalpy of.....

a)

Combustion

b)

Formation

c)

Lattice breaking

d)

Ionisation

49.

Heat released when 1 mole of aqueous ions is formed from gaseous ions.

a)

Enthalpy of Formation

b)

Enthalpy of Combustion

c)

Enthalpy of Hydration

d)

Enthalpy of Lattice Dissociation

50.

Which equation represents the process that occurs when the standard enthalpy of atomisation of iodine is measured?

a)

½ I₂(s) → I (g)

b)

I₂(s) → 2I (g)

c)

½ I₂(g) → I (g)

d)

I₂(g) → 2I (g)

51.

Lowest lattice enthalpy (∆_LEH)?

a)

NaBr

b)

LiCl

c)

KI

52.

First electron affinity of an element is usually ________ and second electron affinity is ____________.

a)

exothermic, endothermic.

b)

exothermic, exothermic.

c)

endothermic, endothermic.

d)

endothermic, exothermic.

53.

Which one of the following statements is not correct?

a)

The first ionisation energy of iron is greater than its second ionisation energy.

b)

The magnitude of the lattice enthalpy of magnesium oxide is greater than that of barium oxide.

c)

The oxidation state of iron in [Fe(CN)₆]³⁻ is greater than the oxidation state of copper in [CuCl₂]⁻

d)

The boiling point of C₃H₈ is lower than that of CH₃CH₂OH

54.

Given the Born-Haber cycle below, which of the following enthalpy change is not represented by steps a-g

a)

Enthalpy of formation

b)

Enthalpy of atomization of calcium

c)

1^st electron affinity of bromine

d)

2^nd electron affinity of bromine

55.

Identify the enthalpy change represented by letter g.

a)

Enthalpy of formation

b)

Enthalpy lattice of formation

c)

Electron affinity of bromine

d)

Enthalpy of atomisation of zinc

56.

What enthalpy change does the number 4 represent?

a)

Formation

b)

1^st ionisation energy of Na

c)

1^st electron affinity of Na

d)

Atomisation of Na

57.

Enthalpy change is represented by step y ?

a)

Enthalpy of formation calcium bromide

b)

Enthalpy of atomisation of calcium

c)

1^st electron affinity of bromine

d)

Enthalpy of atomisation of bromine

58.

What is value of the enthalpy of lattice energy for NaCl (s)?

a)

+766

b)

+230

c)

-230

d)

-766

59.

Calculate value x.

a)

-1970 kJ

b)

-266 kJ

c)

+266 kJ

d)

+1970 kJ

60.

What is the value of enthalpy of formation of NaF (s)?

a)

-928 kJ mol^-1

b)

-574 kJ mol^-1

c)

+107 + 496

d)

+107 + 496 - 928

61.

What two factors govern whether a collision between reacting particles will be effective?

a)

orientation and potential energy

b)

kinetic energy and temperature

c)

kinetic energy and orientation

d)

potential energy and kinetic energy

62.

Under the collision theory, the particles must collide with ____ and ____ for a reaction to occur.

a)

sufficient rate and sufficient energy

b)

sufficient surface area and correct orientation

c)

sufficient catalyst and sufficient energy

d)

sufficent energy and correct orientation

63.

More collisions correspond to a:

a)

Faster reaction rate

b)

Slower reaction rate

c)

Constant reaction rate

d)

None of the above

64.

Why does a higher temperature increase the rate of a reaction?

a)

it increases both the frequency and energy of particle collisions

b)

it only increases the frequency of particle collisions

c)

it only increases the energy of particle collisions

d)

it reduces the activation energy of the reaction

65.

In the chemical reaction, the successful formation of products depends on the following factors EXCEPT:

a)

The collisions between reacting molecules must be effective

b)

The colliding molecules must approach each other in the right orientation

c)

The energy of colliding molecules must be equal to or more than activation energy

d)

A collision between molecules must produce sufficient energy to overcome the activation energy

66.

The catalyst alters the rate of a chemical reaction by,

a)

providing an alternative pathway with a lower Ea

b)

changing the products formed in the direction of the reaction

c)

providing a surface on which the molecules react

d)

increasing the frequencies of collisions between molecules

67.

The ____________ states that atoms, ions, or molecules must collide in order to react.

a)

transition state

b)

activation energy

c)

rate law

d)

collision theory

68.

A ______________ is a substance that increases the rate of a reaction without being used up during the reaction.

a)

catalyst

b)

product

c)

reactant

d)

solute

69.

Increasing the temperature will increase the kinetic energy of particles, therefore increasing the collisions between particles.

a)

false

b)

true

70.

Liz observes 2 reactions, A and B.

Reaction A is performed in a beaker placed on top of a hotplate.

Reaction B is is performed in a beaker placed within a container of ice.

Which of the following statements is true?

a)

Reaction A will react slower because the higher temperature will cause more collisions because the heat makes the molecules tired.

b)

Reaction A will react slower because the higher temperature will cause give the reactants higher speed and kinetic energy.

c)

Reaction B will react faster because the lower temperature causes the particles to move slower, allowing them to have more time to attach to one another.

d)

Reaction B will react slower because the lower temperature makes the particles move slower, resulting in less collisions.

71.

Which of the following best describes why increasing temperature increase reaction rate?

a)

Activation energy is reduced.

b)

Collisions become more frequent.

c)

Collisions become more energetic.

d)

Collisions become more frequent and more energetic.

72.

To measure rate of reaction, we can monitor the concentration change in which of the following?

a)

Reactants only

b)

Products only

c)

Reactants or products

73.

What is the activation energy?

a)

The energy of the products

b)

The energy of the reactants

c)

The minimum energy needed by the reactants to form the products.

d)

The energy lost in the reaction

74.

What is the best definition of rate of reaction?

a)

The time it takes to use up all the reactants

b)

The rate at which all the reactants are used up

c)

The change in concentration of a reactant/product per unit time

d)

The time it takes for one of the reactants to be used up

75.

The question below is about the Maxwell–Boltzmann distribution shown for a sample of a gas, X, at two different temperatures.

Which statement is correct for the higher temperature?

a)

The area under the curve to the left of Ea decreases.

b)

The total area under the curve increases.

c)

The activation energy decreases.

d)

More molecules have the mean energy.

76.

This question is about the Maxwell–Boltzmann distribution of molecular energies in a sample of a gas shown in the figure below.

What does the area under the curve represent?

a)

The total energy of the particles.

b)

The total number of particles.

c)

The number of particles that can react with each other.

d)

The total number of particles that have activation energy.

77.

The graph below shows a typical energy distribution for particles of an ideal gas in a sealed container at a fixed temperature.

Which of the following statements is true?

a)

Position A represents the mean energy of a molecule in the container.

b)

Addition of a catalyst moves the position of E_A to the right.

c)

The area under the curve to the right of E_A represents the number of molecules with enough energy to react.

d)

The position of the peak of the curve at a higher temperature is further away from both axes.

78.

The graph shows the Maxwell−Boltzmann distribution of molecular energies in a sample of gas at a fixed temperature.

The horizontal axis label is

a)

energy

b)

J mol⁻¹

c)

E

d)

mean energy

79.

What is another name for enzymes?

a)

chemical catalysts

b)

microorganisms

c)

biological catalysts

d)

inorganic catalysts

80.

If the reactant particles collide with less than the activation energy, the particles will be rebound, and no reaction will occur. (BONUS)

a)

True

b)

False

81.

The diagram represents the Maxwell‒Boltzmann energy distribution curve of the reactants for a chemical reaction with different activation energies, E_a1 and E_a2.

What is the reason why the rate of the reaction with activation energy E_a2 is greater?

a)

More frequent collisions between the particles occur.

b)

More energetic collisions between the particles occur.

c)

A catalyst has been added.

d)

The temperature is higher.

82.

Which is the reactant - and why?

a)

A, as concentration decreases

b)

B, as concentration decreases

c)

A, as concentration increases

d)

B, as concentration increases

83.

When the system X + 2 Y <=> Z has reached equilibrium, which of the following is TRUE?

a)

Forward reaction and Backward reaction stop.

b)

Forward reaction has sped up and backward reaction has slowed down.

c)

Forward and backward reactions occur at the same rate.

d)

Forward reaction has slowed down and backward reaction has sped up.

84.

At what time the reaction reached equilibrium?

a)

t₁

b)

t₂

c)

t₃

d)

t₄

85.

For the reaction...
SO₂ + O₂ <=> SO₃
If the concentration of SO₂is increased, the equilibrium position of the reaction will shift ___________.

a)

to the left

b)

to the right

c)

to the left and right

d)

neither left nor right

86.

For the reaction...
SO₂ + O₂ <=> SO₃If the equilibrium position shifts to the right, the concentration of O₂ will ___________.

a)

increase

b)

decrease

c)

remain the same

d)

double

87.

For the reaction...
N_{2 (g)}+ 3 H₂ _(g) <=> 2 NH_{3 (g)}
If the pressure in the system is increased, which substance(s) will increase in concentration?

a)

N₂ (as 2 mol gas → 4 mol gas)

b)

H₂ (as 2 mol gas → 4 mol gas)

c)

N₂ and H₂(as 2 mol gas → 4 mol gas)

d)

NH₃ (as 4 mol gas → 2 mol gas)

88.

For the reaction...
N₂ + O₂ <=> 2NO: Δ H = +182 kJ mol^-1.
If the temperature is increased the equilibrium position will shift _______.

a)

to the left

b)

to the right

c)

to the left and right

d)

neither left nor right

89.

What is the K_c expression for this reaction?
2 NO_(g) + O_2(g) ⇌ 2 NO_2(g)

a)

K_c= [NO₂]² / [NO]²[O₂]

b)

K_c= [NO]²[O₂] / [NO₂]²

c)

K_c= [NO]²[O₂] [NO₂]²

d)

K_c= [NO₂]² / [NO]² + [O₂]

90.

An equilibrium constant with a large value, e.g. K_c = 1000, indicates…

a)

A very fast reaction

b)

Mostly products at equilibrium

c)

Mostly reactants at equilibrium

d)

Nothing useful at all

91.

Le Chatelier's Principle states that....
If a (dynamic) equilibrium is disturbed by changing the conditions, the position of equilibrium................

a)

moves to increase the change

b)

moves to counteract the change

c)

does not change

92.

Given: 2A(g) <=> 2B(g) + C(g). At a particular temperature, K_c = 16000.
Raising the pressure, by decreasing the volume of the container, will...

a)

cause the value of K_c to increase

b)

cause the value of K_c to decrease

c)

have no effect on the value of K_c as temperature does not change

d)

favour the forward reaction

93.

What are [A] and [B]?

a)

concentration of reactants

b)

concentration of products

c)

energy of reactants

d)

energy of products

94.

What is the equilibrium expression for:
Fe₃O₄(s) + 4H₂(g) <=> 3Fe(s) + 4H₂O(g)
K_c =

a)

[Fe]³[H₂O]⁴ / [Fe₃O₄] [H₂]⁴

b)

[Fe₃O₄] [H₂]⁴/ [Fe]³[H₂O]⁴

c)

[H₂O]⁴ / [H₂]⁴

d)

[Fe] [H₂O] / [Fe₃O₄] [H₂]

95.

Identify the incorrect statement about achieving equilibrium.

a)

achieved when product and reactant concentrations are equal

b)

achieved when forward and reverse reaction rates are same

c)

achieved when concentration of reactants is stable/constant

d)

achieved when concentration ratio of reactants to products becomes stable, thus fixing the equilibrium constant (equilibrium position)

96.

PCl_{3 (g)} + Cl_{2 (g)} ↔ PCl_{5 (g)} + energy

Which of the following causes an increase in the moles of PCl₅ present at equilibrium shown?

a)

decreasing the volume of the container

b)

raising the temperature

c)

adding a mole of He gas at constant volume

d)

adding a mole of He gas at constant pressure

97.

What kind of equilibrium does the reaction below show?

H_2(g) +I_2(g) ↔ 2HI_(g)

a)

Heterogeneous equilibrium, all the reactants and products are in the same physical state.

b)

Homogeneous equilibrium, all the reactants and products are in the same physical state.

c)

Heterogeneous equilibrium, the reactants and products are present in more than one physical state

d)

Homogeneous equilibrium, the reactants and products are present in more than one physical state

98.

Given: A + B --> 2C K = 1.2 x 10^-3

What is the K for the reaction 2C --> A + B?

a)

1.2 x 10^-3

b)

1.2 x10³

c)

833

d)

83333

99.

If red/orange region is absorbed, what is the color that we see. Refer to color wheel

a)

Red

b)

Orange

c)

Blue

d)

Yellow

100.

If violet region is absorbed, what is the color that we see. Refer to color wheel

a)

Violet

b)

Orange

c)

Blue

d)

Yellow

101.

Color of metal complexes are due to the following factors EXCEPT

a)

Types of ligands

b)

Nature of transition metal

c)

Oxidation states

d)

Type of orbitals

102.

3d orbitals are degenerate. What does it means?

a)

All 3d orbitals are in same order

b)

All 3d orbitals have different electron arrangement

c)

All 3d orbitals are in the same energy level

d)

All 3d orbitals have the same shape

103.

Ti³⁺ solution in the presence of water, is violet in color. Why?

a)

Presence of water ligands cause color formation

b)

Electron can be excited

c)

Wavelength of light can be absorbed

d)

Presence of water ligand cause the splitting the 3d orbitals

104.

Sc³⁺ solution in the presence of water is colorless. Why?

a)

Presence of ligands cause the splitting of 3d orbital.

b)

There is no splitting of 3d orbitals

c)

There is no electrons to be excited

d)

Energy level is too high for electron to reach

105.

Zn²⁺ solution in the presence of water is colorless. Why?

a)

Presence of ligands cause the splitting of 3d orbital.

b)

There is no splitting of 3d orbitals

c)

There is no electrons to be excited

d)

3d orbitals are fully filled and no electrons can be excited

106.

Cu²⁺ solution in the presence of water is blue. Why?

a)

Splitting of 3d orbital and electron can be easily excited

b)

Blue light is aborbed to excite elec to higher energy level

c)

Blue light is emitted when elec move to lower energy level

d)

3d orbital are partially filled and electrons can be excited

107.

The following statements are true about spectrochemical series EXCEPT

a)

CN^- ion is the strongest ligand

b)

Strong ligand displaces weak ligand

c)

Strong ligand causes greater split in 3d orbital

d)

Weak ligand will not split the 3d orbital

108.

Strong ligand causes a greater split of 3d orbital because

a)

greater repulsion between ligands and 3d electrons

b)

strong ligand has greater energy

c)

strong ligand causes a greater displacement

d)

strong ligand has greater electron density

109.

V⁵⁺ causes a greater split than Fe³⁺. Why?

a)

Stronger ESF between nucleus and ligands

b)

Weaker ESF between nucleus and ligands

c)

Stronger ESF between nucleus and ligand cause greater repulsion

d)

Weaker ESF bet nucleus and ligand cause weaker repulsion

110.

Which of the following have the greatest ΔE splitting of 3d orbitals?

a)

A

b)

B

c)

C

111.

Which of these ions is violet in colour?

a)

V²⁺

b)

V³⁺

c)

VO²⁺

d)

VO₂⁺

112.

An element where the last electron added was into a d subshell is the definition of...

a)

Transition Metal

b)

d-block element

c)

catalytic ion

d)

complex ion

113.

What is a ligand

a)

An atom, molecule or ion that forms a dative covalent bond to a central metal ion in a complex.

b)

An atom, molecule or ion that accepts a pair of electron from a central metal ion in a complex.

c)

A central metal ion with one or more covalent bonds to it

d)

A d-block element that has a partially filled d subshell

114.

Tetraaquadihydroxocopper (II) is

a)

dark blue

b)

a pale blue ppt

c)

a blue solution

d)

yellow (but looks green)

115.

Which Hydroxide Precipitate is "Amphoteric"

a)

[Cu(H₂O)₄(OH)₂]

b)

[Fe(H₂O)₄(OH)₂]

c)

[Cr(H₂O)₃(OH)₃]

d)

[Fe(H₂O)₃(OH)₃]

116.

VO₂⁺ --> VO²⁺ --> V³⁺ --> V²⁺

a)

Yellow --> (Green) --> Blue ---> Violet --> Green

b)

Yellow --> Orange --> Green ---> Blue

c)

Yellow --> (Green) --> Blue ---> Green --> Violet

d)

Violet --> Green --> Blue ---> (green)--> Yellow

117.

The coordination number of [Cu(en)₃]²⁺ is

a)

3

b)

6

c)

2

d)

12

118.

The number of en ligands dative covalently bonded to Cu in [Cu(en)₃]²⁺ is

a)

3

b)

6

c)

2

119.

[Cu(H₂O)₆]²⁺ + edta^4- <−−>[Cu(edta)]^2- + 6H₂O

a)

The entropy decreases from left to right

b)

The entropy increases from right to left

c)

The entropy increases from left to right

120.

Which is not an example of a ligand substitution reaction?

a)

[Cu(H₂O)₆]²⁺ + 2OH^- --> [Cu(H₂O)₄(OH)₂] + 2H₂O

b)

[Cu(H₂O)₆]²⁺ + 4Cl^- --> [Cu(Cl)₄]^2- + 6H₂O

c)

[Cu(H₂O)₄(OH)₂] + 4NH₃--> [Cu(H₂O)₂(NH₃)₄]²⁺ + 2H₂O + 2OH^-

d)

[Cu(H₂O)₆]²⁺ + 3en --> [Cu(en)₃]²⁺ + 6H₂O

121.

What is entropy?

a)

The measure of randomness or disorder in a system

b)

The measure of energy dispersal in a system

c)

The measure of temperature change in a system

d)

The measure of pressure change in a system

122.

Which of the following is a spontaneous change? (Select all that apply)

a)

Sugar dissolving in water

b)

Water turning to ice at 10°C

c)

Ethanol vaporizing at 20°C

d)

Limestone decomposing at room temperature

123.

What happens to the entropy when a solid melts?

a)

Entropy increases

b)

Entropy decreases

c)

Entropy remains the same

d)

Entropy cannot be determined

124.

Which substance has the highest entropy?

a)

Diamond

b)

Graphite

c)

Calcium

d)

Lead

125.

What is the standard molar entropy of water?

a)

48.0 J/K mol

b)

69.9 J/K mol

c)

188.7 J/K mol

d)

213.6 J/K mol

126.

Which type of substances have low entropy?

a)

Gases

b)

Solids

c)

Liquids

d)

Aqueous solutions

127.

In the reaction 2N₂O₅(g) → 4NO₂(g) + O₂(g), what is the entropy change of the system?

a)

Positive

b)

Negative

c)

Zero

d)

Cannot be determined

128.

What is the entropy change of the system for the reaction 2Ca(s) + O₂(g) → 2CaO(s)?

a)

Increases

b)

Decreases

c)

Remains Same

d)

Information missing for correct evaluation

129.

Which of the following reactions has a positive total entropy change?

a)

S(s) + H₂(g) → H₂S(g)

b)

2C(graphite) + N₂(g) → C₂N₂(g)

c)

4P(s) + 5O₂(g) → P₄O₁₀(s)

d)

C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(l)

130.

How does increasing the temperature affect the ability of zinc carbonate to undergo thermal decomposition?

a)

Increases the entropy change of the system

b)

Decreases the entropy change of the system

c)

Has no effect on the entropy change of the system

d)

Cannot be determined

131.

Calculate the entropy change of the system for the reaction:

2Ca(s) + O₂(g) -----> 2CaO(s)

a)

ΔSθ (system) =

−208.4 J.K^-1.mol^-1

b)

ΔSθ (system) =

−200.4 J.K^-1.mol^-1

c)

ΔSθ (system) =

−208.4 J.K.mol^-1

d)

ΔSθ (system) =

−200.4 J.K.mol^-1

132.

Calculate the entropy change of the system for the reaction:

CH₄(g) + 2O₂(g) --------> CO₂ + 2H₂O

a)

ΔSθ (system) =

−5.2 J.K^-1.mol^-1

b)

ΔSθ (system) =

−52.2 J.K^-1.mol^-1

c)

ΔSθ (system) =

5.2 J.K^-1.mol^-1

d)

ΔSθ (system) =

52.2 J.K^-1.mol^-1

133.

Calculate the entropy change of the surroundings for the reaction

a)

ΔSθ (surroundings) =

+4.262 J.K^-1.mol^-1

b)

ΔSθ (surroundings) =

-4.262 J.K^-1.mol^-1

c)

ΔSθ (surroundings) =

+4.262 kJ.K^-1.mol^-1

d)

ΔSθ (surroundings) =

-4.262 kJ.K^-1.mol^-1

134.

Calculate the entropy change of the surroundings for the reaction

a)

ΔSθ (surroundings) =

+2.987 J.K^-1.mol^-1

b)

ΔSθ (surroundings) =

-2.987 J.K^-1.mol^-1

c)

ΔSθ (surroundings) =

+2.987 kJ.K^-1.mol^-1

d)

ΔSθ (surroundings) =

-2.987 kJ.K^-1.mol^-1

135.

If we look at Table 22.3, we can find some trends about absolute entropies. What can you infer about Sθ values? (Pick one correct answer)

a)

increase as A_r increases

b)

increase from solids to liquids to gases

c)

decrease if the element is soft

d)

All options are correct options

136.

What is ΔSθ for the reaction 2Na(s) + Cl₂(g) → 2NaCl(s)?

Remember: At 0 K, ΔS = 0 as ΔS for all the substances = 0

a)

ΔSθ = –123.2 J.K^–1.mol^–1

b)

ΔSθ = 411.6 J.K^–1.mol^–1

c)

ΔSθ = 123.2 J.K^–1.mol^–1

d)

ΔSθ = –411.6 J.K^–1.mol^–1

137.

Calculate ΔG^θ for the reaction:

C₂H₆(g) + 4O₂(g) → 2CO₂(g) + 3H₂O(l)

a)

ΔG^θ = -1467.5 kJ.mol^-1

b)

ΔG^θ = -1533.3 kJ.mol^-1

c)

ΔG^θ = 1467.5 kJ.mol^-1

d)

ΔG^θ = 1533.3 kJ.mol^-1

138.

Calculate the value of ΔG^θ at 1000 K.

Remember: ΔG^θ = ΔH^θ - TΔS^θ

a)

ΔG^θ = +141 kJ.mol^-1

b)

ΔG^θ = +336 kJ.mol^-1

c)

ΔG^θ = -141 kJ.mol^-1

d)

ΔG^θ = -336 kJ.mol^-1

139.

For the following processes 'a' & 'b', choose the correct reason why ΔH^θ and ΔS^θ have the signs (positive or negative) shown.

a)

(a) ΔH^θ is negative because relatively weak bonds in magnesium and oxygen are converted into strong ionic bonds in magnesium oxide. ΔS^θ is negative because of the difference between no. of moles of reactants and products is decreasing from left to right of the reaction.

(b) ΔH^θ is positive because a relatively weak π bond is being formed. ΔS^θ is positive because of the difference between no. of moles of reactants and products is increasing from left to right of the reaction.

b)

(a) ΔH^θ is negative because relatively strong metallic bonds in magnesium and weak bonds in oxygen are converted into strong ionic bonds in magnesium oxide. ΔS^θ is negative because of the difference between no. of moles of reactants and products is decreasing from left to right of the reaction.

(b) ΔH^θ is positive because a relatively weak π bond is being formed. ΔS^θ is positive because of the difference between no. of moles of reactants and products is increasing from left to right of the reaction.

c)

(a) ΔH^θ is negative because relatively strong metallic bonds in magnesium and strong double covalent bonds in oxygen are converted into strong ionic bonds in magnesium oxide. ΔS^θ is negative because of the difference between no. of moles of reactants and products is decreasing from left to right of the reaction.

(b) ΔH^θ is positive because a relatively strong double covalent bond is being formed. ΔS^θ is positive because of the difference between no. of moles of reactants and products is increasing from left to right of the reaction.

d)

(a) ΔH^θ is negative because relatively strong metallic bonds in magnesium and a weak π bond in oxygen are converted into strong ionic bonds in magnesium oxide. ΔS^θ is negative because of the difference between no. of moles of reactants and products is decreasing from left to right of the reaction.

(b) ΔH^θ is positive because a relatively weak π bond is being formed. ΔS^θ is positive because of the difference between no. of moles of reactants and products is increasing from left to right of the reaction.

140.

What is a redox reaction?

a)

A reaction where both reactants get reduced.

b)

A reaction where both reactants get oxidized.

c)

A reaction that involves a transfer of electrons between reactants.

d)

A reaction that involves the combustion of at least one reactant.

141.

What is reduction?

a)

It is the gain of electrons.

b)

It is the loss of electrons.

c)

It is the creation of electrons.

d)

It is the destruction of electrons.

142.

What is oxidation?

a)

It is the gain of electrons.

b)

It is the loss of electrons.

c)

It is the creation of electrons.

d)

It is the destruction of electrons.

143.

What is a reducing agent?

a)

It is the species that gets reduced -- gains electrons.

b)

It is the species that gets oxidized -- loses electrons.

c)

It is the species that creates electrons.

d)

It is the species that destroys electrons.

144.

What is an oxidizing agent?

a)

It is the species that gets reduced -- gains electrons.

b)

It is the species that gets oxidized -- loses electrons.

c)

It is the species that creates electrons.

d)

It is the species that loses electrons.

145.

What is a galvanic (voltaic) cell?

a)

It is a cell that destroys electrons on one side and creates electrons on the other side.

b)

It is a cell that contains only one metal bar and one aqueous ion solution.

c)

It is a type of battery that drives a redox reaction when electricity is applied.

d)

It is a type of battery that generates an electrical current from redox reactions.

146.

What is a standard half-cell?

a)

It is one cell that comprises half of a whole battery.

b)

It is a cell containing only a metal bar with no aqueous ion solution.

c)

It is a cell containing only an aqueous ion solution and no metal bar.

d)

It is a cell that generates half of a volt of electricity.

147.

What is a half-reaction?

a)

It is a reaction at equilibrium where half of the substances are reactants and half of the substances are products.

b)

It is the reaction that occurs when the switch of a galvanic cell is open.

c)

It shows EITHER the reduction component or the oxidation component of a redox reaction.

d)

It shows BOTH the reduction and oxidation components of a redox reaction.

148.

What is an anode?

a)

It is the electrode where oxidation takes place.

b)

It is the electrode where reduction takes place.

c)

It is an aqueous solution containing an electrode.

d)

It is the salt bridge that connects half-cells.

149.

What is a cathode?

a)

It is the electrode where oxidation takes place.

b)

It is the electrode where reduction takes place.

c)

It is an aqueous solution containing an electrode.

d)

It is the salt bridge that connects half-cells.

150.

Electrons always flow from

a)

cathode to anode

b)

anode to cathode

151.

Oxidation occurs at the

a)

anode

b)

cathode

152.

Electrons are gained at the

a)

anode

b)

cathode

153.

Given their standard reduction potentials, which of the species is going to be oxidized?

Cu²⁺/Cu = 0.34V

Zn²⁺/Zn = -0.76V

a)

Cu

b)

Zn

c)

CuSO4

d)

ZnSO4

154.

An iron nail is put into a solution of copper nitrate, iron is above copper in the activity series of metals, what will happen?

a)

iron will be reduced

b)

bubbles of oxygen gas will form on iron nail

c)

the iron nail will become copper plated

d)

no reaction occurs

155.

If Al is above Co in the activity series of metals, which of the following will occur if Al metal is put into a solution of cobalt nitrate?

a)

a redox reaction takes place

b)

the Al strip dissolves

c)

the Al strip becomes coated with cobalt

d)

all of the above

156.

What occurs to the mass of copper electrode in the following reaction?

Zn/Zn²⁺ // Cu²⁺/Cu

a)

increases

b)

decreases

c)

remains the same

157.

Which metal is the negative electrode?

Zn/Zn²⁺ // Cu²⁺/Cu

a)

zinc

b)

copper

158.

What reaction occurs at the anode?

Ag⁺/Ag = 0.80V

Ni²⁺/Ni = -0.25V

a)

Ag⁺ + e^- →Ag

b)

Ag → Ag⁺ + e^-

c)

Ni²⁺ + 2e^- → Ni

d)

Ni → Ni²⁺ + 2e^-

159.

What would be the theoretical cell potential of the previous electrochemical cell?

Ag⁺/Ag = 0.80V

Ni²⁺/Ni = -0.25V

a)

1.05V

b)

-1.05V

c)

0.55V

d)

-0.55V

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