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Chemistry 11th Ch 7-11

Total questions: 100

Worksheet time: 2hrs 44mins

Name
Class
Date
1.
Heat flows from _______ to ______ objects
a)
cooler to warmer
b)
warmer to warmer
c)
warmer to cooler
d)
cooler to cooler
2.
Energy can not be created or destroyed, it can only be transferred refers to the....
a)
law of conservation of energy
b)
law of conservation of mass
c)
life facts
d)
law of conservation of chemistry
3.
Heat travels from the sun to the earth by the process of...
a)
conduction
b)
convection
c)
radiation
d)
insulation
4.

Which of the following is an example of an endothermic process or reaction?

a)

condensing

b)

freezing

c)

combustion

d)

melting

5.

Which of the following is an example of an exothermic process or reaction?

a)

melting

b)

boiling

c)

photosynthesis

d)

freezing

6.

.


How does an exothermic reaction compare with an endothermic reaction?

a)

Both changes absorb energy as reactants become products.

b)

Both changes release energy as reactants become products.

c)

An exothermic reaction absorbs energy, and an endothermic reaction releases energy.

d)

An endothermic reaction absorbs energy, and an exothermic reaction releases energy.

7.
The symbol for specific heat is .......
a)
c
b)
Q
c)
m
d)
t
8.
What  is the formula to calculate heat energy required to raise the temperature of any substance?
a)
Q=mc∆t
b)
Q=mc
c)
Q= ½mv
d)
m=QC
9.

The specific heat of platinum is 0.133 J/g°C. How much heat(Q) is released when a 10 g piece of platinum cools from 100°C to 50°C?

a)

66.5 J

b)

665 J

c)

0.0266 J

d)

0.665 J

10.
The temperature of an unknown piece of metal with a mass of 30.00 g changes from 25.0 °C to 35.0 °C when the metal absorbs 350.0 J of energy. What is the specific heat of the metal?
a)
1.17 J/g°C
b)
-1.17 J/g°C
c)
0.857 J/g°C
d)
-0.857 J/g°C
11.

20 g of water. specific heat of water is 4.18 J/g°C. temperature changes from 25° C to 20° C, how much heat energy (Q) moves from the water to the surroundings?

a)

418 J

b)

209 J

c)

83 J

d)

4.18 J

12.

What information do we need to perform a calculation of Gibbs Free Energy?

a)

Enthalpy of the reaction.

b)

Entropy of the reaction.

c)

The temperature of the reaction in Kelvin.

d)

All of the above are needed.

13.
A reaction has a positive ΔH and a positive ΔS. Which of the following is true?
a)
It will be spontaneous at all temperatures.
b)
It will be nonspontaneous at all temperatures.
c)
It will be spontaneous at low temperatures.
d)
It will be spontaneous at high temperatures.
14.
When ammonium nitrate dissolves in water, the solution gets cold. Which is true?
a)
Reaction is ENDOTHERMIC with positive ΔH
b)
Reaction is ENDOTHERMIC with -ΔH
c)
Reaction is EXOTHERMIC with a -ΔH
d)
Reaction is EXOTHERMIC with +ΔH
15.

how many laws are there in thermodynamics

a)

2

b)

3

c)

4

d)

6

16.

The second law of Thermodynamics define

a)

Entropy

b)

Energy

c)

Enthalpy

d)

Heat

17.

According to kinetic theory of gases, the absolute zero temperature is attained when

a)

When mass is zero

b)

Kinetic Energy of gas molecules is zero

c)

Pressure of gas molecules is zero

d)

Volume of gas is zero

18.

A concept that addresses the conservation of energy

a)

2nd law of thermodynamics

b)

1st law of thermodynamics

c)

3rd law of thermodynamics

19.

A system's organization is measured by ______

a)

equilibrium

b)

internal energy

c)

entropy

20.

Transfer of heat from low temperature to high temperature is

a)

None of these

b)

Possible by doing some external work

c)

Possible by keeping both the bodies in contact

d)

Impossible

21.

A piece of ice is added to water in a cup. The entropy

a)

decreased

b)

increased

c)

no change

d)

decreases then increases.

22.

The diagram represents two solids and the temperature of each. What occurs when the two solids are placed in contact with each other?

a)

Heat energy flows from solid A to solid B. Solid A decreases in temperature

b)

Heat energy flows from solid A to solid B. Solid A increases in temperature

c)

Heat energy flows from solid B to solid A. Solid B decreases in temperature.

d)

Heat energy flows from solid B to solid A. Solid B increases in temperature.

23.

Find the enthalpy of formation for the following chemical equation.

CH4 (g) + 2 O2 (g) --> CO2 (g) + 2 H2O (g)

a)

-802.3 kJ

b)

802.3 kJ

c)

951.9 kJ

d)

-951.9 kJ

24.

Find the enthalpy change that occurs over the following two-step equations.

2 H2 (g) + O2 (g) --> 2 H2O (l) ΔH = -572 kJ

2 H2O2 (l) --> 2 H2 (g) + 2 O2 (g) ΔH = 376 kJ

a)

-196 kJ

b)

-948 kJ

c)

196 kJ

d)

948 kJ

25.

Match the definition below to the correct term.

The enthalpy change that takes place when one mole of a compound is formed from its elements in their standard states under standard conditions.

a)

The enthalpy of neutralisation

b)

The enthalpy of combustion

c)

The enthalpy of formation

d)

The enthalpy of reaction

26.

How do you calculate Enthalpy of a reaction?

a)

ΔH = ΔHproducts - ΔHreactants

b)

ΔT = q / mC

c)

ΔG = ΔH -TΔS

d)

E = mc2

27.

"Heat change when 1 mole of gaseous atom is formed from its element at standard states" is the definition for _______.

a)

standard enthalpy of formation

b)

standard enthalpy ofhydration

c)

electron affinity

d)

standard enthalpy of atomisation

28.

Name the type of enthalpy for the following reaction:

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

a)

Standard enthalpy of neutralisation, ΔHneuto

b)

Standard enthalpy of solution, ΔHsolno

c)

Standard enthalpy of hydration, ΔHhydo

d)

lattice energy, ΔHlatticeo

29.

Define standard enthalpy of combustion.

a)

Heat released when one mole of substance is burnt completely in excess oxygen.

b)

Heat absorbed when one mole of substance is burnt completely in excess oxygen under standard state.

c)

Heat released when one mole of substance is burnt completely in excess oxygen under standard state.

d)

Heat change when one mole of substance is burnt partially in excess oxygen under standard state.

30.
What is the specific heat of an unknown substance if 100.0 g of it at 200.0 °C reaches an equilibrium temperature of 27.1 °C when it comes in contact with a calorimeter of water.  The water weighs 75. g and had an initial temperature of 20.00 °C?  (Specific heat of water is 4.18 J/g°C) (show your work)
a)
0.111 J/g°C
b)
1.29 J/g°C
c)
0.129 J/g°C
d)
22225.85 J
31.

2016 Q10 (80% correct ans)


A student calibrated a calorimeter using an electric heating coil. A current of 1.50 A with a potential difference of

4.50 V was applied for two-and-a-half minutes. A digital probe recorded a temperature rise of 5.35 °C.

The value of the calibration factor, in J °C–1, is

a)

189

b)

42.1

c)

3.15

d)

0.317

32.
For the reaction...
SO2 + O2  <−>  SO3
If the concentration of SO2 is increased, the equilibrium of the reaction will shift ___________.
a)
left
b)
right
c)
left and right 
d)
neither left nor right
33.
For the reaction...
SO2 + O2 <−>  SO3
If the equilibrium shifts to the right, the concentration of O2 will ___________.
a)
increase
b)
decrease
c)
remain the same
d)
double
34.
For the reaction...
heat  +  N2  +  O2  <−>  2NO
If  NO is removed  from the system, the concentration of N2 will _______.
a)
increase
b)
decrease
c)
remain the same
d)
double
35.
For the reaction...
heat  +  N2  +  O2  <−>  2NO
If  NO is removed  from the system, the concentration of N2 will _______.
a)
increase
b)
decrease
c)
remain the same
d)
double
36.
For the reaction...
heat  +  N2  +  O2  <−>  2NO
If  NO is removed  from the system, the concentration of N2 will _______.
a)
increase
b)
decrease
c)
remain the same
d)
double
37.
When K is large, 
a)
There are more products than reactants when the reaction reached equilibrium.
b)
There are more reactant than products when the reaction reached equilibrium.
c)
The amount of reactants is equal to the amount of products.
38.

What would be the equilibrium expression based on the following chemical equation:

jA+kB → lC+mDjA+kB\ \rightarrow\ lC+mD  

a)

K=[A]j[B]k[C]l[D]mK=\frac{\left[A\right]^j\left[B\right]^k}{\left[C\right]^l\left[D\right]^m}  

b)

K=[C]l[D]m[A]j[B]kK=\frac{\left[C\right]^l\left[D\right]^m}{\left[A\right]^j\left[B\right]^k}  

c)

K=[C]l[B]k[A]j[D]mK=\frac{\left[C\right]^l\left[B\right]^k}{\left[A\right]^j\left[D\right]^m}  

d)

K=[D]m[B]k[A]j[C]lK=\frac{\left[D\right]^m\left[B\right]^k}{\left[A\right]^j\left[C\right]^l}  

39.
What are the two factors to look for when determining if the reaction is at equilibrium?
a)
Forward reaction rate is faster than the reverse and concentrations are equal
b)
Forward and reverse reaction rates are equal and concentration is constant
c)
Forward and revers reaction rates are equal and concentration is equal
d)
Forward reaction rate is faster than the reverse and concentration is equal
40.
What are the two factors to look for when determining if the reaction is at equilibrium?
a)
Forward reaction rate is faster than the reverse and concentrations are equal
b)
Forward and reverse reaction rates are equal and concentration is constant
c)
Forward and revers reaction rates are equal and concentration is equal
d)
Forward reaction rate is faster than the reverse and concentration is equal
41.

HCl(aq) + H2O(l) ⇄ H3O+(aq) + Cl−(aq)

In 1.0 M HCl(aq), HCl is nearly 100 percent dissociated, as represented by the equation above. Which of the following best helps to explain why, in 0.01 M HCN(aq), less than 1 percent of HCN is dissociated?

a)

The CN− ion is not very soluble in water, and a solid precipitate would form if more of the HCN dissociated.

b)

Compared to the HCl(aq) solution, the concentration of the HCN(aq) solution is much too dilute to achieve 100 percent dissociation.

c)

The equilibrium constant for the dissociation of HCN(aq) is much smaller than that for the dissociation of HCl(aq).

d)

HCN(aq) reacts with water to form a basic solution, and the high concentration of OH−(aq) interferes with the dissociation process.

42.
What is the Kc expression for this reaction?
   2 NO(g)  +  O2(g) ⇌  2 NO2(g)
a)
Kc = [NO2]2 / [NO]2 [O2]
b)
Kc =  [NO]2 [O2] / [NO2]2
c)
Kc = [NO]2 [O2] [NO2]2
d)
Kc = [NO2]2 / [NO]2 +  [O2]
43.
What are [A] and [B]?
a)
concentration of reactants
b)
concentration of products
c)
energy of reactants
d)
energy of products
44.
What are [A] and [B]?
a)
concentration of reactants
b)
concentration of products
c)
energy of reactants
d)
energy of products
45.

A chemical equilibrium may be established by starting a reaction with ____________

a)

reactants only

b)

products only

c)

equal quantities of reactants and products

d)

any quantities of reactants and products

e)

all of the above

46.

1. The nationality of Henri-Louis Le Chatelier

a)

German

b)

American

c)

French

47.

3. In change in pressure and volume, if the pressure decreases the system will shift to the?

a)

location of more moles

b)

shift to the left

c)

no shift

48.

4. When a system at equilibrium is subject to change the system tries to minimize the change by?

a)

dissappearing

b)

going to the direction where it will relieve the stress

c)

expand

49.

5. Reversible reaction is?

a)

cannot be changed

b)

stays the same

c)

the reactants will produce products from products it can go back to be reactants

50.

2SO2(g)+O2(g) ⇌ 2SO3(g) + Heat


Adding SO2(g) will

a)

shift equilibrium right

b)

shift equilibrium left

c)

increase rate of reaction

d)

have no change

51.

2SO2(g)+1O2(g) ⇌ 2SO3(g) + Heat


Increasing the volume of the container will...

a)

shift equilibrium right

b)

shift equilibrium left

c)

slow rate of reaction

d)

have no change

52.
What are the two factors to look for when determining if the reaction is at equilibrium?
a)
Forward reaction rate is faster than the reverse and concentrations are equal
b)
Forward and reverse reaction rates are equal and concentration is constant
c)
Forward and revers reaction rates are equal and concentration is equal
d)
Forward reaction rate is faster than the reverse and concentration is equal
53.
When the rate of the forward reaction is equal to the rate of backward reaction, the system is said to be in
a)
Chemical Equilibrium
b)
Chemical Balance
c)
Stoichiometry
d)
Chemical Reaction
54.
When the rate of the forward reaction is equal to the rate of backward reaction, the system is said to be in
a)
Chemical Equilibrium
b)
Chemical Balance
c)
Stoichiometry
d)
Chemical Reaction
55.
Which of the following is NOT true at equilibrium?
a)
The forward and reverse reactions proceed at the same rate.
b)
The concentrations of reactants and products do not change.
c)
The concentration of the reactants is equal to the concentration of the products.
d)
The forward and reverse reactions continue to occur.
56.

Which law states that energy cannot be created nor destroyed, but it can be transformed or transferred?

a)

Dalton's Law

b)

law of conservation of energy

c)

Hess's Law

d)

E=mc^2

57.

What happens when kinetic energy increases?

a)

Temperature increases.

b)

Temperature decreases.

c)

Temperature is not affected by kinetic energy.

58.

A sample of a compound has a mass of 30 g and a specific heat of 0.258 J/g*K. If its temperature drops 40 K, how much heat was released? ( q=mcΔTq=mc\Delta T  )

a)

309.6 J

b)

-309.6 J

c)

465 J

d)

-465 J

59.

In the following reaction, A + 2B --> 3C, ΔH = 640 kJ\Delta H\ =\ 640\ kJ  . How much energy is needed to react 4 moles of A?

a)

2560 kJ

b)

1280 kJ

c)

1920 kJ

d)

3840 kJ

60.
Which is an example of a solute?
a)
Egg whites
b)
Sugar
c)
Water
d)
Acetone
61.
Which is an example of a solvent?
a)
Salt
b)
Egg whites
c)
Water
d)
Sugar
62.

Which of the following statements is held true about the common-ion effect on the equilibrium of solubility of salt ?

(you may choose more than one answer)

a)

Adding common-ion to the solution containing dissolved salt cause the solubility equilibrium shifts to the formation of undissolved salt

b)

Adding common-ion to the solution containing dissolved salt cause the numerical value of Ksp of corresponding salt changes

c)

Solubility of salt in solution containing common-ion is greater than that of pure water.

63.

At which solution will PbI2 have the lowest molar solubility ?

(Ksp PbI2 = 7.1 x 10-9)


PbI2(s) ↔ Pb2+(aq) + 2I-(aq)

a)

0.02 M Pb(NO3)2 solution

b)

0.02 M CaI2 solution

c)

0.02 M Pb(ClO4)2 solution

d)

0.02 M KI solution

64.

At which solution will Zn(OH)2 have the lowest molar solubility ?

(Ksp Zn(OH)2 =1.2 x 10-27)


Al(OH)3(s) ↔ Al3+(aq) + 3OH-(aq)

a)

buffered solution of pH 12

b)

10-3 M KOH solution

c)

10-3 M ZnCl2 solution

d)

10-2 M ZnSO4 solution

65.

if you add an acid to a base what will happen?

a)

BOOM

b)

it will neutralize

c)

it wont mix

66.
The pH of a solution is tested, and it is found to be a basic solution. Of the following choices, what could the pH have been?
a)
3
b)
9
c)
7
d)
5
67.

What might happen if buffers did not exist within the human body?

a)

Our blood and other bodily fluids might become too acidic or basic.

b)

Our stomach acid would not be able to break down food.

c)

We would not be able to process glucose within our cells.

d)

We would not be able to inhale oxygen into our lungs.

68.

Two beakers have liquids in them. When you add a base such as ammonia to beaker 1, the pH did not change. But when you add ammonia to beaker 2, the pH changes drastically. Choose the best explanation.

a)

Beaker 1 had only water and Beaker 2 contained a buffer.

b)

Beaker 1 contained a buffer and Beaker 2 contained only water.

c)

Beaker 2 contained an acid.

d)

The beaker that changed pH contained a buffer.

69.

moles of solute/liters of solution

a)

Solubility

b)

Molarity

c)

Solution

d)

Solute

70.

The measure of how much exposed area a solid object has, expressed in square units

a)

Surface Area

b)

Temperature

c)

Solubility

d)

Supersaturated

71.

The measure of how much exposed area a solid object has, expressed in square units

a)

Surface Area

b)

Temperature

c)

Solubility

d)

Supersaturated

72.

A mixture that is not uniform in composition; components are not evenly distributed throughout the mixture

a)

Heterogeneous mixture

b)

Unsaturated

c)

Concentration

d)

Homogeneous mixture

73.

To stir or mix a solution as to increase the particle movement of the solute particles in solution

a)

Agitation

b)

Unsaturated

c)

Solution

d)

Concentration

74.

the substance that is dissolved in a solution.

a)

Solvent

b)

Solubility

c)

Solute

d)

Solution

75.

A solution that contains less than the maximum amount of dissolved solute in a concentration.

a)

Unsaturated

b)

Mixtures

c)

Saturated

d)

Supersaturated

76.
Which is an example of a solution?
a)
Chex Mix
b)
Mixed fruit
c)
Juice
d)
milk and cereal
77.

Which of the following is NOT TRUE about Heterogeneous Mixtures?

a)

They will settle out over time

b)

Individual particles are often distinguishable

c)

Solutions are a type of heterogenous mixture

d)

Emulsions, Suspensions, and Collides are types of Heterogeneous Solutions.

78.
When materials combine to form a mixture, they
a)
A. keep their original properties.
b)
B. react to form a new substance with new properties.
c)
C. combine in a specific ratio.
d)
D. always change their physical state.
79.

Solubility refers to the ____ of solute that can dissolve in a certain volume or mass of solvent, at a certain temperature.

a)

Mean

b)

Range

c)

Amount

d)

Mode

80.

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.

81.

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.

82.

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.

83.

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.

84.

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.

85.

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.

86.

Which of the following factors can affect the value of the activation energy of a reaction?

1. the presence of a catalyst

2. changes in temperature

3. changes in the concentration of the reactants

a)

1 only

b)

1 and 2 only

c)

3 only

d)

1,2, and 3

87.

When the pressure of a fixed mass of gaseous reactants is raised at a constant temperature, the rate of reaction increases.

Which of the following statements explain this observation?

(You can choose more than 1 answer)

a)

rasing the pressure lowers the activation energy

b)

more molecules have energy greater than the activation energy at the higher pressure

c)

More collisions occur per second when the pressure is increased

88.

What is the main reason for the increase in reaction rate with increasing temperature?

a)

The activation energy decreases

b)

The activation energy increases

c)

The molecules collide more frequently

d)

More molecules have an energy greater than the activation energy

89.

Which of the following statements regarding the half-life of first order reaction is true?

a)

The half-life is affected by pressure.

b)

The half-life depends on the concem9of the reactant.

c)

The half-life is shorter at lower temperature.

d)

At the end of fourth half-life, the percentage of element left is 6.25% of the original quantity.

90.

Which of the following statements about the rate constant for a chemical reaction is true?

a)

The unit is s-1.

b)

If does not change if a catalyst is added.

c)

It remains constant at a constant temperature.

d)

It decreases as the concentration of the reactant decreases.

91.

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

92.
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
93.
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
94.
Which of the following increase the reaction rate?
a)
less surface area
b)
lower temperature
c)
an inhibitor
d)
increased concentration
95.
Products will form faster if____________
a)
the particle size of the reactants are larger.
b)
temperature is decreased.
c)
concentration of the reactants are increased.
d)
the reaction is not stirred.
96.

In a closed system...

a)

only matter can transfer

b)

only energy can transfer

c)

both energy and matter can transfer

d)

nothing can transfer

97.

You put your hot pan in the sink and added running cold water, how is the heat transferred?

a)

Heat from the pan flows into water

b)

Coldness of water takes away heat

c)

Coldness of water gets into the pan

d)

temperature does not change

98.

When a candle is burning, which type of reaction is it? What is the ΔH?

a)

Endothermic / positive ΔH

b)

Endothermic / negative ΔH

c)

Exothermic / positive ΔH

d)

Exothermic / negative ΔH

99.

In the equation q = m ⋅\cdot c ⋅\cdot  ΔT

What does c stand for?

a)

calories

b)

catalyst

c)

specific heat

d)

calcium

100.

A can of coke is a(n)...

a)

open system

b)

closed system

c)

isolated system

d)

individual system

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