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Unit 4 Part 2 OnRamps Chemistry

Total questions: 53

Worksheet time: 45mins

Name
Class
Date
1.

Is this an example of an isolated, closed, or open system?

a)

isolated

b)

closed

c)

open

2.

Is this an example of an isolated, closed, or open system?

a)

isolated

b)

closed

c)

open

3.

Is this an example of an isolated, closed, or open system?

a)

isolated

b)

closed

c)

open

4.

When work is being done by the system the sign is.

a)

positive

b)

negative

5.

When work is being done on the system the sign is.

a)

positive

b)

negative

6.

q represents _____.

a)

heat

b)

enthalpy

c)

reaction energy

d)

kinetic energy

7.

w represents ______

a)

work

b)

wasted energy

c)

Waltzmann's constant

d)

heat

8.

H represents ___ .

a)

Entropy

b)

Enthalpy

c)

Internal Energy

d)

Gibbs Free Energy

9.

S represents ___ .

a)

Entropy

b)

Enthalpy

c)

Internal Energy

d)

Gibbs Free Energy

10.

U represents ___ .

a)

Entropy

b)

Enthalpy

c)

Internal Energy

d)

Gibbs Free Energy

11.

G represents ___ .

a)

Entropy

b)

Enthalpy

c)

Internal Energy

d)

Gibbs Free Energy

12.

A reaction is spontaneous when...

a)

ΔGrxn\Delta G_{rxn}   is positive

b)

ΔSuniv\Delta S_{univ}  is positive

c)

ΔGrxn\Delta G_{rxn}   is negative

d)

ΔSuniv\Delta S_{univ}  is negative

13.

Which of the following reactions have an increase in entropy?

a)

2 H2 (g)+ O2 (g)  → 2 H2O (l)2\ H_{2\ \left(g\right)}+\ O_{2\ \left(g\right)}\ \ \rightarrow\ 2\ H_2O\ _{\left(l\right)}  

b)

Mg(s) +2 HCl (aq) → H2 (g) + MgCl2 (aq)Mg_{\left(s\right)}\ +2\ HCl\ _{\left(aq\right)}\ \rightarrow\ H_{2\ \left(g\right)\ }+\ MgCl_2\ _{\left(aq\right)}  

c)

H2O (l) → H2O (g)H_2O\ _{\left(l\right)\ }\rightarrow\ H_2O\ _{\left(g\right)}  

14.

Which of the following statements concerning the first law of thermodynamics is true

a)

The internal energy of the universe is conserved

b)

The change in internal energy of a system and the surroundings can have the same sign

c)

The entropy of the universe is always decreasing

d)

The internal energy changes are only determined by gases.

15.

A chromebook battery has a change in internal energy, ΔE\Delta E  , of 1000 kJ as it powers the computer. If 50 kJ is warming the desk it is sitting on, how much work was done to watch a video.

a)

1050 kJ

b)

-1050 kJ

c)

9050 kJ

d)

-9050 kJ

16.

What is the total change in internal energy when your cell phone uses 30 kJ to light up and releases 30 kJ of heat into your hand.

a)

60 kJ

b)

-60 kJ

c)

0 kJ

d)

900 kJ

17.

If the specific heat of water is 4 J/g∙°C, how much heat is required to increase the temperature of 10 g of water from 20 °C to 40 °C? (show your work)

a)

-800 J

b)

800 J

c)

800 KJ

d)

-800 KJ

18.

For a skillet, used for cooking, do you want a high or low specific heat

a)

High, so that it will need more energy to heat up

b)

Low, so that it will change temperature quickly

19.
A sample of iron receives 50.J of heat energy that raises the temperature of the iron by 25.0°C. If iron has a specific heat of 0.10 J/g°C, what is the mass of the iron sample?  (show your work)
a)
25g
b)
30g
c)
20g
d)
50g
20.

A reaction is performed in a beaker with a temperature probe recording the temperature changes of the surroundings.  If the temperature began at 15.0 degrees Celsius and ended at 27.5 degrees Celsius. Is the reaction endothermic or exothermic?

a)
Exothermic
b)
Endothermic
21.
In an exothermic process, the surroundings are gaining energy.
a)
True
b)
False
22.

The specific heat of water is 4J/g°C.

If 2400 J of energy is added to 6 g of water at 18.0 °C, what is the final temperature of the water?

a)

18 °C

b)

78 °C

c)

100 °C

d)

118 °C

23.

What does "ΔT" mean?

a)

A change in health

b)

A change in heat

c)

A change in height

d)

A change in temperature

24.

Using the equations below:

C(s) + O2(g) → CO2(g) ∆H = –400 kJ

Mn(s) + O2(g) → MnO2(s) ∆H = –500 kJ

what is ∆H (in kJ) for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g)

a)

900

b)

100

c)

-100

d)

-900

25.

Using the equations below

Cu(s) + 1/2O2(g) → CuO(s) ∆Hf = –150 kJ

2Cu(s) + 1/2O2(g) → Cu2O(s) ∆Hf = –170 kJ

what is the value of ∆H (in kJ) for the following reaction?

2CuO(s) → Cu2O(s) + 1/2O2(g)

a)

130

b)

20

c)

-20

d)

-130

26.

Consider the following equations.

Mg(s) + O2(g) → MgO(s) ∆H = –600 kJ

H2(g) + O2(g) → H2O(g) ∆H = –250 kJ

What is the ∆H value (in kJ) for the following reaction?

MgO(s) + H2(g) → Mg(s) + H2O(g)

a)

-850

b)

-350

c)

+350

d)

+850

27.

Which of the following statements are true for the reaction:

SO2(g) + 1/2O2(g) ↔ SO3(g)

ΔH = –92 kJ mol-1

Where ↔ indicates that the reaction can proceed in the forward and the reverse direction.

a)

The forward and reverse reaction both produce 92 kJ of energy.

b)

Oxidizing 2 moles of SO2 would produce twice as much energy.

c)

The reverse reaction has an enthalpy of +92 kJ mol-1.

d)

Collecting the SO3 produced in the liquid state would not change the measured enthalpy.

28.

The standard enthalpy change of formation values of two oxides of phosphorus are:

P4(s) + 3O2(g) → P4O6(s) ΔHf = –1600 kJ mol–1

P4(s) + 5O2(g) → P4O10(s) ΔHf = –3000 kJ mol–1

What is the enthalpy change, in kJ mol–1, for the reaction below?

P4O6(s) + 2O2(g) → P4O10(s)

a)

+4600

b)

+1400

c)

–1400

d)

–4600

29.

If N2 (g) + 2O2 (g) ⟶\longrightarrow   2NO2(g) has a ΔHrxn \Delta H_{rxn\ }  = 68, then what is the ΔHrxn\Delta H_{rxn}  if you reverse the reaction?

a)

86 kJ

b)

- 86 kJ

c)

68 kJ

d)

-68 kJ

30.

Which of the following statements must be true if ΔG \Delta G\  is negative.

a)

The reaction is nonspontaneous

b)

The reaction is exothermic

c)

The reaction is at equilibrium

d)

The entropy could be positive or negative

31.

What is the symbol for entropy?

a)

H

b)

G

c)

S

d)

E

32.

Entropy is a measure of

a)

accuracy

b)

precision

c)

possible microstates

d)

the attraction of a nucleus for an electron

33.

Which of these processes represents a decrease in entropy?

a)

sublimation of CO2

b)

boiling water

c)

condensation of steam on cold glass

d)

NaCl dissolving in water

34.

Which of these reactions shows a DECREASE in entropy?

a)

CaCO3(s) → CaO(s) + CO2(g)

b)

3O2(g) → 2O3(g)

c)

2NH3(g) → 3H2(g) + N2(g)

d)

C6H6(l) → C6H6(g)

35.

The entropy will usually increase when

a)

a molecule is broken into two or more smaller molecules

b)

a reaction occurs that results in an increase in the number of moles of gas

c)

a solid changes to a liquid

d)

all of these

36.

Which reaction has a +ΔS ?

a)

AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

b)

H2O(g) + CO2(g) → H2CO3(aq)

c)

H2(g) + I2(g) → 2Hl(g)

d)

C2H2O2(g) → 2CO(g) + H2(g)

37.

Which sample has the lowest entropy?

a)

1 mole of KNO3(l)

b)

1 mole of KNO3(s)

c)

1 mole of H2O(l)

d)

1 mole of H2O(g)

38.

Which phase change represents a decrease in entropy?

a)

solid to liquid

b)

gas to liquid

c)

liquid to gas

d)

solid to gas

39.

Which sample has the greatest entropy?

a)

1 mole of KNO3(l)

b)

1 mole of KNO3(s)

c)

1 mole of H2O(l)

d)

1 mole of H2O(g)

40.

Which sample has the greatest entropy?

a)

1 mole of CH4(l)

b)

1 mole of C3H8(s)

c)

1 mole of C2H4(l)

d)

1 mole of C6H12(g)

41.

Which sample has the greatest entropy?

a)

1 mole of CH4(l)

b)

1 mole of C3H8(s)

c)

1 mole of C2H4(l)

d)

1 mole of C6H12(g)

42.

Your enthalpy is high and negative but your entropy is also negative. What type of temperature would you want to get a spontaneous reaction?

a)

low

b)

high

43.

Your enthalpy is positive and your entropy is positive. What type of temperature would you want to ensure a spontaneous reaction.

a)

very high

b)

very low

44.

Which of the following situations demonstrates high entropy?

a)

water freezing

b)

water vaporizing

c)

steam condensing to water

45.

For which of the following processes would ΔS have a negative value (look at the states s, l, or g) ?


i. 2Fe2O3 (s) → 4Fe(s)+3O2(g)

ii. Mg2+(aq)+2OH-(aq) → Mg(OH)2(s)

iii. H2(g) + C2H4(g) → 3C2H6(g)

a)

I only

b)

I and II only

c)

I and III only

d)

II only

46.

NH3(g) + 2CH4(g) + 5/2 O2(g) → H2NCH2COOH(s) + 3H2O(l), ΔS is...

a)

increasing

b)

decreasing

c)

staying the same

d)

speeding up

47.

What is the equation for Gibbs Free Energy Change?

a)

ΔG = ΔH + TΔS

b)

ΔH = ΔG - TΔS

c)

ΔS = ΔH - TΔG

d)

ΔG = ΔH - TΔS

48.

1. Calculate the standard entropy change for the following reaction: Cu(s)+ 1/2 O2(g) → CuO(s). Given that

S0 (Cu(s)) = 30 J/K*mol

S0 (O2(g)) = 200 J/K*mol

S0 (CuO(s)) = 40 J/K*mol

a)

-200 J/K.mol

b)

90 J/K.mol

c)

-90 J/K.mol

d)

+200 J/K.mol

49.

For the process at 250C : I2(g) →I2(s). What are the signs for ΔG, ΔH and ΔS?

a)

ΔG ΔH ΔS

+ - -

b)

ΔG ΔH ΔS

- - -

c)

ΔG ΔH ΔS

- + +

d)

ΔG ΔH ΔS

- + +

50.

If a process is exothermic and not spontaneous, then what must be true?

a)

ΔSsys > 0

b)

ΔHsys > 0

c)

ΔGsys = 0

d)

ΔSuniv < 0

51.

In which of the following reactions do you expect to have a decrease in entropy?

a)

Fe(s) →Fe(l)

b)

H2O2(l) →2H2O(l) +O2(g)

c)

2Fe(s) +3/2O2(g)→Fe2O3(s)

d)

HF(l) →HF(g)

52.

If ΔH and ΔS are both negative or positive, then ΔG has a -------------------sign.

a)

Positive

b)

Negative

c)

Large

d)

Variable

53.
An ice cube is placed on a table and begins to melt.  The following conclusion can be made.
a)
Energy is being absorbed by the table which causes the cube to melt.
b)
Energy is being transferred from the table to the ice cube.
c)
Energy is being created by the air and table and making the cube heat up.
d)
Energy  is being created by the ice cube and is being released to the air

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