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ACHEM - testing > Chemistry Unit #6: Energetics.............................. Test on March 14, 2017 > Flashcards

Chemistry Unit #6: Energetics.............................. Test on March 14, 2017 Flashcards

1
Q

Define Exothermic 3

A

Is a chemical reaction that releases energy in the form of heat

Temperature of reaction increases 

May be shown as product A + B => C + D + heat
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2
Q

Define Calorimeter

A

A calorimeter is a device used to measure the heat flow of a chemical reaction or physical change.

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3
Q

Define Standard Enthalpy Change 7

A

Measures the amount of heat energy transferred under standard conditions

Enthalpy difference between products and reactants under standard conditions

Pressure is @100.0 kPa

Temperature is @298 K or 25°C

Standard conditions are indicated by °

It can't measure internal energy of substance 

It can only measure change in heat content 

It can measure the heat energy transferred for a chemical reaction
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4
Q

Define calorie (Not Calorie)

A

It is the amount of heat required to raise the temperature of one gram of water from 14.5°C to 15.5°C. 1 calorie = 4.184 joules.

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5
Q

Define Activation Energy 2

A

It is the minimum energy required for reaction to occur

The activation energy is usually represented by the symbol Ea
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6
Q

Define Endothermic 3

A

Is a chemical reaction that absorbs energy from the surroundings

Temperature of reaction decreases

May be shown as product A + B + heat => C + D

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7
Q

State what a calorimeter is and how it works (including how to adjust the temperature for heat loss) 3-4

A

A calorimeter can as simple as a Styrofoam container.

Using a thermometer, we can measure the hamount of heat involved in a chemical or physical process from the surroundings of the calorimeter.

To adjust the temperature for heat loss, we extrapolate the maximum temperature and get ∆T by subtracting T(extrapolated) by T(initial).

This means that the amount of heat produced or consumed in the reaction equals the amount of heat absorbed or lost by the solution.

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8
Q

Define Specific Heat Capacity

A

Is the specific amount of energy needed to raise the temperature of 1g of substance by 1 K or (1°C)

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9
Q

Define Heat 2

A

Heat is a measure of the total kinetic energy of a substance

The total energy depends on the amount of material present
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10
Q

Define Temperature 2

A

Measures average kinetic energy of the a substance

Number of particles is unimportant
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11
Q

When the solids Ba(OH)2 and NH4SCN are mixed, a solution is produced and the temperature drops.

Ba(OH)2(s) + 2NH4SCN(s) → Ba(SCN)2(aq) + 2NH3(g) + 2H2O(l)

Which statement about the energetics of this reaction is correct?

A. The reaction is endothermic and ∆H is negative.
B. The reaction is endothermic and ∆H is positive.
C. The reaction is exothermic and ∆H is negative.
D. The reaction is exothermic and ∆H is positive.

A

B

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12
Q

A sample of a metal is heated. Which of the following are needed to calculate the heat absorbed by the sample?

I. The mass of the sample
II. The density of the sample
III. The specific heat capacity of the sample

A. I and II only
B. I and III only
C. II and III only
D. I, II and III

A

B

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13
Q

Using the equations below:
Cu(s) + O2(g) → CuO(s)∆Hο = –156 kJ
2Cu(s) + O2(g) → Cu2O(s)∆Hο = –170 kJ

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

2CuO(s) → Cu2O(s) + O2(g)

A. 142 B. 15
C. –15 D. –142

A

A

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14
Q
  1. How much energy, in joules, is required to increase the temperature of 2.0 g of aluminium from 25 to 30°C? (Specific heat of Al = 0.90 J g–1 K–1).

A. 0.36
B. 4.5
C. 9.0
D. 54

A

C

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15
Q

Which statements about exothermic reactions are correct?

I. They have negative ∆H values.
II. The products have a lower enthalpy than the reactants.
III. The products are more energetically stable than the reactants.

A. I and II only
B. I and III only
C. II and III only
D. I, II and III

A

D

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16
Q

The average bond enthalpies for O—O and O==O are 146 and 496 kJ mol–1 respectively.
What is the enthalpy change, in kJ, for the reaction below?

H—O—O—H(g) ===> H—O—H(g) + ½O==O(g)

A. – 102 B. + 102
C. + 350 D. + 394

A

A

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17
Q

What energy changes occur when chemical bonds are formed and broken?

A. Energy is absorbed when bonds are formed and when they are broken.
B. Energy is released when bonds are formed and when they are broken.
C. Energy is absorbed when bonds are formed and released when they are broken.
D. Energy is released when bonds are formed and absorbed when they are broken.

A

D

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18
Q

Which statements are correct for an endothermic reaction?

I. The system absorbs heat.
II. The enthalpy change is positive.
III. The bond enthalpy total for the reactants is greater than for the products.

A. I and II only
B. I and III only
C. II and III only
D. I, II and III

A

A

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19
Q

Using the equations below:

C(s) + O2(g) → CO2(g)	∆H = –390 kJ Mn(s) + O2(g) → MnO2(s) 	∆H = –520 kJ

What is ∆H (in kJ) for the following reaction?
MnO2(s) + C(s) ===> Mn(s) + CO2(g)

A. 910
B. 130
C. –130
D. –910

A

B

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20
Q

When 40 joules of heat are added to a sample of solid H2O at –16.0°C the temperature increases to –8.0°C. What is the mass of the solid H2O sample?

[Specific heat capacity of H2O(s) = 2.0 J g–1K–1]

A. 2.5 g
B. 5.0 g
C. 10 g
D. 160 g

A

A

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21
Q

Which statement about bond enthalpies is correct?

A. Bond enthalpies have positive values for strong bonds and negative values for weak bonds.
B. Bond enthalpy values are greater for ionic bonds than for covalent bonds.
C. Bond breaking is endothermic and bond making is exothermic.
D. The carbon–carbon bond enthalpy values are the same in ethane and ethene.

A

C

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22
Q

What two types of reactions are always exothermic?

A

Combustion and neutralization

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23
Q

Is Ice melting an exothermic reaction?

A

No

24
Q

Consider the following equations.

Mg(s) +  O2(g) → MgO(s)    ∆Hο = –602 kJ H2(g) +  O2(g) → H2O(g)      ∆Hο = –242 kJ

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

MgO(s) + H2(g) → Mg(s) + H2O(g)

A. –844 B. –360
C. +360 D. +844

A

C

25
Q

For which of the following is the sign of the enthalpy change different from the other three?

A. CaCO3(s) → CaO(s) + CO2(g)
B. Na(g) → Na+(g) + e–
C. CO2(s) → CO2(g)
D. 2Cl(g) → Cl2(g)

A

D

26
Q

Separate solutions of HCl(aq) and H2SO4(aq) of the same concentration and same volume were completely neutralized by NaOH(aq). X kJ and Y kJ of heat were evolved respectively.
Which statement is correct?

A. X = Y B. Y = 2X
C. X = 2Y D. Y = 3X

A

B

27
Q

Which statement is correct for an endothermic reaction?

A. The products are more stable than the reactants and ∆H is positive.
B. The products are less stable than the reactants and ∆H is negative.
C. The reactants are more stable than the products and ∆H is positive.
D. The reactants are less stable than the products and ∆H is negative.

A

C

28
Q

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

P4(s) + 3O2(g) ===> P4O6(s) ∆HӨf= –1600 kJ mol–1

P4(s) + 5O2(g) ===> P4O10(s) ∆HӨf= –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

A

C

29
Q

Which statements are correct for all exothermic reactions?

I. The enthalpy of the products is less than the enthalpy of the reactants.
II. The sign of ∆H is negative.
III. The reaction is rapid at room temperature.

A. I and II only
B. I and III only
C. II and III only
D. I, II and III

A

A

30
Q 
The mass m (in g) of a substance of specific heat capacity c (in J g–1 K–1 ) increases by t°C. What is the heat change in J?
A.	mct 
B.	mc(t + 273)
C.	mct/1000
D.	[mc(t + 273)]/1000
A

A

31
Q

Define the term average bond enthalpy, illustrating your answer with an equation for methane, CH4. 2

A

Average value of the energy required to break a bond (average of each C-H in Ch4) {1} When molecule and atoms are in gaseous states {1}

32
Q

Using the equations below:

C(s) + O2(g) → CO2(g) ∆Hο = –394 kJ mol–1
Mn(s) + O2(g) → MnO2(s) ∆Hο = –520 kJ mol–1

What is ∆H, in kJ, for the following reaction?

MnO2(s) + C(s) → Mn(s) + CO2(g)

A. 914
B. 126
C. –126
D. –914

A

B

33
Q

Which combination is correct for a chemical reaction that absorbs heat from the surroundings?

Type of reaction	ΔH at constant pressure

A. Exothermic Positive
B. Exothermic Negative
C. Endothermic Positive
D. Endothermic Negative

A

C

34
Q

Consider the specific heat capacity of the following metals.

Metal	Specific heat capacity / J kg–1 K–1
Cu	385
Ag	234
Au	130
Pt	134

Which metal will show the greatest temperature increase if 50 J of heat is supplied to a 0.001 kg sample of each metal at the same initial temperature?

A. Cu
B. Ag
C. Au
D. Pt

A

C

35
Q

The temperature of a 2.0 g sample of aluminium increases from 25°C to 30°C.
How many joules of heat energy were added? (Specific heat of Al = 0.90 J g–1K–1)

A. 0.36
B. 2.3
C. 9.0
D. 11

A

C

36
Q

Consider the following reactions.

S(s) + O2(g) ===> SO3(g) ∆HӨ = -395 kJ mol1

SO2(s) + O2(g) ===> SO3(g) ∆HӨ = -98 kJ mol1

What is the ∆HӨ value (in kJ mol–1) for the following reaction?

	S(s) + O2(g) ===> SO2(g)

A. –297
B. +297
C. – 493
D. +493

A

A

37
Q

Which statement is correct for an endothermic reaction?

A. Bonds in the products are stronger than the bonds in the reactants.
B. Bonds in the reactants are stronger than the bonds in the products.
C. The enthalpy of the products is less than that of the reactants.
D. The reaction is spontaneous at low temperatures but becomes non-spontaneous at high temperatures.

A

B

38
Q

Comparing a 100 ml solution with the temperature of 10°C, to a 100 ml solution with the temperature of 20°C, what solution has the highest temperature and has the highest heat? 2

A

The 100 ml solution with the temperature of 20°C has the highest temperature

The 100 ml solution with the temperature of 20°C has the highest heat

39
Q

Describe the changes in bonding that take place during the combustion of C₃H₈. [2]

A

[C₃H₈ + O₂] bonds are broken, requiring energy as they separate.

[H₂O + CO₂] bonds are formed, releasing energy as they form.

40
Q

Describe an experiment to determine the enthalpy change of the reaction between dilute hydrochloric acid and aqueous sodium hydroxide. Show how the value of ΔH would be calculated from the data obtained. [10]

A

Take a known volume of sodium hydroxide solution

Take a known molarity/concentration of sodium hydroxide solution

Record the temperature of the sodium hydroxide solution

Place sodium hydroxide solution in an isolated vessel

Add the solution of HCI

Such that an equal/excess number of moles/known quantity is added

Stir/mix the solution

Record the final tempearature

HCI + NaOH ===> NaCl + H₂O

ΔH = Total mass (or volume) X temperature change X specific heat capacity

Divide ΔH by the # of mole of limiting reactant = kJmol

41
Q

In an enthalpy level diagram, what does the symbol R signifies?

A

Initial Enthalpy of Reactants

42
Q

Comparing a 100 ml solution with the temperature of 10°C, to a 500 ml solution with the temperature of 10°C, what solution has the highest temperature and has the highest heat? 2

A

The temperature for both solutions are the same

The 500 ml solution with the temperature of 10°C has the highest heat

43
Q

A negative ΔH is an _______ reaction.

A

Exothermic

Heat was absorbed by the sorroundings

44
Q

What is the symbol for Activation energy?

A

Ea

45
Q

In an enthalpy level diagram, what does the symbol P signifies?

A

Final Enthalpy of Products

46
Q

What reaction is it where Products are more stable than the reactants?

A

Exothermic reaction

47
Q

A positive ΔH is an _______ reaction.

A

Endothermic

Heat was absorbed from the sorroundings

48
Q

In an enthalpy level diagram, what does the symbol ΔH signifies?

A

Change in Enthalpy

49
Q

What are conclusions from Enthalpy Level Diagrams? 6

A

Stronger Bonds release more energy when formed and are more difficult to break

 In an exothermic reaction, Bonds in products are stronger than in reactants

      Products are more stable than the reactants

 In an endothermic reaction, Bonds in products are less strong than in reactants

      Reactants are more stable than products

      Generally need extra energy to proceed
50
Q

What are the limitations of Bond Enthalpies? 2

A

➫ Values from table are only suitable for substances in gaseous states

      Producing a liquid or solid would release extra energy

           Recall energy changes involved in altering states of matter

      Breaking bonds in solids or liquids require more energy than stated

           First must become gaseous

➫ The values are only average bond enthalpies

      Average obtained by considering a number of similar compounds

           Neighboring atoms will affect the exact energy value of a bond
51
Q

How do you calculate ΔT in a graph?

A

ΔT = T(extrapolated) - T(initial)

52
Q

What are we assuming in when graphing a Calorimeter Data?

A

The rate of heat loss after reaction is constant

53
Q

What reaction is it where Reactants are more stable than the Products?

A

Endothermic reaction

54
Q

State Hess’ Law and one reason why it could be used. 2-3

A

Hess’ Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps.

Hess’s Law lets us find the enthalpy change of a reaction we cannot measure using other reactions that lead to the one we can’t measure because we know their values.

Official definition: The enthalpy change accompanying a chemical reaction is independent of the pathway between the initial and final steps.

55
Q

What is an assumption we make about the heat of the reaction in a calorimeter? 2-3

A

We assume that all heat of the reaction was transferred to/from the water.

 It cannot actually occur

We assume that the rate of energy change that occurs is the same for the entire time

56
Q

What are we assuming for Graphing Calorimeter Data?

A

We assume that the rate of heat loss for the calorimeter is constant

57
Q

What are the two 0= _______ + _____________?

A

0 = energy lost + energy gained

0 = ΔH(reaction) + ΔH(solution)

ACHEM - testing (5 decks)

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