Saturday Study Session 1, 3Rd Class Student Handout

Saturday Study Session 1, 3rd Class Student Handout

Thermochemistry

Multiple Choice

Identify the choice that best completes the statement or answers the question.

1.C2H4(g) + 3 O2(g)  2 CO2(g) + 2 H2O(g)

For the reaction of ethylene represented above, Hrxn is -1323 kJ.

What is the value of H if the combustion produced liquid water H2O(l), rather than water vapor H2O(g)?

(H for the phase change H2O(g) H2O(l) is -44 kJ mol-1.)

A) / -1235 KJ
B) / -1279 kJ
C) / -1323 kJ
D) / -1411 kJ

2.What is the standard enthalpy change H°, for the reaction:

3C2H2(g)  C6H6(g)

H°f of C2H2(g) is 230 kJmol-1

H°f of C6H6(g) is 83 kJmol-1

A) / -607 kJ
B) / -147 kJ
C) / -19 kJ
D) / +19 kJ

3.True for the evaporation of water at 1 atm and 25 C.

A) / H>0
B) / H<0
C) / H=0
D) / 

4.A 10 g sample of a metal was heated to 100°C and then quickly transferred to an insulated

container holding 100 g of water at 20°C. The temperature of the water rose to reach a final

temperature of 50°C. Calculate the heat absorbed by the water. Specific heat of water is 4 J/(gx°C)

A) / 400 kJ
B) / 12 kJ
C) / 8 kJ
D) / 1.2 kJ

The dissolution of an ionic solute in a polar solvent can be imagined as occurring in three steps, as shown in the figure above. In step 1, the separation between ions in the solute is greatly increased, just as will occur when the solute dissolves in the polar solvent. In step 2, the polar solvent is expanded to make spaces that the ions will occupy. In the last step, the ions are inserted into the spaces in the polar solvent. Which of the following best describes the enthalpy change, H, for each step?

A) / Step 1 and Step 3 are exothermic and Step 2 takes no energy so is neither endothermic nor exothermic.
B) / Step 1 and 2 add together to be 2X Step 3.
C) / Steps 2 and 3 add together to be Step 1
D) / Step 3 releases energy while Step 1 and 2 absorb energy

Na(s) + ½ Br2(g)  NaBr(s) H = -361 kJ/molrxn

The elements Na and Br react directly to form the compound NaBr according to the equation above. Refer to the information above and the table below to answer the questions 6-8.

Process / H (kJ/molrxn)
Na(s)  Na(g) / m
Na(g)  Na+(g) + e- / n
Br2(g)  2 Br(g) / p
Br (g) + e-  Br-(g) / q
Na+(g) + Br-(g)  NaBr(s) / r

6.How much heat is released or absorbed when 0.050 mol of Br2(g) is formed from NaBr(s)?

A) / 72.2 kJ is released / C) / 36.1 kJ is absorbed
B) / 36.1 kJ is released / D) / 72.2 kJ is absorbed

Short FRQ #1

A sample of CH3CH2NH2 is placed in an insulated container, where it decomposes into ethene and ammonia

according to the reaction represented above.

(a) Using the data in the table below, calculate the value, in kJ/molrxn , of the standard enthalpy change, ΔH°,

for the reaction at 298 K.

Bond / C–C / C = C / C–H / C–N / N–H
Average Bond Enthalpy (kJ/mol) / 348 / 614 / 413 / 293 / 391

(b) Based on your answer to part (a), predict whether the temperature of the contents of the insulated container will increase, decrease, or remain the same as the reaction proceeds. Justify your prediction.

FRQ #2
MgO(s) + 2 H+(aq) Mg2+(aq) + H2O(l)

3. A student was assigned the task of determining the enthalpy change for the reaction between solid MgO and aqueous HCl represented by the net-ionic equation above. The student uses a polystyrene cup calorimeter andperforms four trials. Data for each trial are shown in the table below.

Trial / Volume of 1.0 M HCl (mL) / Mass of MgO(s) Added (g) / Initial Temperature of Solution(°C) / Final Temperature of Solution (°C)
1 / 100.0 / 0.25 / 25.5 / 26.5
2 / 100.0 / 0.50 / 25.0 / 29.1
3 / 100.0 / 0.25 / 26.0 / 28.1
4 / 100.0 / 0.50 / 24.1 / 28.1

(a) Which is the limiting reactant in all four trials, HCl or MgO? Justify your answer.

(b) The data in one of the trials is inconsistent with the data in the other three trials. Identify the trial with

inconsistent data and draw a line through the data from that trial in the table above. Explain how you

identified the inconsistent data.

For parts (c) and (d), use the data from one of the other three trials (i.e., not from the trial you identified in part(b) above). Assume the calorimeter has a negligible heat capacity and that the specific heat of the contents of thecalorimeter is 4.18 J/(g°C). Assume that the density of the HCl(aq) is 1.0 g/mL.

Include units with your answer.

(d) Determine the student’s experimental value of H°for the reaction between MgO and HCl in units of

kJ/molrxn.

(e) Enthalpies of formation for substances involved in the reaction are shown in the table below. Using the

information in the table, determine the accepted value of H°for the reaction between MgO(s) andHCl(aq).

Substance / Hf°(kJ/mol)
MgO(s) / -602
H2O(l) / -286
H+(aq) / 0
Mg2+(aq) / -467

(f) The accepted value and the experimental value do not agree. Suggest a possible error in the procedure and

explain.

Thermochemistry – Entropy and Free Energy

Quick Notes:

Potential energy- energy of bonds

Kinetic energy-energy of motion (molecules moving) KE=1/2 mv2

Three laws:

1. Energy is neither created or destroyed

2. The disorder of the universe is increasing

3. The entropy of a perfect crystal at 0K is zero (why ∆S is not zero)

Heat is a transfer of energy.

Temperature reflects motion of particles

Enthalpy is heat content at constant pressure (put mol-1 on answers)

Pay attention to definitions of ∆H. Some are based on the formation of one mole which sometimes changes a reaction.

Endothermic-- feels cold; system is absorbing heat (+H)

Sketch an endothermic graph. Label reactants, products, EA, ∆H

Exothermic – feels hot; system is releasing heat (-H)

Sketch an exothermic graph. Label reactants, products, EA, ∆H

Entropy is the measure of chaos; -S disorder leaving, +S disorder entering

Gibbs free energy: -G spontaneous, +G nonspontaneous

Formulas

∆E = q (heat) + w(work)w=-P∆V

q=mC∆T use for calorimetry problems (water’s specific heat is 4.184 J/g °C)

∆H, S, G = (sum of products) – (sum of reactants)

∆H bond energy = (sum of reactants) – (sum of products); (bond energy is the backwards equation)

Hess’s law match the equation and add up ∆Hrxn (changes affect equation and ∆H)

∆G = H-TS (be careful with units on S they are in joules and H is in KJ, must make them the same.) (temperature affects –TS. If S is negative then TS is positive and as temperature increases G becomes more positive or less negative. If S is positive then TS is negative and as

temperature increases G becomes more negative or less positive.)

Entropy

Symbol______
Units______
Negative sign______
Positive sign______
Calculations (Two Ways)

Free Energy

Symbol______
Units______
Negative sign______
Positive sign______
Calculations (Four Ways)

∆G=-RTlnK Use to find the Equilibrium constant. (units on R are in joules. G in thermo is in KJ usually)

∆G=-nFE connection to Electrochem

Relationship between G, K, and E

If G=0 K = 1 E=0 at equilibrium

If G<0 K>1 E>0 products are favored

If G>0 K<1 E<0 reactants are favored

Any transition points such as melting point is considered @ equilibrium

Le Chatlier’s

∆H can be inserted into reactions: endo on reactants and exo in products if temp increases it shifts to the side without a ∆H, a decrease shifts to the side with a ∆H.

Pressure causes a reaction to shift to the side with the fewest gases

Multiple Choice Questions

question 1

Which of the following is a graph that describes the pathway of reaction that is exothermic and has high activation energy?

A. B.

question 2

When solid NH4SCN is mixed with solid Ba(OH)2 in a closed container, the temperature drops and a gas is produced. Which of the following indicates the correct signs for ΔG, ΔH, and ΔS for the process?

ΔG ΔHΔS

A) – – –

B) – + –

C) – + +

D) + – +

question 3

N2(g) + 3 H2(g) → 2 NH3(g)

The reaction indicated above is thermodynamically spontaneous at 298 K, but becomes nonspontaneous at higher temperatures. Which of the following is true at 298 K?

A) ΔG, ΔH, and ΔS are all positive.

B) ΔG, ΔH, and ΔS are all negative.

C) ΔG and ΔH are negative, but ΔS is positive.

D) ΔG and ΔS are negative, but ΔH is positive.

E) ΔG and ΔH are positive, but ΔS is negative.

question 4

Of the following reactions, which involves the largest decrease in entropy?

A) 2 CO(g) + O2(g) → 2 CO2(g)

B) Pb(NO3)3(s) + 2 KI(s) → PbI2(s) + 2 KNO3(s)

C) C3H8(g) + O2(g) → 3 CO2(g) + 4 H2O(g)

D) 4 La(s) + 3 O2(g) → 2 La2O3(s)

question 5

Assume the data graphed was collected at a constant pressure of 0.97 atm and represents four different temperature samples of pure neon gas. Which of the following temperatures most likely corresponds to the data graphed for sample “D”?

A) 273 KB) 298 KC) 305 KD) 338 K

question 6

At 298 K, as the salt MX dissolves spontaneously to form an aqueous solution, ∆S and ∆H are

positive. Which describes the value of ∆G and the absolute values of its components, T∆S and ∆H?

A) ∆G < 0; T∆S∆HB)∆G < 0; T∆S∆H

C) ∆G > 0; T∆S∆HD) ∆G > 0; T∆S∆H

Free Response # 1 (1998)

C6H5OH(s) + 7 O2(g) → 6 CO2(g) + 3 H2O(l)

When a 2.000-gram sample of pure phenol, C6H5OH(s), is completely burned according to the equation above, 64.98 kilojoules of heat is released. Use the information in the table below to answer the questions that follow.

Substance / Standard Heat of Formation, ΔH°f,
at 25°C (kJ/mol) / Absolute Entropy, S°,
at 25°C (J/mol-K)
C(graphite) / 0.00 / 5.69
CO2(g) / −393.5 / 213.6
H2(g) / 0.00 / 130.6
H2O(l) / −285.85 / 69.91
O2(g) / 0.00 / 205.0
C6H5OH(s) / ? / 144.0

(a)Calculate the molar heat of combustion of phenol in kilojoules per mole at 25°C.

(b)Calculate the standard heat of formation, ΔH°f, of phenol in kilojoules per mole at 25°C.

(3) Is reaction thermodynamically favorable? Justify answer

Free Response # 2 ( 2000)

O3(g) + NO(g) → O2(g) + NO2(g)

Consider the reaction represented above.

(a) Referring to the data in the table below, calculate the standard enthalpy change, ΔH, for the reaction atv25°C. Be sure to show your work.

(b)Make a qualitative prediction about the magnitude of the standard entropy change, ΔS°, for the reaction at25°C. Justify your answer.

(c) On the basis of your answers to parts (a) and (b), predict the sign of the standard free-energy change, ΔG°, for the reaction at 25°C. Explain your reasoning.