NOTES: Thermochemistry Part 1 - Heat

Honors text: Chapter 16Unit 09

NOTES: Thermochemistry Part 1 - Heat

HEAT-

TEMPERATURE -

Thermochemistry: the study of energy (in the form of heat) changes that accompany physical & chemical changes

heat flows from high to low (hot cool)
endothermic reactions: absorb energy in the form of heat; show a positive value for quantity of heat (q > 0)
exothermic reactions: release energy in the form of heat; show a negative value for quantity of heat (q < 0)

Magnitude of Heat Flow:

Units of heat energy:

For a pure substance of mass m, the expression of q can be written as:

where, q =

m =

c =

t =

Specific heat = the amount heat that must be added to raise the temp. of 1 g of a substance by 1C, with no change in state.

Specific heat values (in J / g C):

Examples:

How much heat is given off by a 50.0 g sample of copper when it cools from 80.0 to 50.0C?

Iron has a specific heat of 0.446 J/gC. When a 7.55 g piece of iron absorbs 10.33 J of heat, what is the change in temperature? If it was originally at room temp. (22.0C), what is the final temperature?

The specific heat of copper is 0.382 J/gC. How much heat is absorbed by a copper plate with a mass of 135.5 g to raise its temperature from 25.0C to oven temperature (420F)?

Thermochemistry Part 2 - Calorimetry

Honors text: Chapter 16Unit 09

Q: If you leave your keys and your chemistry book sitting in the sun on a hot summer day, which one is hotter?

Q: Why is there a difference in temperature between the two objects?

Honors text: Chapter 16Unit 09

Heat required to melt ice (a.k.a. latent heat of fusion) cannot be measured directly, but calorimetryprovides an experimental method allowing this heat transfer to be measured indirectly.

Calorimetry: measurement of the amount of heat evolved or absorbed in a chemical reaction, change of state or formation of a solution.

The enthalpy change associated with a chemical reaction or process can be determined experimentally.

Measure the ______during a reaction at CONSTANT pressure.
A ______ is a device used to measure the heat absorbed or released during a chemical or physical process

Simple/Coffee Cup Calorimeter – Picture of coffee cup calorimeter:

The cup is filled with water,which

absorbs the heat evolved(or given off)

by the reaction.
qrxn = -qcal

What happens in a calorimeter?

One object will ______, and the other will ______the heat
System loses heat to surroundings =
System absorbs heat from surroundings =

When a hot chunk of metal is dropped in a cool glass of water, the metal cools off.

Where did the heat from the metal go?

Did the metal lose more heat than the water gained?
Magnitude of ______= ______(ALWAYS!)

To do calorimetry problems…

First, make a chart -

Measurement / Water (cal) / Pebble (rxn)
Heat (q)
Mass (m)
Specific Heat (c) / 4.184
Final Temp (Tf)
Initial Temp (Ti)

Example1 : A small pebble is heated and placed in a foam cup calorimeter containing 25.0 g of water at 25.0 C. The water reaches a maximum temperature of 26.4 C. How many joules of heat were released by the pebble? The specific heat of water is 4.184 J/g C.

Example 2: When 1.00 g of ammonium nitrate, NH4NO3, is added to 50.0 g of water in a coffee cup calorimeter, it dissolves, NH4NO3 (s)  NH4+(aq) + NO3-(aq), and the temperature of the water drops from 25.00C to 23.32C. Calculate q for the reaction system.

Example 3: Suppose that 100.00 g of water at 22.4 °C is placed in a calorimeter. A 75.25 g sample of Al is removed from boiling water at a temperature of 99.3 °C and quickly placed in a calorimeter. The substances reach a final temperature of 32.9 °C . Determine the SPECIFIC HEAT of the metal.The specific heat of water is 4.184 J/g C.

Bomb Calorimeter Picture of bomb calorimeter:

NOTE: In a bomb calorimeter, heat is transferred from the

sample to the oxygen-enriched chamber, to the metal that

makes up the chamber, to the water… thus we cannot just

use the specific heat of water; instead heat capacity of the

calorimeter, Ccal, can be used or calculated.

It is possible to calculate the amount of heat absorbed or

evolved by the reaction if you knowthe heat capacity, Ccal,

and the temp. change, Δt, of the calorimeter.

Everything else is the same (remember, the heat lost from

the reactiongoes into the calorimeter)

EXAMPLE 4: The reaction between hydrogen and chlorine, H2 + Cl22HCl, can be studied in a bomb calorimeter. It is found that when a 1.00 g sample of H2 reacts completely, the temp. rises from 20.00C to 29.82C. Taking the heat capacity of the calorimeter to be 9.33 kJ/C, calculate the amount of heat evolved in the reaction.

EXAMPLE 5: When 1.00 mol of caffeine (C8H10N4O2) is burned in air, 4.96 x 103 kJ of heat is evolved. Five grams of caffeine is burned in a bomb calorimeter. The temperature is observed to increase by 11.37C. What is the heat capacity of the calorimeter in J/C?

EXAMPLE 6: When twenty milliliters of ethyl ether, C4H10O. (d=0.714 g/mL) is burned in a bomb calorimeter, the temperature rises from 24.7C to 88.9C. The calorimeter heat capacity is 10.34 kJ/C.

(a)What is q for the calorimeter?

(b) What is q when 20.0 mL of ether is burned?

Thermochemistry Part 3–Enthalpy and Thermochemical Equations

Enthalpy

Enthalpy = a type of chemical energy (thermodynamic potential), sometimes referred to as “heat content”
Enthalpies of Reaction (∆Hrxn): the enthalpy change (reported in kJ/mol) that accompanies a chemical reaction is called the enthalpy of reaction. Also called “heat of reaction”

If ∆Hrxn = negative

Exothermic
Heat is evolved, or given off
Under conditions of constant pressure, q =ΔH < 0 (negative sign)

If ∆Hrxn = positive

Endothermic
Heat isabsorbed
Under conditions of constant pressure, q = ΔH > 0 (positive sign)

Thermochemical Equations

Thermochemical equations are balanced chemical equations that show the associated enthalpy change (H)

balanced equation
enthalpy change (ΔHrxn)

Rules of Themochemistry:

Rule #1) The magnitude of H is directly proportional to the amount of reactant consumedand product produced.

Example 1: H2 + Cl22HclH = - 185 kJ

Calculate H when 1.00 g of Cl2 reacts.

Example 2: When an ice cube weighing 24.6 g of ice melts, it absorbs 8.19 kJ of heat. Calculate H when 1.00 mol of solid water melts.

Example 3: Methanol burns to produce carbon dioxide and water:

2CH3OH + 3O2 2CO2 + 4H2O + 1454 kJ

What mass of methanol is needed to produce 1820 kJ?

Example 4: How much heat is produced when 58.0 liters of hydrogen (at STP) are also produced?

Zn + 2HCl  ZnCl2 + H2 + 1250 kJ

Rule #2) H for a reaction is equal in the magnitude but opposite in sign to H for the reverse reaction. (If 6.00 kJ of heat absorbed when a mole of ice melts, then 6.00 kJ of heat is given off when 1.00 mol of liquid water freezes)

Example: Given: H2 + ½ O2 H2OH = -285.8 kJ

Calculate H for the equation:

2H2O 2H2 + O2

Rule #3) The value of H for a reaction is the same whether it occurs in one step or in a series of steps.

for the overall equation is the sum of the H’s for the individual equations:

Hess’s Law:H = H1 + H2 + …

Example 1: Calculate H for the reaction:C + ½ O2 CO

Given:

C + O2 CO2H = -393.5 kJ

2CO + O2 2CO2H = -566.0 kJ

Example 2: Find the heat of reaction (enthalpy) for the following reaction: NO + ½ O2 NO2 H = ?

Given the following equations….

½ N2 + ½ O2 NOH = +90.4 kJ

½ N2 + O2NO2H = +33.6

Thermochemistry Part 4–Phase Changes and Heats of Formation

specific heat = the amt of heat that must be added to a stated mass of a substance to raise its temp by 1C, with no change in state.

Ex: How much heat is released by 250.0 g of H2O as it cools from 85.0C to 40.0C? (specific heat of water = 4.18 J/gC)

Heat changes involving phase changes

LATENT HEAT OF FUSION, :the enthalpy change (energy absorbed) when a compound is converted from a solid to a liquid without a change in temperature. “Latent” means hidden; the heat absorbed/released during a phase change does not cause the temperature to change.

Note:for water is ______

LATENT HEAT OF VAPORIZATION, - the enthalpy change (energy absorbed) when a compound is converted from a liquid to a gas without a change in temperature.

Note:for water is ______

When substances change state, they often have different specific heats:

cice= 2.09 J/gCcwater= 4.18 J/gCcsteam= 2.03 J/gC

Example 1: How much heat is released by 250.0 g of H2O as it cools from 125.0C to -40.0C

Five steps…

Cool the steam m∙csteam∙T =
Condense m(-Hvap) =
Cool the liquid waterm∙cwater∙T =
Freezem(-Hfus) =
Cool the solid icem∙cice∙T =

Example 2: How much heat energy is required to bring 135.5 g of water at 55.0C to its boiling point (100C) and then vaporize it?

Example 3: How much heat energy is required to convert 15.0 g of ice at –12.5C to steam at 123.0C?

Enthalpies of Formation

Hf = enthalpy of formation

usually exothermic
see table for Hf values
enthalpy of formation of an element in its stable state =
these can be used to calculate H for a reaction
standard enthalpy change, H, for a given thermochemical equation is = to the sum of the standard enthalpies of formation of the product – the standard enthalpies of formation of the reactants.

elements in their standard states can be omitted:

2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s)

the coefficient of the products and reactants in the thermochemical equation must be taken into account:

2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)

Example: Calculate H for the combustion of one mole of propane:C3H8 (g) + 5O2 (g)  3CO2 (g) + 4H2O (l)

Example: The thermochemical equation for the combustion of benzene, C6H6, is:

C6H6(l)+ 15/2 O2(g)  6CO2 (g) + 3H2O(l)H = -3267.4 kJ

Example: When hydrochloric acid is added to a solution of sodium carbonate, carbon dioxide gas is formed. The equation for the reaction is:2H+ (aq) + CO32-(aq)  Co2 (aq) + H2O (l)

Calculate H for this thermochemical equation.

Thermochemistry Part 5–Spontaneity

THERMODYNAMICS = the study of energy changes that accompany physical and chemical changes.

Enthalpy (H):): the total energy “stored” within a substance

Enthalpy change (H):a comparison of the total enthalpies of the product & reactants.

Exothermic v. Endothermic:

Exothermic reactions/changes: release energy in the form of heat; have negative H values
H2O(g)  H2O(l)ΔH = -2870 kJ
Endothermic reactions/changes: absorb energy in the form of ehat; have positive H values.
H2O(l)  H2O(g)ΔH = +2870 kJ

Changes that involve a decrease in enthalpy are favored!

Reaction pathways:

• Entropy (S): the measure of the degree of disorder in a system; in nature, things tend to increase in entropy, or disorder.

All physical & chemical changes involve a change in entropy, or S. (remember that high entropy is favorable)
Enthalpy and entropy are DRIVING FORCES for spontaneous reactions (rxns that happen at normal conditions)
It is the interplay of these 2 driving forces that determines whether or not a physical or chemical change will actually happen.

• Free Energy (G): relates enthalpy and entropy in a way that indicates which predominates; the quantity of energy that is available or stored to do work or cause change.

where:G =

H =

T =

S =

G: pos value means change is NOT spon.

G: neg value means change IS spon.

Relating Enthalpy and Entropy to Spontaneity

Example of reaction / H / S / Spontaneity
H2O(g) H2O(l)
H2O(s) H2O(l)

Examples:

1) For the decomposition of O3 (g) to O2(g): 2O3(g)3O2(g)

H = -285.4 kJ/mol and S = 137.55 J/mol•K at 25°C.

a) Calculate G for the reaction.

b) Is the reaction spontaneous?

c) Is H or S (or both) favorable for the reaction?

2) What is the minimum temperature (in °C) necessary for the following reaction to occur spontaneously?

Fe2O3 (s) + 3CO(g)  2Fe(s) + 3CO2 (g)

H = +144.5 kJ/mol; S = +24.3 J/K•mol

(Hint: assume G = -0.1 kJ/mol)

Standard Heats of Formation, Gibbs Free Energy, and Entropy

Substance
(compounds at 1 atm, aqueous ions at 1M) /
at 25C
(kJ/mol) /
at 25C
(kJ/mol) /
at 25C
(J/mol·K)
Al3+(aq) / -531.0 / -485.0 / -321.7
Al2O3 (s) / -1676.0 / -1576.4 / 50.99
Br2 (g) / 30.91 / 3.14 / 245.3
Br2 (l) / 0.0 / 0.0 / 152.3
C (s, diamond) / 1.9 / 2.866 / 2.439
C (s, graphite) / 0.0 / 0.0 / 5.694
CH4 (g) / -74.86 / -50.79 / 186.2
C3H8(g) / -103.8 / -23.5 / 269.9
CO (g) / -110.5 / -137.2 / 197.9
CO2 (g) / -393.5 / -394.4 / 213.6
CO32-(aq) / -677.1 / -527.8 / -56.9
CaCO3 (s) / -1207.0 / -1127.7 / 88.7
CaO (s) / -635.1 / -604.2 / 39.75
Cl2 (g) / 0.0 / 0.0 / 223.0
Cu (s) / 0.0 / 0.0 / 33.2
Cu+(aq) / 71.7 / 50.0 / 40.6
Cu2+(aq) / 64.8 / 65.5 / -99.6
F2 (g) / 0.0 / 0.0 / 203
Fe (s) / 0.0 / 0.0 / 27.2
Fe2O3 (s) / -822.1 / -741.2 / 90.0
H+(aq) / 0.0 / 0.0 / 0.0
H2 (g) / 0.0 / 0.0 / 130.6
H2O (g) / -241.8 / -228.6 / 188.7
H2O (l) / -285.8 / -237.2 / 69.94
H2O2 (l) / -187.8 / -114.0 / 92
HCl (g) / -92.31 / -95.27 / 186.7
H2S (g) / -20.1 / -33.02 / 205.6
I2 (g) / 62.4 / 19.4 / 260.6
I2(s) / 0.0 / 0.0 / 117
Mg(OH)2(s) / -924.5 / -833.6 / 63.2
N2 (g) / 0.0 / 0.0 / 191.5
NH3 (g) / -46.19 / -16.64 / 192.5
NO (g) / 90.37 / 86.69 / 210.6
NO2 (g) / 33.85 / 51.84 / 240.5
Na2CO3 (s) / -1131.1 / -1048 / 136
NaCl (s) / -411.2 / -384.03 / 72.4
O2 (g) / 0.0 / 0.0 / 205.0
O3 (g) / 142.0 / 163.4 / 238
P (s, white) / 0.0 / 0.0 / 44.4
P (s, Red) / -18.4 / -14 / 29
S (s, rhombic) / 0.0 / 0.0 / 31.9
S (s, monoclinic) / 0.30 / 0.096 / 32.6
SO2 (g) / -296.8 / -300.4 / 248.5
SO3 (g) / -395.7 / -370.4 / 256.2

Ch 16 Problem Set 1: Calculating Heat

Although the calorimeters are used to measure heat, the fact is that heat cannot be measured directly. Instead, the calorimeter is used to measure the change in temperature of a measured amount of a given substance. Heat is then calculated from this data; heat is a derived quantity. Recall that the textbook defines specific heat capacity (c) as the quantity of heat required to raise the temperature of a gram of the substance 1C. The equation is as follows.

The equation used to calculate heat from experimental data can be obtained from the above equation simply by rearranging terms. Multiplying both sides of the equation by (mass  temperature change) gives:

To solve problems involving heat calculations, simply rearrange terms to isolate the unknown variable.

Example A

How much heat will be absorbed by 320 g of water when its temperature is raised by 35C? The specific heat of water is 4.18 J/(g C).

Solution

Example B

Calculate the specific heat for aluminum if 16,500 J of heat are absorbed in raising the temperature of 1.50  102 g of aluminum by 125C.

Solution

Now you try it!

How much heat will be given off by 55 g of water as it cools from 87C to 25C?

Calculate the specific heat of glass from the following data. The temperature of a piece of glass with a mass of 65 grams increases by 26C when it absorbs 840 J of heat energy.

Calculate the temperature change for mercury if 160 grams of the metal absorb 1500 J of heat energy. Mercury’s specific heat is 0.14 J/(g C).

The temperature of a 150 gram drinking glass decreased from 25C to 14C as 200 grams of water were added. Determine the original temperature of the water. (Hint: Heat loss must equal heat gain).

How much heat energy is required to raise the temperature of 200 grams of water from 25C to 100C?

The temperature of an iron bar with a mass of 87.0 g is raised from 31C to 543C. In the process 4.90  103cal of heat energy were absorbed. What is the specific heat of iron?

Ch 16 Problem Set 2: Calorimetry

Imagine that you’re working outdoors on a hot, humid day. If you drink four glasses of ice water at 0C, how much heat energy is transferred as this water is brought to body temperature? Assume that each glass contains 250.0 g of water and that your body temperature is 37C.

How much heat energy is released to your body when a cup of hot tea containing 200.0 g of water is cooled from 65C to body temperature, 37C?

How much heat energy is needed to raise the temperature of a 355 g aluminum baking sheet from room temperature, 25C, to a baking temperature of 200C? (the specific heat of aluminum is 0.897 J/gC)

A nutritional chemist burns one pulverized peanut with a mass of 0.887 g in a bomb calorimeter. The calorimeter contains 2.50 kg f water, and its temperature increases from 25.0C to 27.0C as the peanut burns. What is the energy content of the peanut? What is the energy content of peanuts in kJ/g?

What is the energy content of a 1.28 g sample of oatmeal that raises the temperature of 2.50 kg of water within a calorimeter from 25.0C to 27.2C? What is the energy content of oatmeal in kJ per gram?

Predict the final temperature of 2.50 kg of water within a calorimeter if the water is at 25.0C before a 1.8 g piece of dried peach with an energy content of 18.5 kJ is burned.

A swimming pool measures 6.0 m x 12.0 m and has a uniform depth of 3.0 m. The pool is full of water at a temperature of 20.0C. How much energy must be released by the pool’s heater to raise the water temperature to 25.0C? (the density of water is 1 g/mL).

Predict the final temperature of 3.50 kg of water in a calorimeter if the water is at 27.5C before 0.77 oz of noodles containing 81.0 kcal are burned.

On a cold winter day with a temperature of 4.0C, you pick up a penny from the ground and put it in your pocket. If the penny has a mass of 1.85 g, how much heat energy must be transferred to the coin to warm it to your body temperature, 37C? (assume the penny is pure copper, and copper’s specific heat is 0.385 J/gC)

Isooctane is a primary component of gasoline and gives gasoline its octane rating. Burning 1.00 mL of isooctane (d = =0.688 g/mL) releases 33.0 kJ of heat. When 10.00 mL of isooctane is burned in a bomb calorimeter, the temperature in the bomb rises from 23.2C to 66.5C. What is the heat capacity of the bomb calorimeter (in J/C)?

Urea, (NH2)2CO, is a commonly used fertilizer. When 237.1 mg of urea is burned, 2.495 kJ is given off. If 500.0 mg of urea is burned in a bomb calorimeter (heat capacity = 5326 J/C) initially at 23.15C, what is the calorimeter temperature when combustion is complete?

Ch 16 Problem Set 3:

Thermochemical Equations & Hess’s Law

The production of iron and carbon dioxide from iron (III) oxide and carbon monoxide is an exothermic reaction : Fe2O3 (s) + 2CO (g)  2Fe (s) + 3CO2(g) + 26.3 kJ

How much kilojoules of heat are produced when 3.40 grams of Fe2O3 reacts with an excess amount of CO?

The burning of magnesium in oxygen is a very exothermic reaction:2Mg + O2  2MgO + 1204 kJ

How many kilojoules are given off when 6.55 g of Mg reacts with an excess amount of oxygen?

A considerable amount of heat is required for the decomposition of aluminum oxide:

2Al2O3  4Al + 3O2H = +3352 kJ

How many grams of Al are produced when 5783 kJ of heat is absorbed by the reaction?

The combustion of ethane, C2H4, is an exothermic reaction:

C2H4 + O2  2CO2 2H2OH = -1390 kJ

Calculate the heat liberated when 4.79 g of ethene burns.

Calculate H for the formation of lead (IV) chloride by the reaction of lead (II) chloride with chlorine:

PbCl2 + Cl2  PbCl4H = ?

Use the following thermochemical equations:

Pb + 2Cl2  PbCl4H = -329.2 kJ

Pb + Cl2  PbCl2H = -359.4 kJ

From the following reactions…

½ N2 + ½ O2  NOH = +90.4 kJ

½ N2 + O2  NO2H = +33.6 kJ

Calculate the heat of reaction for : NO + ½ O2  NO2

Calculate the heat change for the formation of copper (I) oxide from the elements:

Cu + ½ O2  CuO

Use the following two thermochemical equations to make the calculations:

CuO + Cu  Cu2OH = -11.3 kJ

Cu2O + ½ O2  2CuOH = -114.6 kJ

Find the enthalpy of reaction, Hrxn, for the formation of phosphorus pentachloride from its elements.

2P + 5Cl2  2PCl5

Use the following equations:

PCl5  PCl3 + Cl2H = +87.9 kJ

2P + 3 Cl2  2PCl3H = -574 KJ

Calculate the enthalpy of reaction, Hrxn, for the formation of nitrogen monoxide from its elements:
N2 + O2  2NO

Use these equations:

4NH3 + 3O2  2N2 + 6H2OH = -1530 kJ

4NH3 + 5O2  4NO + 6H2OH = -1170 kJ

Calculate the enthalpy of reaction, Hrxn, for the formation of nitrogen monoxide from its elements:

H2S(g) + 4F2(g)  2HF(g) + SF6(g)

Use these equations:

½ H2(g) + ½ F2(g)  HF(g)H = -273 kJ

S(s) + 3F2(g) SF6(g)H = -1220 kJ

H2(g) + S(s) H2S(g) H = -21 kJ

Ch 16 Problem Set 4:

Heat transfer in Liquids and Solids

Equations and constants:

q = mct (for problems involving changes in temperature)

q = m·Hfusor q = m·Hvap (for problems involving phase changes)

Specific heat of ice = 2.09 J/g·ºC

Specific heat of water = 4.18 J/g·ºC

Specific heat of steam = 2.03 J/g·ºC

Heat if fusion of water = 334 J/g

How much heat must your body transfer to 500.0g of water to heat it from 25.0ºC to body temperature, 37.0ºC?

How much heat energy from the sun is needed to heat a 1575g puddle of water from 5.00ºC in the morning to 20.0ºC by the afternoon?

How much heat is needed to warm a 50.0g piece of solid copper from 25.0ºC to 200.0ºC? (the specific heat of copper is 0.385 J/g·ºC)

How much energy is needed to heat a 35.5g sample of ice at -17.5ºC to liquid water at 77.3ºC?

How much energy is needed to heat a 68.9g sample of water at 88.5ºC to steam at 103.7ºC?

How much heat energy is released when a 234.7g sample of steam at 114.5ºC is cooled until it’s ice at a temperature of -8.50ºC?

Prob Set 4 Cont’d:Heats of Formation

#7-10 Use the standard heats of formation table (Table C-13 of Appendix C, p. 921) to calculate the enthalpy change (H) for these reactions:

7)Br2(g)  Br2(l)

8)CaCO3(s)  CaO(s) + CO2(g)

9)2NO(g) + O2(g)  2NO2(g)

10)4NH3(g) + 5O2(g)  2NO(g) + 6H2O(l)

Ch 16 Problem Set 5:

Reaction Spontaneity (∆H, ∆S & ∆G)

Enthalpy change

Entropy change

Free energy change

Gibbs-Helmholtz equation

Is the entropy (degree of disorder) increasing or decreasing in these reactions?

H2(g) + Br2(l)  2HBr(g)

CuSO4∙5H2O(s)  CuSO4(s) + 5H2O(g)

2XeO3(s) 2Xe(g) + 3O2(g)

Classify each of these systems as always spontaneous (A), never spontaneous (N), or depends on the relative magnitude of the heat and entropy changes (D).

Heat is released; entropy decreases

Heat is absorbed; entropy decreases

Heat is absorbed; entropy increases

Heat is released; entropy increases

Calculate the standard entropy change associated with each reaction:

2H2O2(l)  2H2O(l) + O2(g)

I2(g) I2(s)

2CO(g) + O2(g)  2CO2(g)

A reaction is endothermic (positive ΔH) and has a positive entropy. Would this reaction more likely be spontaneous at high or low temperatures? Justify your answer.

A reaction has a ΔS of -122 J/K·mol and a ΔH of -78 kJ/mol at 285C.

Calculate ΔG for the above reaction.

Is this reaction spontaneous?

Calculate the standard free energy change for the reaction between iron (III) oxide and carbon (graphite).

2Fe2O3(s) + 3C(s)  4Fe(s) + 3CO2(g)

A student warned his friend not to swim in a river close to an electric plant. He claimed that the ozone produced by the plant turned the river water to hydrogen peroxide, which would bleach hair. The reaction is

O3(g) + H2O(l)  H2O2(aq) + O2(g)

Assuming that the river water is at 25C and all species are at standard concentrations, show by calculation whether his claim is plausible. Take at 25C to be +163.2kJ/mol and = -134 kJ/mol.