Chapter 6: Thermochemistry - Moorpark...

C h e m i s t r y 1 A : C h a p t e r 6 P a g e | 1

Chapter 6: Thermochemistry

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Chemical Thermodynamics:

Chemical thermodynamics is the study of the changes and the transfers of energy

that accompany chemical and physical processes.

It is VERY IMPORTANT to always identify the phases in chemical reactions. This unit

also will occasionally require balancing reactions with fractions.

Thermodynamics addresses 3 fundamental questions:

1) Will two or more substances spontaneously react when mixed under specified

conditions or will it be nonspontaneous (will not react)?

2) What energy changes and transfers are associated with the reaction?

3) To what extent does a reaction occur (equilibrium)? (learn in Chem 1B)

Systems tend toward…

1) Minimizing potential energy

2) Maximizing entropy (degrees of freedom or ways of being).

Thermochemistry tells us nothing of the reaction rate. Reactions may be

spontaneously favored under specified conditions, but may not react at an

observable rate, or reactions may need a large initial amount of energy to initiate

and reach the activation energy.

Thermodynamic examples:

Methane will spontaneously react and combust in oxygen, but the reaction does

not initiate without a spark.

Ice will spontaneously melt into water, but only if the temperature is above

freezing.

We will study physical and chemical processes accompanied by the transfer of heat or

work. Nuclear energy, where the atom changes, is different.


Example 1:

Systems tend toward minimizing potential energy and maximizing entropy.

Identify the side that minimizes the potential energy. Identify the side that maximizes

ways of being (entropy). Assume the same temperature and atmospheric pressure for

each.

a) Top of waterfall or Bottom of waterfall

b) 1 mole H2 molecules or 2 mole of H atoms

c) H2O (g) at 100°C or H2O (l) at 100°C

Nature of Energy:

Energy is the capacity to do work or transfer heat. Energy (a quantity) can exist in

a variety of forms and can be transformed from one form to another. It is measured

by the amount of work done, usually in joules or watts. Chemistry generally uses

joules. SI units of energy: 1 Joule = 1 kg m2/s2

Energy Conversions

1 calorie (cal) = 4.184 joules (J) (exact)

1 Calorie (Cal, food) = 1000 cal = 1 kcal = 4184 J

1 kilowatt-hour (kWh) = 3.60 x 106 J

Some Forms of Energy:

• Electrical

kinetic energy associated with the flow of electrical charge

• Mechanical

kinetic energy associated with moving parts

• Heat or thermal energy

kinetic energy associated with molecular motion

• Light or radiant energy

kinetic energy associated with energy transitions in an atom

• Nuclear

potential energy in the nucleus of atoms

• Chemical

potential energy due to the structure of the atoms, the attachment between

atoms, the atoms’ positions relative to each other in the molecule, or the

molecules, relative positions in the structure


Energy can be classified into two major categories, potential and kinetic.

Potential Energy is the energy by virtue of its position of composition.

a) For an object that has height, potential energy is the mass x gravity (9.8 m/s2)

x height; Ep = mgh.

b) Another form is electrostatic potential energy from interactions with charged

particles; Eel = kQ1Q2/d, where k = 8.99 x 109 J m/C2 and C is coulomb that

is a unit of charge, Q is charge, and d is distance. Like charges repel and

opposites attract.

c) Potential energy also comes from the arrangement of chemical compositions.

When bonds break and new ones form in a chemical reaction, the energy

change generally is due to the changes in potential energy/composition.

Kinetic Energy is the energy of motion. For a moving object kinetic energy is half

the mass x velocity squared; Ek = ½ mv2

Example 2: What is the potential energy of a 400 g ball on top of a building that

is 30.0 m tall?

Example 3: What is the kinetic energy of a 400 g ball moving at 30.0 m/s?

Some Thermochemistry terms…

System: the portion we study

Surroundings: everything else affected by the portion we study

Universe: the combined system plus surroundings. In the lab experiments our

universe will be a calorimeter and its contents.


Laws of Thermodynamics:

Laws of thermodynamics are useful in predicting outcomes, but unlike a theory do

not explain the expected behavior.

Zeroth Law: Temperature Concept: Temperature measures the intensity of

hotness or coldness of an object. When two objects are brought

together heat always flows spontaneously from a hotter object to a

colder object until thermal equilibrium is reached. Think of the heat

flow direction when placing your hand in ice water verses hot water.

First Law: Law of Conservation of Energy: Energy cannot be created or

destroyed. The energy of the universe is constant.

This law works for ordinary chemical and physical processes. It was

later combined with another (conservation of matter) to include

nuclear reactions after E =mc2. It was changed to the Law of

Conservation of Matter and Energy

Example 4: What is the difference between temperature and heat?

Example 5: Review your temperature conversions between °F, °C, and Kelvin.

a) 98.6 °F = ? °C = ? K

b) 788 K = ? °C = ? °F

When you see X where X can be anything remember that it is just the change from the

final minus the initial values. X = Xfinal – Xinitial

Definitions:

Heat is the flow of energy caused by temperature differences

Enthalpy,H; the energy accounting for heat flow in constant pressure processes.

Units are generally kJ/mol. When the energy is measured in J or kJ for a particular

system, qp may be used.


Internal Energy, E or U; the sum of all the kinetic and potential energies of all

components. This includes all motions of vibration, rotation, movement through

space of the object, its atoms, nuclei, and even including subatomic particles like

electrons. We cannot truly measure all of this, so we generally try to solve for the

difference or change in internal energy in our studies. E = Efinal – Einitial

When solving for E we work with a constant volume system. Units are generally

kJ/mol. When the energy is measured in J or kJ for a particular system, qv may be

used.

Work, w; Work is force acting over a distance, (work = force x distance), the

energy used to cause an object to move. Pressure-Volume work involves the

compression or expansion of gases.

Work = -PV ; if using atm and liter: convert to J with R/R (101.3 J/L atm)

Work = -nRT; use the R with joules: R = 8.314 J/mol K

For most reactions the energy from work is quite small.

Expansion work is (-); work done by the system on the surroundings.

Compression work is (+); work done on the system by the surroundings

Relating together; qv = qp + w , or it can be written E = H + w

Endothermic: a process that absorbs heat (+)

Exothermic: a process that releases heat (-)

State Function: a property that

depends only on the present

state and not the path the system

took to reach the state. In

Thermochemistry state

functions are generally

capitalized. (Notice work is a

small w and is not a state

function)

Examples of state functions

include: E, H, V, P, T.


Example 6:

a) As 1.00 mole of H2O gas becomes H2O liquid at 100°C, 40.7 kJ of energy

are transferred at 1.00 atm pressure. What is the enthalpy sign for this

conversion? Is this process endothermic or exothermic?

b) During this process at 1.00 atm pressure, the volume decreases from 30.6

liters down to 18 ml. Solve for the work in L atm and in J. Is the sign of

work positive or negative? Is work done on the system by the

surroundings (compression) or by the system on the surroundings

(expansion) ?

Example 7:

Predict the sign for work, determine if the system is expanding or compressing and

state if work is done on or by the system.

a) 2 NO (g) + O2 (g) 2 NO2 (g)

b) 2 C4H10 (g) + 13 O2 (g) 10 H2O (l) + 8 CO2 (g)

Example 8:

Which of the following are state functions and which are not state functions ?

a) The temperature of an ice cube.

b) The enthalpy of fusion for water.

c) The time to travel 1-mile


Enthalpies of Reaction: Thermochemical Equations with Enthalpy:

1) Always include phases (s, l, g, aq) of substances. Energy varies with phase.

2) Assume temperature is constant for reactants at the beginning and products at

the end, even though the temperature may change during the reaction. H is

the energy required to return a system to the starting temperature at the

completion of the reaction.

3) Reactions may be balanced using fractions and not whole numbers.

4) Enthalpy (H) is an extensive property.

When a reaction is doubled, the enthalpy doubles.

5) When a reaction is reversed the enthalpy changes its sign.

6) Change in enthalpy may be experimentally determined using calorimeters.

7) The ° indicates standard conditions (1 atm and 25°C).

8) Scientists arbitrarily agreed to set a zero Hf° defined as the enthalpy of

formation of elements in their most common form under standard state

conditions. All other Hf° values are calculated from these

9) Enthalpy of formation Hf°) is the formation of 1 mole of a compound from

its elements in their most common form under standard conditions.

10) Tabulated enthalpy (Hf°) values are found in reference sources.

Example 9:

Given: 4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g); H = -906 kJ/mol

a) Is the reaction above endothermic or exothermic?

b) How many grams of NH3 gas are required to produce 374 kJ of energy?

c) What is the enthalpy term for… 2 NO (g) + 3 H2O (g)2 NH3 (g) + 5/2 O2 (g)


Calorimetry:

Using the First Law of Thermodynamics and a calorimeter it is possible to solve for

quantitative energy transfers.

Heat capacity (C) is the energy required to raise an object by 1°C or 1K (kJ/°C)

Specific heat capacity (C or s) is energy to raise one gram of a specific substance by

1°C or 1K (units: J/g°C)

Molar heat capacity (J/mol °C)

Qgain = - Qlost

Qcalorimeter = - Qreaction

Coffee- cup calorimetry: pressure is constant; Qp = mcT for calorimeter

Bomb calorimetry: volume is constant; Qv = CT for calorimeter (Chem1B)

It is important to remember that the system you are studying is going to have equal

energy, but opposite sign when compared to what happens in the calorimeter.

WATCH THE SIGNS!

Example 10:

A 7.82 g sample at 99.2C is added to 30.0g of water at 22.10C in a coffee-cup

calorimeter. The final temperature is 24.20C. The specific heat capacity of water is

4.184 J/gC.

a) Solve for the specific heat of the metal

b) Identify the metal. (choices: Al = 0.903 J/gC,

Fe = 0.449 J/gC, Pb = 0.128 J/gC)


Example 11:

Test the Law of Dulong and Petit which states that the molar heat capacity of

elements are all approximately 25 J/molC by solving for the molar heat capacities

of Al, Fe, Pb. Al = 0.903 J/gC, Fe = 0.449 J/gC, Pb = 0.128 J/gC)

Example 12:

Before the days of refrigerators and icemakers, drinking water in hot countries was

often cooled by storage in clay pots. Evaporation of water through the pot cooled the

remaining water. How many grams of water must evaporate to cool 946 grams (1

quart) of water from 33.5°C to 19.2°C? For water, specific heat is 4.184 J/g°C; Hvap

is 44.0 kJ/mol.

Enthalpies of Formation:

Enthalpy (or Heat) of Formation,H°f: The enthalpy change associated with the

formation of one mole of a substance from its constituent elements in their most

common form and state. H°f, the f represents formation and ° indicates the

standard state conditions of 25°C and 1 atm pressure.

Example 14:

Fix the following to make it an enthalpy of formation equation…

2K (l) + H2 (g) + C (diamond) + O3 (g) KHCO3 (s)


Example 15:

Write balanced equations that describe the formation of the following compounds

from their elements in their standard states. Use the appendix to obtain the values

of their standard enthalpy of formation:

a) NH4NO3 (s)

b) Ca3(PO4)2 (s)

c) C3H8 (g)

Hess’s Law:

Hess’s Law: If a reaction is carried out in a series of steps, H for the overall

reaction will equal the sum of the enthalpy changes in the individual steps.

It is possible to calculate enthalpy of a reaction in both the lab setting, and by using

reference sources. Enthalpy is an extensive property, so the amount is important. If

a series of reactions are used to create an overall reaction, the enthalpy values can

be added to solve for the overall enthalpy of a reaction.

Hrxn = Ha + Hb + Hc

Any reaction can be broken down to formation reactions from elements. When this

is done we have… H°rxn = nH°f (products) - nH°f (reactants)


Example 16: Use Hess’ Law to solve for the following…

a) Use reference values of H°f to solve for the H°rxn below…

C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l)

b) Calculate Hrxn for… 3 H2 (g) + O3 (g) 3 H2O (g) GOAL

Given: 3 O2 (g) 2 O3 (g) H = +284.6 kJ/mol

Hf˚ for H2O(g) Hf = -241.8 kJ/mol

c) Calculate the standard enthalpy of formation (H°f) of diborane gas, B2H6,

using the following…

Given: 4 B (s) + 3 O2 (g) 2 B2O3 (s) H = -2509.1 kJ/mol

2 H2 (g)+ O2 (g) 2 H2O (l) H = -571.7 kJ/mol

B2H6 (g)+3 O2 (g) B2O3 (s)+3 H2O (l) H = -2147.5 kJ/mol

Foods and Fuels:

Foods:

Most of the energy we need for life comes from carbohydrates and fats in food.

Carbohydrates (starches) decompose in our intestines into glucose (C6H12O6, blood

sugar). Sugars and carbohydrates have an average fuel value of 17 kJ/g (4 kcal/g)

C6H12O6 (s) + 6 O2 (g) 6 CO2 (g) + 6 H2O (l) H° = -2803kJ

Fats also react with oxygen to give energy. This energy can be stored for later use.

Fats average fuel value is 38 kJ/g (9 kcal/g).


Fuels:

Coal, petroleum and natural gas (fossil fuels) are the primary sources of energy. In

units of energy per gram, the fuel value generally increases for hydrocarbons as the

quantity of H is added.

In the United States, each person uses over 105 kWh of energy per year

C(s) + O2(g) → CO2(g) H°rxn = −393.5 kJ

CH4(g) +2 O2(g) → CO2(g) + 2 H2O(g) H°rxn = −802.3 kJ

C8H18(g)+12.5 O2(g) → 8 CO2(g)+9 H2O(g) H°rxn = −5074.1 kJ

Convert these values to kJ/gram of fuel…

C, CH4, C8H18

• Most of the energy for US

citizens comes from combustion

of fossil fuels, combustible

materials that originate from

ancient life.

• Fossil fuels cannot be

replenished

1 Quad = quadrillion British

thermal units = 1.06 x 1018

J.

The US consumes close to 100

quads of energy annually.

Example 17:

Citizens of the world burn the fossil fuel equivalent of 7 x 1012 kg of petroleum each year.

a) If all 7 x 1012 kg was in the form of octane, calculate the kg of CO2 produced

annually. C8H18 (l) + O2 (g) CO2 (g) + H2O (l) (unbalanced)

b) If the atmosphere currently contains 3 x 1015 kg of CO2, how long will it take

to double the atmospheric carbon dioxide, assuming no photosynthesis or

other reactions remove it?


Global Warming:

• CO2 is a greenhouse gas which allows light from the sun to reach the earth, but does

not allow the heat (infrared light) reflected off the Earth to escape back into outer

space. It acts like a blanket

• CO2 levels in the atmosphere are steadily increasing

• Observations suggest that

the average global air

temperature has risen 0.6

°C in the past 100 yrs.

• Atmospheric models

suggest that the warming

effect could worsen if

CO2 levels are not curbed

• Some models predict that

the result will be more

severe storms, more

floods and droughts, shifts

in agricultural zones,

rising sea levels, and

changes in habitats

Renewable Energy:

• Our greatest unlimited supply of energy is the sun. Technologies are being

developed to capture the energy of sunlight: parabolic troughs, solar power towers,

and dish engines concentrate the sun’s light to generate electricity; solar energy

may also be used to decompose water into H2(g) and O2(g); the H2 can then be used

by fuel cells to generate electricity.

H2(g) + ½ O2(g) → H2O(l) H°rxn = −285.8 kJ

• Hydroelectric power

• Wind power

Chapter 6: Thermochemistry - Moorpark...

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