CHM 122 Chapter 8 -Thermochemistry: Chemical Energy

CHM 122 Chapter 8 -Thermochemistry: Chemical Energy

8.1 Energy and Its Conservation

A. Definitions

Thermochemistry is a part of Thermodynamics dealing with energy changes associated with physical and chemical reactions

Why do we care?

- Will a reaction proceed spontaneously?

- If so, to what extent?

However, it won’t tell us:

- How fast the reaction will occur

- The mechanism by which the reaction will occur

Energy is the capacity to do work or to transfer heat

For example if you climb a mountain, you do some work against the force of gravity as you carry yourself and your equipment up the mountain. You can do this because you have the energy, or capacity to do so, the energy being supplied by the food that you have eaten. Food energy is chemical energy –energy stored in chemical compounds and released when the compounds undergo the chemical process of metabolism

- Kinetic Energy: energy associated with mass in motion

- Potential Energy: energy associated with the position of an object relative to other objects (energy that is stored - can be converted to kinetic energy)

System: portion of the universe under study

Surroundings: everything else

Open System: can exchange energy and matter through boundary

Closed System: can exchange energy through boundary

Isolated System: can exchange neither with surroundings

We can define the system and surroundings however we want!

Example: You place 10.0g of LiOH in a Styrofoam cup of water. What is the system, and what are the surroundings?

Energy and units

One of the earliest units of energy to be devised was the unit of calorie, or cal. Since the calorie represents a small amount of heat energy and we usually work with larger quantities of matter the unit of kilocalorie or kcal is used instead.

The SI unit is Joules 1 cal = 4.184 J (exactly)

8.2 Internal Energy and State Functions

Energy (E) can be transferred two different ways:

1. By doing work (w) (applying a force over a distance)

w = F x d

Work: Can be electrical, mechanical, etc.

2. Transferring heat (q) (results in a change in temperature)

Heat: The amount of thermal energy transferred from one object to another as the result of a temperature difference between the two

Thermal Energy: The kinetic energy of molecular motion and is measured by finding the temperature of an object

We will symbolize heat energy transferred by the letter q.

HOT OBJECT COLD OBJECT

Heat transfer occurs by way of collisions between randomly moving particles of matter. The particles with higher thermal energy are moving more quickly, when they collide with slower moving particles, some of their energy is transferred to the slower particle as heatenergy increasing its speed

Note: w, q, and E all have the same units (Joule), but:

- w & q depend on path (path function)

- E is independent of path (state function)

• A function or property whose value depends only on the present state, or condition, of the system, not on the path used to arrive at that state

First Law of Thermodynamics

“The total energy of the universe is constant.”

“Energy is neither created nor destroyed in a process, only converted to another form.” -Conservation of Energy

“You can’t win . . . you can only break even.”

The sum of all the kinetic and potential energies of particles making up a substance is called internal energy, E

8.3 Expansion Work

8.4 Energy and Enthalpy

When there is no work done by the system, theheat given off or absorbed by the reaction would be equal to the change in the internal energy of the system.

• Esys= q (if and only if w = 0)
• Esys= qv (at constant volume)

However, if a gas is driven out of the flask during the reaction, the system does work on its surroundings. If the reaction pulls a gas into the flask, the surroundings do work on the system. The amount of heat given off or absorbed will no longer equal to the change of the internal energy of the system because some of the heats has been converted to work.

Esys= q + w

qP = ΔE + PΔV (at constant pressure)

The amount of heat absorbed or released during a chemical reaction by a system is also call Enthalpy, ΔH. Under constant pressure, qp = ΔH

qP = ΔE + PΔV = H

Enthalpy is a state function whose value depends only on the current state of the system, not on the path taken to arrive at that state. Enthalpy is an extensive property because the magnitude of ΔH depends on the quantity of reactant

ΔH = Hfinal - Hinitial = Hproducts - Hreactants

For convenience, the sign and value of ΔH is always provided in a thermochemical equation.

2 Al(s) + Fe2O3(s)  2Fe(s) + Al2O3(s) ΔH = -825 kJ

Write the above thermochemical equation in the reverse reaction

8.6The Thermodynamic Standard State - Enthalpies of Physical and Chemical Change

Thermodynamic Standard State: Most stable form of a substance at 1 atm pressure and at a specified temperature, usually 25 °C; 1 M concentration for all substances in solution.

Standard enthalpy of reaction is indicated by the symbol ΔHo

Example: Applying Stoichiometry to Heats of Reaction

A propellant for rockets is obtained by mixing the liquids hydrazine, N2H4, and dinitrogen tetroxide, N2O4. These compounds react to give gaseous nitrogen, N2 and water vapor, evolving 1049 kJ of heat at constant pressure when 1 mol N2O4 reacts.

Write the thermochemical equation for this reaction

Write the thermochemical equation for the reverse of the reaction

How much heat evolves when 10.0 g of hydrazine reacts according to the reaction described in (a) ?

Example: Calculate the ΔH for the reaction that occurs when 4.1 g CH4decomposes according to the following reaction:

CH4(g) + 2O2(g) -> CO2(g) + 2H2O(g) ΔH = - 802 kJ/mol

8.6Calorimetry and Heat Capacity

In general, when heat transfers to an object, the temperature increased. When heat lost, the temperature decreased. This relationship can be describe in this equation

q = C x ΔT at constant volume

q = SH x m x ΔT at constant pressure

The heat evolved in a chemical reaction can be determined by a process called calorimetry. Calorimeter is a device that use to measure the heat flow. It is well-insulated so that, ideally, no heat enters or leaves the calorimeter from the surroundings

At constant pressureat constant volume

Example

Iron metal has a specific heat of 0.449 J/goC. How much heat (in kJ) is transferred to a 5.00g piece of iron, initially at 20.0oC, when it is placed in a pot of boiling water? Assume that the temperature of the water is 100.0oC and that the water remains at this temperature, which is the final temperature of iron

ExampleYou dissolve 5.25 g of NaOH into 100.0 mL of water in a Styrofoam cup. The temperature of the water increases from 23.3 oC to 40.5 oC. What is the enthalpy of reaction in kJ/mol

Example: A 400. 0 g piece of iron (C (iron) = 0.38 J/g°C) is heated in a flame and dropped into a beaker containing 1000. 0 g of water at 20.0°C. The final temperature is 32.8°C.What was the initial temperature of the iron bar?

8.9 Hess’s Law

Hess’s Law

Energy conservation is the basis of Hess’s Law which states that, if a reaction is the sum of two or more other reactions, the ΔH overall is the sum of theΔH values of the constituent reactions

ΔH°total = ΔH°1 + ΔH°2 + etc….

For a reaction of 3 H2(g) + N2(g)  2NH3(g)

Example: Find ΔHorxn for the following reaction:

3H2(g) + O3(g)  3 H2O(g) ΔHorxn = ??

Use the following reactions with known ΔH’s

2H2(g) + O2(g)  H2O(g)ΔHo = -483.6 kJ

3O2(g)  2 O3 (g)ΔHo =. +285.4 kJ

Example

Find ΔHorxn for the following reaction

C(s) + H2O(g)  CO(g) + H2(g) Horxn = ?

Use the following reactions with known H’s

C(s) + O2(g)  CO2(g) ΔHo = -393.5 kJ

2CO(g) + O2(g)  2CO2(g) ΔHo = -566.0kJ

2H2(g) + O2(g)  2H2O(g) ΔHo = -483.6 kJ

Standard Heats of Formation

Theenthalpy change which occurs for a reaction which forms one mole of a compound from its constituent elements in their standard states at standard conditions is called the

STANDARD MOLAR ENTHALPY OF FORMATION or Enthalpy of Formation for short and is signified byΔH°f .

Ag(s) + ½ Cl2(g)  AgCl(s) ΔHrxn = -127.0 kJ

Then ΔH°f (AgCl(s)) = -1270 kJ/mol

Ho = Hof (Products) - Hof (Reactants)

Example: Using standard heats of formation, calculate the standard enthalpy of reaction for the photosynthesis of glucose (C6H12O6) and O2 from CO2 and liquid H2O.

6CO2(s) + 6H2O(l)  C6H12O6(s) + 6O2(g)

Calculate the enthalpy change for the following reaction

8Al(s) + 3 Fe3O4(s) 4Al2O3(s) + 9 Fe(s)

Given ΔH°f (Fe3O4) = -1120.9 kJ/mol

ΔH°f (Al2O3) = -1669.8 kJ/mold

An Introduction to Entropy

Entropy (S): The amount of molecular randomness in a system

Spontaneous Process: A process that, once started, proceeds on its own without a continuous external influence

Spontaneous processes are

•favored by a decrease in H (negative ΔH).

•favored by an increase in S (positive ΔS).

Nonspontaneous processes are

•favored by an increase in H (positive ΔH).

•favored by a decrease in S (negative ΔS).

Example

Predict whether ΔSo is likely to be positive or negative for each of the following reactions

a.H2C=CH2(g) + Br2(g)  BrCH2CH2Br(l)

b.Consider the following figures

Example

Is the Haber process for the industrial synthesis of ammonia spontaneous or nonspontaneous under standard conditions at 25.0oC.

3 H2(g) + N2(g)  2NH3(g)

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