Ch. 9: Thermochemistry Overview and Key Concepts

Thermodynamics

The study of energy and its transformation.

Thermochemistry

The study of energy changes associated with a chemical reaction.

Energy

The capacity to do work.

Work

A force acting over a distance.

Kinetic Energy

Energy due to the motion of an object.

Kinetic Energy Formula

Kinetic Energy = ½ mv² (m = mass, v = velocity).

Potential Energy

Energy due to the position of an object.

SI Unit of Energy

Joule.

Calorie to Joule Conversion

1 calorie (cal) = 4.184 J.

L·atm to Joule Conversion

101.32 J = 1 L·atm.

Kilowatt Hour to Joule Conversion

1 kWh = 3.60 x 10⁶ J.

System

A limited and well defined part that is under study.

Surroundings

Everything else other than the system.

Zeroth Law of Thermodynamics

If two systems are at the same time in thermal equilibrium with a third system, they are in thermal equilibrium with each other.

First Law of Thermodynamics

Energy cannot be created or destroyed; it can be converted from one form to another.

Total Energy of the Universe

Total Energy of the universe is constant.

Change in Energy Equation

ΔEnergy_universe = 0 = ΔEnergy_system + ΔEnergy_surroundings.

Conservation of Energy

The sum of the energy changes in the system and the surroundings must be zero.

Exothermic Process

If the reactants have a lower internal energy than the products, the change in energy will be positive.

Endothermic Process

If the reactants have a higher internal energy than the products, the change in energy will be negative.

Change in Internal Energy Formula

ΔE = change in internal energy.

Heat Symbol

q is heat.

Work Symbol

W is work.

Problem 9.1

A system absorbs 105 kJ of heat and does 29 kJ of work. Calculate the change in internal energy.

Work Done in Process

A system releases 57.5 kJ of heat and its internal energy decreases by 85 kJ. How much work was done in the process?

P-V Work

Work: w = -PΔV where P = pressure, ΔV = change in volume

Work done by the system

Energy lost by the system, w < 0

State Function

A state function is independent of pathway. Only the initial and final positions or states are required.

Path Function

A path function depends on the pathway.

Example of State Function

Internal energy

Example of Path Function

Work and heat

Heat Exchange

Heat is the exchange of thermal energy between the system and surroundings, occurring when there is a temperature difference.

Thermal Equilibrium

Heat flows from matter with high temperature to matter with low temperature until both reach the same temperature.

Heat Capacity

The proportionality constant in the equation q = C × ΔT, with units of J/°C or J/K.

Definition of Heat Capacity

The amount of energy needed to raise the temperature of a substance by 1 K (or 1°C).

Specific Heat Capacity

The amount of energy required to raise the temperature of 1 g of a substance by 1 K (or 1°C).

Specific Heat of Water

Water can absorb a lot of heat energy without a large increase in temperature due to its high specific heat.

Heat Transfer

When two objects at different temperatures are placed in contact, heat flows from the higher temperature material to the lower temperature material until they reach the same final temperature.

Thermal Energy Transfer

A block of metal at 55 ºC is added to water at 25 ºC, transferring heat from the metal to the water.

Factors Affecting Temperature Change

The exact temperature change depends on the mass of the metal, the mass of water, and the specific heat capacities of the metal and of water.

Specific Heat of Granite

0.79 J/g °C

Specific Heat of Lead

0.128 J/g °C

Problem 9.2a

How much heat is needed to raise the temperature of 125.0 g of water from 24.6°C to 46.2 °C? Specific heat of water = 4.18 J/g °C.

Problem 9.2b

A 5.00 g sample of a metal was heated to 100.0 °C and dropped into 100.0 g of water at 25.0 °C.

Specific heat of water

4.18 J/g °C

Specific heat of the metal sample

To be determined from the problem involving a 5.00 g sample of metal heated to 100.0 °C dropped into 100.0 g of water at 25.0 °C with a final temperature of 28.0 °C.

ΔE

ΔE = q + w

Calorimetry at Constant Volume

At constant volume, ΔEsystem = qsystem.

Bomb calorimeter

A sealed, insulated container filled with water used to measure heat changes.

Calorimeter constant

The heat capacity of the calorimeter, the amount of heat absorbed for each degree rise in temperature.

Ccal

34.65 kJ/°C, the calorimeter constant used in the problem involving burning 1.765 g of Ethanol.

Enthalpy (H)

Defined as H = E + PV.

Enthalpy change (ΔH)

The heat evolved in a reaction at constant pressure, ΔHreaction = qreaction at constant pressure.

Exothermic reaction

When ΔH is negative, heat is released by the system into the surroundings, causing the surroundings to feel hot.

Endothermic reaction

When ΔH is positive, heat is absorbed by the system from the surroundings, causing the surroundings to feel cold.

Enthalpy of Reaction

An extensive property where the enthalpy change increases with the amount of reactants used.

Reaction: C3H8 + 5 O2 → 3 CO2 + 4 H2O

ΔH = -2044 kJ.

Reaction: 2CH3OH + 3O2 → 2 CO2 + 2 H2O

ΔH = -1199 kJ.

Heat released from 8.57 g of KClO3

To be determined from the reaction 2 KClO3 → 2 KCl + 3 O2 with ΔH = -89.4 kJ.

Grams of O2 produced from 126.0 kJ

To be determined from the reaction 2 KClO3 → 2 KCl + 3 O2 with ΔH = -89.4 kJ.

Temperature change in calorimetry

Measured indirectly by observing the temperature change in the surroundings.

qsystem

The heat absorbed or released by the system, equal to -qsurroundings.

qsurroundings

The heat absorbed or released by the surroundings, equal to -qsystem.

ΔErxn

The change in energy for a reaction, measured using a bomb calorimeter.

Combustion of propane

An example of an exothermic reaction.

Water boiling on stove

An example of an endothermic reaction.

Calorimetry

The measurement of heat transfer in chemical reactions, typically performed at constant pressure.

qreaction

The heat absorbed or released by the reaction, equal to -qsolution.

DHreaction

The change in enthalpy for a reaction at constant pressure, equal to qreaction.

Hess's Law

A principle stating that the total enthalpy change for a reaction is the sum of the enthalpy changes for the individual steps.

ΔH

The change in enthalpy, which depends only on the initial and final states of a system.

Enthalpy of formation (ΔHf)

The change in enthalpy that accompanies the formation of one mole of a compound from its elements in their standard states.

Standard state

The most stable form of an element under 1 atm pressure and 25°C.

Bond enthalpy

The energy required to break a bond, always a positive quantity.

Average bond energies

Values used to estimate the ΔHrxn based on the number of bonds formed and broken.

ΔHrxn

The change in enthalpy for a reaction, calculated using bond energies or enthalpies of formation.

Problem 9.6

Calculate ΔHrxn per mol of HCl reacted when 0.200 L of 0.200 M HCl is mixed with 0.200 L of 0.200 M NaOH, resulting in a temperature increase from 22.15°C to 23.48°C.

Problem 9.7

Calculate ΔH for the reaction 3H2 (g) + O3 (g) → 3 H2O(g) using given ΔH values.

Problem 9.8

Calculate the ΔH° for the reaction H2 (g) + ½ O2 (g) → H2O(l) using provided enthalpy changes.

Problem 9.9

Calculate the ΔH° for the reaction HCl (g) + NaNO2 (s) → HNO2 (l) + NaCl (s) using given ΔH values.

Calculating Hrxn from Hf

Using the formula ΔH = Σ n ΔHf(products) - Σ m ΔHf(reactants) to find the enthalpy change for a reaction.

Problem 9.10

Calculate the standard enthalpy change for the reactions Al2O3(s) + 2Fe (s) → 2 Al (s) + Fe2O3(s) and NH3 (g) + 5 O2 (g) → 4 NO (g) + 6 H2O (g).

Bond strengths

Measured in terms of bond enthalpy; stronger bonds have higher bond enthalpy and are shorter.

Thermochemistry

The study of the heat energy associated with chemical reactions and changes.