EnergyA First Look at

Thermodynamics…

Definition…

• Ability to do “work”; capacity of a system to do “work”

• “Work” is defined for macroscopic systems as moving an object against a force. But chemistry deals with many microscopic systems so the physics definition of “work” does not always seem to apply.

Forms…

• Potential – energy due to position

• Kinetic – energy due to motion

• These forms of energy are interconvertible.

EnergyEnergy is the capacity to do work.

less stable

more stable

change in potential energy EQUALSkinetic energy

A gravitational system. The potential energy gained when a lifted weight is converted to kinetic energy as the weight falls.

EnergyEnergy is the capacity to do work.

less stable

more stable

change in potential energy EQUALSkinetic energy

A system of two balls attached by a spring. The potential energy gained by a stretched spring is converted to kinetic energy when the moving balls are released.

EnergyEnergy is the capacity to do work.

less stable

more stable

change in potential energy EQUALSkinetic energy

A system of oppositely charged particles. The potential energy gained when the charges are separated is converted to kinetic energy as the attraction pulls these charges together.

EnergyEnergy is the capacity to do work.

less stable

more stable

change in potential energy EQUALSkinetic energy

A system of fuel and exhaust. A fuel is higher in chemical potential energy than the exhaust. As the fuel burns, some of its potential energy is converted to the kinetic energy of the moving car.

Types of Energy• There are also different types of

energy:chemical

electrical gravitational

heatlight (electromagnetic radiation)

magneticmechanical

nuclear These are also interconvertible.

Laws!

• No matter what the conversion the system must obey the Law of Conservation of Energy.

• Energy cannot be created or destroyed. (but may be converted to mass under certain conditions)

Thermal Energy

• In chemistry we deal primarily with thermal energy. We also encounter light and electrical energies.

• Thermal energy = energy due to molecular or atomic motion

• Heat (q) = transfer of chemical energy due to a temperature difference

• Thermodynamics = study of heat and its transformations.

• Thermochemistry = branch of thermodynamics that deals with the heat involved with chemical and physical changes.

Temperature

• Measure of the “hotness” or “coldness” of something

• Measure of the effect of heat on an object or system

• Measure of particle motion• Heat always flows from an area of

higher temperature to areas of lower temperatures.

Temperature Measurement

• Typically measured using a thermometer which has an expandable liquid (mercury or alcohol) trapped in a sealed cylinder

• US meteorological and heating unit temperatures are expressed using the Fahrenheit scale.

• Water freezes at 32oF and boils at 212oF.

Figure 1.8 The freezing and boiling points of water.

Temperature Measurement

• Other countries and the scientific community use the Celsius scale.

• Water freezes at 0oC and boils at 100oC.

Figure 1.8 The freezing and boiling points of water.

Silberberg, Principles of Chemistry

Temperature Measurement

• There is a problem using either of these scales in certain systems. Because there are negative temperatures certain mathematical relationships generate impossible results.

• Lord Kelvin generated a new scale with the coldest temperature as 0 K (absolute zero). The unit of temperature is a kelvin (K). The term degree is not used.

• Water freezes at -273.15 K and boils at 373.15 K.

Figure 1.8 The freezing and boiling points of water.

Silberberg, Principles of Chemistry

Temperature Conversions

• oF oC

• oC K

oo o oF -32

= C C x 1.8 + 32 = F1.8

oC + 273.15 = K

Systems…• Changes in thermal energy of systems

are determined by transfer of heat (q) observed by changes in temperature.

• First you need to define what your system and surroundings are…

• When heat moves out of a system to the surroundings the process is exothermic for the system.

• When heat moves into a system from the surroundings the process is endothermic for the system.

A system transferring energy as heat only.

Endothermic

Exothermic

Heat UnitsJoule (J)

Calorie (cal)

British Thermal Unit

1 cal = 4.18J

1 J = 1 kg*m2/s2

1 Btu = 1055 J

Specific Heat Capacity

• Substances respond differently to the same quantity of energy.

• Consider a wooden spoon and a metal spoon in a pot of boiling water. Which spoon gets hotter?

• Specific heat capacity (c) – energy required to raise 1 gram of substance by 1oC

J/g. oC (J g-1 oC-1) or kcal/g. oC (kcal g-1 oC-1)

Heat Calculations

q = specifi c heat capacity x mass x temperature

PROBLEM: A layer of copper welded to the bottom of a skillet weighs 125 g. How much heat is needed to raise the temperature of the copper layer from 250C to 300.0C? The specific heat capacity (c) of Cu is 0.387 J/g*K.

SOLUTION:

PLAN: Given the mass, specific heat capacity and change in temperature, we can use q = c x mass x T to find the answer. T in 0C is the same as for K.

q = 125 g (300-25)0Cx x = 1.33x104 J0.387 J

g*K

Enthalpy

The Meaning of Enthalpy

w = - PV

H = E + PV

qp = E + PV = H

H ≈ E in

1. Reactions that do not involve gases.

2. Reactions in which the number of moles of gas does not change.

3. Reactions in which the number of moles of gas does change but q is >>> PV.

H = E + PV

where H is enthalpy

For most of the processes we will encounter this semester we can use ΔH as heat exchanged by a system.

Silberberg, Principles of Chemistry

Enthalpy

ΔH for a reaction is calculated as: ΔHproducts – ΔHreactants

If the ΔH is positive, energy of the products is higher than reactants and the process is endothermic.

If the ΔH is negative, energy of the reactants is higher than products and the process is exothermic.

Enthalpy diagrams for exothermic and endothermic processes.

Enth

alp

y,

HEnth

alp

y,

H

CH4 + 2O2

CO2 + 2H2O

Hinitial

HinitialHfinal

Hfinal

H2O(l)

H2O(g)

heat out

heat inH < 0 H > 0

A Exothermic process B Endothermic process

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

H2O(l) H2O(g)

Determining the heat exchangein a chemical change

A bomb calorimeter

AMOUNT (mol)

of compound A

AMOUNT (mol)

of compound B

HEAT (kJ)

gained or lost

molar ratio from balanced equation

Hrxn (kJ/mol)

Summary of the relationship between amount (mol) of substance and the heat (kJ) transferred during a reaction.

Using the Heat of Reaction (Hrxn) to Find Amounts

SOLUTION:

PROBLEM: The major source of aluminum in the world is bauxite (mostly aluminum oxide). Its thermal decomposition can be represented by

If aluminum is produced this way, how many grams of aluminum can form when 1.000x103 kJ of heat is transferred?

Al2O3(s) 2Al(s) + 3/2O2(g) Hrxn = 1676 kJ

1.000x103 kJ x 2 mol Al

1676 kJ

26.98 g Al

1 mol Al

= 32.20 g Al

Reaction Progress

• All reactions have an energy barrier that must be overcome.

Energy of reactants

Energy of products

Activation energy (Ea)

This is an exothermicreaction.

Reaction Progress

Energy of reactants

Energy of products

Activation energy (Ea)

This diagram illustrates anendothermic reaction.

Reaction Enthalpy

ΔH

The difference between the starting and ending energiesis ΔH.

Reaction energy diagram of a catalyzed and an uncatalyzed process.

Catalysts reduce the activation energy of a reaction.

Disorder aka Entropy

• Systems tend to greater disorder or greater energy dispersal.

• Disorder is the preferred state.• Energy dispersal or system disorder

is called entropy.

• All systems work to achieve lowest energy and greatest disorder or a balance between these goals.