1

Chapter 9: Energy and Chemistry

• Terminology

• Conservation of Energy

• Heat Capacity and Calorimetry

• Enthalpy

• Hess’s Law and Heats of Reaction

• Energy and Stoichiometry

The rate of a reaction depends on the pathway from reactants to products; this is the domain of kinetics. Thermodynamics tells us whether a reaction is spontaneous based only on the properties of the reactants and products. The predictions of thermodynamics do not require knowledge of the pathway between reactants and products.

2

Thermodynamics

– is concerned with energy transfer.

– predicts whether a reaction can occur or not.

– predicts how much heat is given out or taken in during a reaction.

– it is not concerned with how fast a reaction will occur – that is kinetics

Forms of Energy

• Two broad categories of energy: potential energy and kinetic energy.– Potential energy - associated with the relative position of an object.– Kinetic energy - associated with motion.

Kinetic Energy = ½mv²

• Internal energy - the combined kinetic and potential energies of atoms and molecules that make up an object or system.

• Chemical energy - energy released or absorbed during a chemical reaction. Chemical energy is a form of potential energy

• Other forms of energy include radiant, mechanical, thermal, electrical, and nuclear.

• Thermochemistry - the study of the energetic consequences of chemistry

• The Joule is the SI unit of energy.

– 1 Joule = 1 kg m2/s2

3

Energy Transformation and Conservation of Energy

• During energy transformation, the total energy must be conserved.– The sum of all energy conversions and energy transfers

must equal the total energy present which must remain constant.

• To account for energy transformations and conversions, the system and surroundings must be specified.– System - the part of the universe being considered.– Surroundings - the remainder of the universe.– System + Surroundings = Universe– System and surroundings are separated by a boundary.

Transfer of Energy

In the initial position, ball A has a higher potential energy than ball B.

After A has rolled down the hill, the potential energy lost by A has been converted to random motions of the components of the hill (frictional heading) and to the increase in the potential energy of B.

4

Heat

Energy is transferred between a system and its surroundings as a result of a temperature difference.

• Heat always flows from hotter to colder.– Temperature may change.– Phase may change (an isothermal process).

• This directionality corresponds to the spreading out of energy over the greatest number of atoms, molecules or ions.

Exothemeric and Endothermic

• Heat exchange accompanies chemical reactions.

– Exothermic: Heat flows out of the system (to the surroundings).

– Endothermic: Heat flows into the system (from the surroundings).

Exothermic

System is losing energy

Endothermic

System is gaining energy

5

The combustion of methane releases the quantity of energy change ∆(PE) to the surroundings via heat flow. This is an exothermic process.

The energy diagram for the reaction of nitrogen and oxygen to form nitric oxide. This is an endothermic process: Heat {equal in magnitude to ∆(PE)} flows into the system from the surroundings.

6

Work

• Work is the transfer of energy accomplished by a force moving a mass some distance against resistance.

– Pressure-volume work (PV-work) is the most common work type in chemistry.

• Releasing an inflated balloon before it is tied off illustrates an example of PV-work.

– Electrical work (IVt) is used in electrochemistry.• Batteries, electrolysis

Pressure-Volume Work

• In addition to heat effects chemical reactions may also do work.

• Gas formed pushes against the atmosphere.

• Volume changes.

• Pressure-volume work.

7

Pressure-Volume Work

w = F x d

= (P x A) x h

= PV

w = -PextV

Calculating Pressure-Volume Work.

Suppose the gas in the previous figure is 0.100 mol He at 298 K. How much work, in Joules, is associated with its expansion at constant pressure.

Pressure-Volume Work Examples

What is the maximum flow of energy in the form of work done during an isothermal expansion?

8

Internal Energy

– Total energy (potential and kinetic) in a system.•Translational kinetic energy.

•Molecular rotation.

•Bond vibration.

•Intermolecular attractions.

•Chemical bonds.

•Electrons.

• The total internal energy of a sample depends on:– Sample’s temperature

– The type of material in the sample

– The amount of material in the sample

Energy Transformation

• A system contains only internal energy.– A system does not contain heat or work.

– These only occur during a change in the system.

E = q + w

E = change in system’s internal energy

q = heat flow into system (positive value means energy is flowing into the system)

w = work done on system (positive value means energy is flowing into the system)

9

Work and Heat

State Functions

• Any property that has a unique value for a specified state of a system is said to be a State Function (T, V, H, S, P, E, G).

• Depends only on the present state of the system - not how it arrived there.

• It is independent of pathway

– Water at 293.15 K and 1.00 atm is in a specified state. – d = 0.99820 g/mL– This density is a unique function of the state.– It does not matter how the state was established.

10

Functions of State

• Energy is a function of state.– Not easily measured.

• The difference in energy has a unique value between two states.– Is easily measured.

Path Dependent Functions

• Changes in heat and work are not functions of state.– Remember pressure-volume work example,

w = -113 J in a one step expansion of gas:

– When expansion is done to get maximum work

w = -151 J in a reversibly expansion of gas:

11

Energy Transformation and Conservation of Energy

• First law of thermodynamics states that energy can be transformed from one form to another but cannot be created or destroyed.

Eunivierse = Esurroundings + Esystem = 0

– The energy of an isolated system is constant

Heat Capacity and Calorimetry

• Calorimetry is a laboratory method for observing and measuring the flow of heat into and out of a system.

• Different systems will absorb different amounts of energy based on three main factors.– The amount of material, m or n.

• m is mass and n is number of moles

– The type of material, as measured by c or Cp.• c is the specific heat capacity, or specific heat, and Cp is the

molar heat capacity.

– The temperature change, T.

12

Heat Capacity and Specific Heat

• The specific heat capacity, or specific heat, is a physical property of a substance that describes the amount of heat required to raise the temperature of one gram of a substance by 1ºC.– Represented by c.– Specific heat is compound and phase specific.

• The molar heat capacity is a physical property of a substance that describes the amount of heat required to raise the temperature of one mole of a substance by 1ºC.– Represented by Cp.– Molar heat capacity is compound and phase specific.

Heat Capacity and Specific Heat

• The amount of heat energy absorbed can be quantified.

– Specific heat capacity, c. q = mcT

• System is one gram of substance.

– Molar heat capacity, Cp. q = nCpT

• System is one mole of substance.

– Heat capacity q = CT

• Mass dependent specific heat.

13

Heat Capacity and Specific Heat

• Specific heat and molar heat capacities for some common substances.

Calorimetry

• Heat flow is measured using a calorimeter.

• A calorimeter measures the heat evolved or absorbed by the system of interest by measuring the temperature change in the surroundings.

qsystem - qsurroundings = 0

qsystem = -qsurroundings

qgained = -qlost

14

Determination of Specific Heat

Calculate the specific heat of lead.

100°C

Coffee Cup Calorimeter

• A simple calorimeter.

– Well insulated and therefore isolated.

– Measure temperature change.

• Reaction done at constant pressure

qrxn = -qcal

15

Bomb Calorimeter

Reaction done at constant volume

qrxn = -qcal

qcal = qbomb + qwater + qwires +…

Define the heat capacity of the calorimeter:

qcal = miciT = CT

all i

Calorimetry

• There are two steps in a calorimetric measurement.– Calibration - the calorimeter constant, Ccalorimeter, is

determined by dividing the known amount of heat released in the calorimeter by the temperature change of the calorimeter.

Ccalorimeter = q/T

– Actual Measurement - heat released or absorbed in a reaction of known quantity of material is measured.

qcal = CcalorimeterT

qrxn = -qcal

16

Heat of Reaction Example

• Using Bomb Calorimetry Data to Determine a Heat of Reaction.

• The combustion of 1.010 g sucrose, in a bomb calorimeter, causes the temperature to rise from 24.92 to 28.33°C. The heat capacity of the calorimeter assembly is 4.90 kJ/°C.

(a) What is the heat of combustion of sucrose, expressed in kJ/mol C12H22O11

(b) Verify the claim of sugar producers that one teaspoon of sugar (about 4.8 g) contains only 19 calories.

Reaction Conditions

• The conditions under which heat flow, q, occurs will have an impact on the measurement that is made.

– Combustion of octane releases 5.45 x 103 kJ under constant volume conditions, represented as qv.

– Combustion of octane releases 5.48 x 103 kJ under constant pressure conditions, represented as qp.

17

Heat Flow at Constant Volume

• The internal energy change for a reaction equals the sum of the heat flow and the work.

E = q + w

– During an expansion, w = –PV.

E = q - PV

– Under constant volume conditions, V = 0, and E = qv.

• The change in energy is the heat flow under conditions of constant volume.

Heat Flow at Constant Pressure

• Consider a process done under the condition of constant pressure and with only PV-work.

E = qp + w = qp – PV

• Define Enthalpy, H, as H= E + PV

• The enthalpy change can be expressed as H = E +(PV)

H = qp - PV + PV +VP

H = qp

• The change in enthalpy is the heat flow under conditions of constant pressure.

18

Exothermic and Endothermic

• When a system releases heat, the process is said to be exothermic.– The value of H is less than zero; the sign on H is

negative.– Enthalpy of products is less than enthalpy of reactants

• When a system absorbs heat, the process is said to be endothermic.– The value of H is greater than zero; the sign on H is

positive.– Enthalpy of products is greater than enthalpy of

reactants

H of Phase Changes

• Phase changes occur under constant pressure conditions.– The heat flow during a phase change is an enthalpy change.– During a phase change, temperature does not change with heat flow due to

formation or breaking of intermolecular attractive forces.

19

∆H of Phase Changes

• The heat required to convert a liquid to a gas is the heat of vaporization, Hvap.– Hvap is endothermic with a positive value.

• The heat released to convert a gas to a liquid is the heat of condensation, Hcond.– Hcond is exothermic with a negative value.

• Hcond = –Hvap

– The values of enthalpy changes in opposite directions have equal numeric values and differ only in their signs.

– The magnitude of enthalpy change depends on the substance involved.

∆H of Phase Changes of Water

20

∆H of Phase Changes

• The value of H for a phase change is compound specific and has units of kJ/mol.

– The heat flow can be calculated using the number of moles of substance, n, and the value of the enthalpy change.

H = nHphase change

Heating Curve

• The enthalpy change for the conversion of ice to liquid and then to steam can be calculated.– A heat curve breaks the calculation down into

specific heat calculations (sections of the heat curve where temperature changes) and phase change enthalpy calculations (sections of the heat curve where temperature does not change).

21

Heating Curve for Water

• Heat curve for the heating of 500-g of ice at -50oC to 200oC.

Break the problem into two steps: Raise the temperature of the liquid first then completely vaporize it. The total enthalpy change is the sum of the changes in each step.

Enthalpy Changes Accompanying Changes in States of Matter.

Calculate H for the process in which 50.0 g of water is converted from liquid at 10.0°C to vapor at 25.0°C.

Changes of State of Matter Example

22

Changes of State of Matter Example

A 25.0 g ice cube at 0.0°C is added to 100.0 g of liquid water at 22.0°C in a thermally isolated container. The heat of fusion of ice is 6.01 kJ/mol and the specific heat of liquid water is 4.18 J g-1 °C-1.

(a) What is the final condition reached – liquid only, solid only, or a mixture of solid or liquid?

(b) What is the final temperature?

Vaporization and Electricity Production

• Schematic diagram of the important elements of a standard electric power plant.

23

Heat of Reaction

• Enthalpy changes can be calculated for chemical reactions, in addition to temperature changes and phase transitions.

– The enthalpy change is commonly referred to as the heat of reaction.

Standard States and Standard Enthalpy Changes

• Define a particular state as a standard state.

• Standard enthalpy of reaction, H°

– The enthalpy change of a reaction in which all reactants and products are in their standard states.

• Standard State– Element

• The form [N2(g), K(s)] in which it exists at 1 atm and 25°C.

– Compound• For a gas, pressure is exactly 1 atmosphere.

• For a solution, concentration is exactly 1 molar.

• Pure substance (liquid or solid), it is the pure liquid or solid.

24

Bonds and Energy

• The enthalpy change for a reaction can be estimated using bond energies.

• During a chemical reaction, reactant bonds are broken and product bonds are made.– Breaking bonds requires energy.– Making bonds releases energy.

• If the amount of energy released making product bonds is greater than the amount of energy required to break reactant bonds, the reaction is exothermic. If the energy released is less than the energy required, the reaction is endothermic.

FormedBondsBrokenBonds

DDH

Bonds and Energy

CH4 + 2O2 CO2 + 2H2O

• The combustion of methane breaks 4 C-H bonds and 2 O=O bonds. 2 C=O bonds and 4 O-H bonds are made.

25

Bonds and Energy

• The accuracy of enthalpy changes calculated from tabulated bond energies is not very good.– The bond energies used are averages.– Bond energy method used to estimate enthalpy changes for reactions

involving compounds with no available thermochemical data.

• A thermochemical equation summarizes the overall energetics for a chemical reaction.– The sign on the H indicates whether the reaction is endothermic or

exothermic

CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = -890.4 kJ

• The combustion of methane is an exothermic reaction and releases 890.4 kJ of heat energy when 1 mole of methane reacts with 2 moles of oxygen.

Heats of Reaction for Some Specific Reactions

• Some classes of chemical reactions are given their own labels for heats of reactions.– Heat of combustion, Hcomb

– Heat of neutralization, Hneut

– Heat of formation, Hf, is the heat of reaction for formation of substances.

• Fractional coefficients are allowed for formation reactions because only one mole of product can be formed.

C(s) + ½O2(g) CO(g)