Chapter 7: Energy and Chemical Change

• Energy is the ability to do work and supply heat

• Work is motion against an opposing force

221 KE Energy Kinetic mv

• Potential energy (PE) is the energy of position or internal arrangement

• KE can be converted into PE and vice versaWhen the child is at points (a) and (c) they have only PE; at point (b) only KE. Total energy is conserved (PE + KE = constant).

When fully compressed or extended only PE; at natural length only KE. Total energy is conserved.

• The SI unit of energy is the joule (J)

• A 2 kg object moving at 1 meter per second has 1 J of kinetic energy

• You may also encounter the calorie (cal)

(exactly) J 4.184 cal 1

s m kg 1 kg) (2 J 1 -222

s 1m 1

21

• The energy that is transferred as heat comes from the object’s internal energy

• The energy associated with the motion of the object’s molecules is referred to as its molecular kinetic energy

• The internal energy is often given the symbol E or U

• We are interested in the change in E:

reactantsproductsinitialfinal or - E EE - E EE

• The temperature of an object is related to the average kinetic energy of its atoms and molecules

The temperature for curve (1) is lower than for curve (2) because the average kinetic energy is lower.

• Heat is a transfer of energy due to a temperature difference

• The object we are interested in is called the system

• Everything outside the system is called the surroundings

• A boundary separates the system from the surroundings

– Open systems can gain or lose mass and energy across their boundaries

– Closed systems can absorb or release energy, but not mass, across their boundaries

– Isolated systems cannot exchange energy or matter with their surroundings

– When heat is gained by an object, it is written as a positive number

– When heat is lost by an object, it is written as a negative number

• A spontaneous change is one that continues on its own– Heat flows spontaneously from a warmer to

colder object

• The heat directly gained or lost by an object is directly proportional to the temperature change it undergoes

• The object’s specific heat (C) relates the heat (q) to the objects temperature change

• The heat capacity is the amount of heat needed to raise the object’s temperature by one degree Celsius and has the units J/°C

• C is an extensive property that can be determined from experiment, and is proportional to the sample mass

t C ) - t C (tq initialfinal

• The specific heat capacity (s) is an intensive property, and is unique for each substance

• For example

-1-1 C g J of units with or C/ms m sC

4.18 Water

0.235 Silver

0.129 Gold

0.387 Copper

C) (25 C g J Substance

Heat Specific 1-1-

A large specific heat means the substance releases a large amount of heat as it cools. The heat absorbed or released by an substance is then:

t m s q

Example: The temperature of 251 g of water is changed from 25.0 to 30.0 °C. How much heat was transferred to the water?

ANALYSIS: Connect heat to the temperature change.

SOLUTION:

Note: Heat was absorbed because q is positive

kJ 5.25 J 5250

C25.0) - (30.0)C g J (4.18g) (251

1-1-

t m sq

• Chemical bonds are the net attractive force between nuclei and electrons in compounds

• Breaking a chemical bond requires energy

• Making a chemical bond releases energy• The potential energy that resides in

chemical bonds is called chemical energy

The attraction of the electrons for the nuclei in the hydrogen molecule is strong enough to overcome the nucleus-nucleus and electron-electron repulsions.

• Chemical reactions generally involve both making and breaking chemical bonds

• The net gain or loss of energy is often in the form of heat

• Any reaction where heat is a product is called exothermic

• Reactions that consume energy are called endothermic

• Reactions can release heat by replacing “weak” bonds with “strong” ones

• The amount of heat absorbed or released by a chemical reaction is called the heat of reaction

• A calorimeter can be used to measure the heat of reaction

• Calorimeters are usually designed to measure heats of reaction under conditions of constant volume or constant pressure

• Pressure is the amount of force acting on a unit area: area

force pressure

• Atmospheric pressure is the pressure exerted by the mixture of gases in the atmosphere

• At sea level the atmospheric pressure is about 14.7 lb/in²

• Other common pressure units are the atmosphere (atm) and bar:

14.696 lb/in² = 1.0000 atm = 1.0133 bar

• qv and qp are used to show heats measured at constant volume or pressure, respectfully

• In reactions where gases are produce or consumed qv and qp can be very different

• If the volume change is , work (w) is

Pressure-volume work (a) A gas confined under pressure. (b) The gas does pressure-volume work on the surroundings when it expands.

pressure opposing theis where PVPw V

• Note that the work of expansion is negative

• Work and heat are alternate ways to transfer energy

• Their sum is the change in internal energy the system undergoes

• This is a statement of the first law of thermodynamics, which says that energy cannot be created or destroyed

wqE

• Heat and work are not state functions because they depend on the path between the final and initial state

Heat and work depend on the path. Both paths give the same value for the internal energy change. In path 1, this appears as heat (q). In path 2 this appears mostly as work (w).

• The heat produced by a combustion reaction is called the heat of combustion

• The heats are measured in closed containers because the reactions consume and produce gases

• The instrument used to measure these heats is called a bomb calorimeter

• The reaction is run at constant volume so that

vqE

• Heats of reactions in solution are usually run in open containers at constant pressure

• They may transfer heat and expansion work

A bomb calorimeter. The reaction chamber has a constant volume so no work is done. The heat released by the combustion is absorbed by the “bomb” and surrounding water.

• The heat change measured at constant pressure is the enthalpy, H

• Enthalpy is also a state function– is negative for an exothermic process– is positive for an endothermic process

• The the difference in the values of the internal energy and enthalpy change can be large for reactions that consume or release gases

preactantsproducts qVPEH-HH

HH

• The amount of heat that a reaction produces or absorbs depends on the number of moles of reactant that react

• A set of standard states have been defined for reporting heats of reactions

A coffee cup calorimeter can be used to measure heats of reaction at constant pressure. Heat can be released or absorbed, resulting in a change in temperature of the solution.

• Standard thermodynamic states are: 1 bar pressure for all gases and 1 M concentration for aqueous solutions

• A temperature of 25 °C (298 K) is often specified as well

• The standard heat of reaction is the value of the enthalpy change occurring under standard conditions involving the actual number of moles specified the the equation coefficients

• An enthalpy change for standard conditions is denoted

• For example, the thermochemical equation for the production of ammonia from it elements at standard conditions is:

• The physical states are important

• The law of conservation of energy requires

kJ 38.92 )(2NH)(3H)(N 322 Hggg

kJ 92.38 )(H3)(N)(NH2 223 Hggg

H

• Enthalpy is a state function

• An enthalpy diagram is a graphical representation of alternate paths between initial and final states

Remember to include the physical states of reactants and products in thermochemical equations.

Two paths for the formation of carbon dioxide gas. Each give the same enthalpy change.

• Enthalpy changes for reactions can be calculated by algebraic summation

• This is called Hess’s Law: The value of the enthalpy change for any reaction that can be written in steps equals the sum of the values of the enthalpy change of each of the individual steps.

• Enthalpy changes for a huge number of reactions may be calculate using only a few simple rules

• Rules for Manipulating Thermochemical Equations:1) When an equation is reversed the sign of the

enthalpy change must also be reversed.

2) Formulas canceled from both sides of an equation must be for substances in identical physical states.

3) If all the coefficients of an equation are multiplied or divided by the same factor, the value of the enthalpy change must likewise be multiplied or divided by that factor.

• An enormous database of thermochemical equations have been compiled:– The standard heat of combustion is the amount

of heat released when 1 mol of a fuel completely burns in pure oxygen gas with all products brought to 25 °C and 1 bar

Standard heats of combustion are always negative and produce water in liquid form

kJ -890 )O(2H)(CO)(2O)(CH C2224 Hlggg

– The standard enthalpy of formation of a substance is the amount of heat absorbed when 1 mole of the substance if formed at 25 °C and 1 bar from its elements in their standard states

• The standard enthalpy of formation for elements in their standard states are zero

• These are the values most commonly used to calculated standard enthalpy changes for reactions

• Standard enthalpies of formation are given in Table 7.2 and Appendix C

kJ/mol 9.285 )(OH)(O)(H f2221

2 Hlgg

• Hess’s law can be restated in terms of standard enthalpies of formation:

Example: Calculate the enthalpy of reaction for 2NO(g)+O2(g)2NO2(g)

ANALYSIS: Use Hess’s law and Table 7.2

SOLUTION:

reactants theof

all of of Sum

products theof

all of of Sum ffreaction

HHH

kJ -113.1kJ) 0kJ) (2(90.37-kJ) 2(33.8

2222 O fNO fNO f

HHHH