Chapter 5 Thermochemistry - Philadelphia University · 2015. 3. 30. · Thermochemistry . Transfer...

Chapter 5 Thermochemistry

Topics

Energy and energy changes

Introduction to thermodynamics

Enthalpy

Calorimetry

Hess’s Law

Standard enthalpies of formation

Copyright McGraw-Hill 2009 3

5.1 Energy and Energy Changes

Energy is involved in all types of physical and chemical changes

Energy: the capacity to do work or transfer heat

All forms of energy are either

Kinetic

Potential

Kinetic energy - energy of motion

Defining equation

Where m is mass and u is velocity

Thermal energy - one form of kinetic energy associated with random motion

•Changes in thermal energy are monitored via changes in temperature

Potential energy - energy of position or composition

Chemical energy is stored within structural units of chemical substances.

Electrostatic energy is energy resulting from the interaction of charged particles.

•Dependent on charges and distance between charges (Q = charge and d

= distance)

•Defining equation

•+ Eel: repulsive

•− Eel: attractive

Law of conservation of energy

Energy can be converted from one form to another but not created or destroyed.

The total amount of energy in the universe is constant.

Example

•A chemical reaction (potential) gives off heat (thermal)

System is the part of the universe of interest.

•Example

The reactants NaOH and HCl

Surroundings are the rest of the universe.

•Example

When heat is given off from the reaction of NaOH and HCl, the energy is transferred from the system to the surroundings.

system: what is inside the container. Reactants or products of a chemical reaction

surroundings: container ,room, etc.

Energy changes in chemical reactions


The study of the transfer of heat (thermal energy) in chemical reactions.

Exothermic - transfer of heat from the system to the surroundings

2H2(g) + O2(g) 2H2O(l) + energy

Endothermic - the transfer of heat from the surroundings to the system

energy + 2HgO(s) 2Hg(l) + O2(g)

Thermochemistry

Transfer of Energy

Exothermic process

energy is produced in reaction

energy flows out of system

container becomes hot when touched

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

Endothermic process

energy is consumed by the reaction

energy flows into the system

container becomes cold when touched

N2(g) + O2(g) + Energy 2NO(g)

http://college.hmco.com/chemistry/shared/media/zumdahl/visual/visual.html?name=CH09&vis=_v9a

Comparison of Endothermic and Exothermic Processes

Joule (J) is the SI unit for energy.

•The amount of energy possessed by a 2 kg mass moving at a speed of 1 m/s

Units of Energy


Calorie (cal) - commonly used on food labels

1 cal 4.184 J

1000 cal = 1 Cal = 1 kcal

Food calories (Cal) are really 1000 calories (cal).

Calculate the kinetic energy of a neon atom

moving at a speed of 98 m/s.

226

k kg)(98m/s)10(3.3522

1 E

J101.6 22kE

kg103.352g 1000

kg 1

amu

g101.661amu 20.18 26

24

Example

5.2 Introduction to Thermodynamics

Types of systems:

open (exchange of mass and energy)

closed (exchange of energy)

isolated (no exchange)

State functions depend only on initial and final states of the system and not on how the change was carried out.

Energy (E)

Pressure (P)

Volume (V)

Temperature (T)

State function

State Function

energy change is independent of pathway thus it is a state function

(state property)

State function: a property of a system that depends only on its present state (it does not depend on system’s past or future, i.e, it does not depend on how the system arrived at the present state)

work and heat depend on pathway so they are not state functions

First Law of Thermodynamics

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

Based on the law of conservation of energy

Internal energy (U)

It is the sum of potential and kinetic energy of all particles in the system

U = PE + KE

Kinetic energy - molecular motion

Potential energy - attractive/repulsive interactions

The change in internal energy of a system between final (f) and initial (i) states is defined as:

U = Uf Ui For a chemical system

Cannot calculate the total internal energy with any certainty

Can calculate the change in energy of the system experimentally

U = U(products) U(reactants)

Consider:

S(s) + O2(g) SO2(g)

This reaction releases heat, therefore U is negative.

Usystem + Usurroundings = 0

When a system releases heat, some of the chemical energy is released as thermal energy to the surroundings but this does not change the total energy of the universe.

Usystem = Usurroundings

When a system undergoes a change in energy, the surroundings must undergo a change in energy equal in magnitude and opposite in sign.

where q is heat, w is work

Work and heat

U can be changed (- or + ∆U) by a flow of work, heat, or both

Usys = q + w

Sign Conventions of q and w

Calculate the overall change in internal

energy for a system that absorbs 125 J of

heat and does 141 J of work on the

surroundings.

q is + (heat absorbed)

w is (work done)

Usys = q + w = (+125 J) + (141J) = 16 J

Work

common types of work Expansion- work

done by the system

Compression- work done on the system

A

FP

hAPhFw

VPw

expansion +∆V -w

compression -∆V +w

P is external pressure

Work is force applied over distance

Pressure-volume work, w, done by a system is

w = PV

Usys = q + w

Usys = q P V

Constant volume, V = 0

qv = Usys qv means a constant-volume process

5.3 Reactions carried out at constant volume

or at constant pressure

U = q + w

U = qp PV

qp = U + PV

Reactions carried out at constant pressure

Enthalpy and enthalpy changes

Enthalpy is a property of a system definition: H = U + PV

since U, P and V are all state functions, then H is also a state function

The term enthalpy is composed of the prefix en-, meaning "to put into" and the Greek word -thalpein, meaning "to heat", that is “to put heat into”

http://en.wikipedia.org/wiki/Greek_language

Enthalpy change

H = U + PV

The change in enthalpy

H = U + (PV)

When pressure is held

constant

H = U + PV

Since

qp = U + PV

qp = H

At constant pressure

qp = H

Enthalpy of reaction

H is + for endothermic changes.

H is − for exothermic changes.

Copyright McGraw-Hill 2009

Equations that represent both mass and enthalpy changes

Endothermic reaction H2O(s) H2O(l) H = + 6.01 kJ/mol

•

Exothermic reaction

Thermochemical Equations


Comparison of Endothermic and Exothermic Changes

Always specify state of reactants and products.

When multiplying an equation by a factor (n), multiply the H value by same factor.

Reversing an equation changes the sign but not the magnitude of the H.

Thermochemical equation guidelines


Given the following equation,

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

ΔH = 2803 kJ/mol

calculate the energy released when 45.00 g of glucose is burned in oxygen.

kJ 700.0OHC mol 1

kJ 2803

OHC g 180.2

OHC mol 1OHC g 45.00

61266126

61266126

5.4 Calorimetry

The science of measuring

heat changes

calorimeter- device used to experimentally find the heat associated with a chemical reaction

substances respond differently when heated ( to raise T for two substances by 1 degree, they require different amount of heat)

Specific heat (s)

Specific heat (s) - the amount of heat required to raise the temp of 1 g of a substance by 1C.

Units: J/g C

Relation to amount of heat (q)

where q is heat, m is mass, s is specific heat

and T = change in temp

(T = Tfinal Tinitial)


Heat capacity (C) - the amount of heat required to raise the temp of an object by 1C.

Units: J/C

Relation to amount of heat (q)

where q is heat, C is heat capacity

and T = change in temp

(T = Tfinal Tinitial)

Calculate the amount of energy required

to heat 95.0 grams of water from 22.5C

to 95.5C.

T = Tfinal – Tinitial = 95.5 oC − 22.5oC

T = 73.0 oC

q = (95.0 g) (4.184 J/gC) (73.0C)

q = 2.90 x 104 J or 29.0 kJ

Constant-Pressure Calorimetry

uses simplest calorimeter (like coffee-cup calorimeter) since it is open to air

used to find heat exchange

between the system and

surroundings of a chemical

reaction

i.e., changes in enthalpy

for reactions occurring

in a solution

since qP = ∆H

System: Reactants & Products

Surroundings: the water

System: reactants and products

Surroundings: water

Assuming that the calorimeter

does not leak or absorb heat

coffee cup calorimeter

Constant-pressure calorimetry

For an exothermic reaction:


A metal pellet with a mass of 85.00 grams at an original

temperature of 92.5C is dropped into a calorimeter\

with 150.00grams of water at an original temperature of

23.1C. The final temperature of the water and the pellet

Is 26.8C. Calculate the heat capacity and the specific

heat for the metal.

Example

qwater = msT

= (150.00 g) (4.184 J/gC) (3.7C)

= 2300 J (water gained energy)

= -2300 J (pellet released energy)

Heat capacity of pellet: q = CT

C = q/T

= 2300 J/65.7C = 35 J/C

Specific heat of pellet: J/goC

s =35 J/oC

85.00g= 0.41J/goC

bomb calorimeter

Constant-volume calorimetry

Isolated system

Here, the ∆V = 0 so -P∆V = w = 0

∆U = q + w = qV for constant volume

Typical procedure used in a bomb calorimeter

•Known amount of sample placed in steel

container and then filled with oxygen gas

•Steel chamber submerged in known amount of water

•Sample ignited electrically

•Temperature increase of water is determined


A snack chip with a mass of 2.36 g was burned in a bomb calorimeter. The heat capacity of the calorimeter 38.57 kJ/C. During the combustion the water temp rose by 2.70C. Calculate the energy in kJ/g for the chip


qrxn = − Ccal T

= −(38.57 kJ/C) (2.70C)

= − 104 kJ

Energy content is a positive quantity.

= 104 kJ/2.36 g

= 44.1 kJ/g

Food Calories: 10.5 Cal/g

Example 3

Compare the energy released in the combustion of H2 and CH4 carried out in a bomb calorimeter with a heat capacity of 11.3 kJ/°C. The combustion of 1.50 g of methane produced a T change of 7.3°C while the combustion of 1.15 g of hydrogen produced a T change of 14.3°C. Find the energy of combustion per gram for each.

Example 3 methane: CH4

hydrogen: H2

The energy released by H2 is about 2.5 times the energy released by CH4

gkJgkJH

kJCC

kJH

TCH rcalorimete

/555.1/83

83)3.7()3.11(

gkJgkJH

kJCC

kJH

TCH rcalorimete

/14115.1/162

162)3.14()3.11(

5.5 Hess’s Law

Hess’s Law: The change in enthalpy that occurs when reactants are converted to products is the same whether the reaction occurs in one step or a series of steps.

It is sed for calculating enthalpy for a reaction that cannot be determined directly.

Example 1

N2(g) + 2O2(g) 2NO2(g) ∆H = 68 kJ

OR

N2(g) + O2(g) 2NO(g) ∆H = 180 kJ

2NO(g) + O2(g) 2NO2(g)∆H = -112 kJ

N2(g) + 2O2(g) 2NO2(g) ∆H = 68 kJ

Rules

1. If a reaction is reversed, the sign of ∆H must be reversed as well.

because the sign tells us the direction of heat flow at constant P

2. The magnitude of ∆H is directly proportional to quantities of reactants and products in reaction.

If coefficients are multiplied by an integer, the ∆H must be multiplied in the same way.

because ∆H is an extensive property


Given the following equations:

H3BO3(aq) HBO2(aq) + H2O(l) Hrxn = 0.02 kJ

H2B4O7(aq) + H2O(l) 4 HBO2(aq) Hrxn = 11.3 kJ

H2B4O7(aq) 2 B2O3(s) + H2O(l) Hrxn = 17.5 kJ

Find the H for this overall reaction.

2H3BO3(aq) B2O3(s) + 3H2O(l)

2H3BO3(aq) 2HBO2(aq) + 2H2O(l) x 2

Hrxn = 2(−0.02 kJ) = −0.04 kJ

2HBO2(aq) 1/2H2B4O7(aq) + 1/2H2O(l) reverse, ÷2

Hrxn = +11.3 kJ/2 = 5.65 kJ

1/2H2B4O7(aq) B2O3(s)+ 1/2H2O(l) ÷ 2

Hrxn = 17.5 kJ/2 = 8.75 kJ

2H3BO3(aq) B2O3(s) + 3H2O(l)

Hrxn = 14.36 kJ


5.6 Standard Enthalpies of Formation

Symbol: Hf

The enthalpy change that results when 1 mole of a compound is formed from its elements in their standard states.

Hf for an element in its standard

state is defined as zero.

Standard state: 1 atm, 25C

Values found in reference tables

Used to calculate the Hrxn

Standard States

For a compound:

for gas: P = 1 atm

For pure substances, it is a pure liquid or pure solid state

in solution: concentration is 1 M

For an element:

form that exists in at 1 atm and 25°C

O: O2(g) K: K(s) Br: Br2(l)


Hrxn = nHf

(products) mHf

(reactants)

where denotes summation

n and m are the coefficients in the

balanced equation

Defining equation for enthalpy of reaction:

Calculate the Hrxn for the following reaction

from the table of standard values.

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

Hrxn = nHf

(products) - mHf

(reactants)

= [1(393.5) + 2(285.8)] [1(74.8) + 2(0)]

= 890.3 kJ/mol (exothermic)

Example:


Key Points

First law of thermodynamics

Enthalpy (heat of formation; heat of reaction)

State function

Calorimetry

Specific heat

Hess’s law

Calculations involving enthalpy, specific heat, energy

Chapter 5 Thermochemistry - Philadelphia University · 2015. 3. 30. · Thermochemistry . Transfer...

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