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6. Thermochemistry

Nature of energy

energy: the capacity to do work or to produce heat

law of conservation of energy: energy can be converted from one form to another but can neither be created nor destroyed

potential energy: energy due to position or composition

e.g. water behind a dam, attractive and reactive forces

kinetic energy: energy due to the motion of the object

depends on mass and velocity

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energy can be converted from one form to another

energy is transferrable in two ways

heat: transfer of energy between two objects due to a temperature difference

not a substance contained by an object

frictional heating: kinetic energy transferred to surface as heat

work: force acting over a distance

pathway: specific conditions dictating the way energy transfer is divided between work and heat

total energy transferred will be constant

work and heat are dependent on pathway

state function/state property: property of the system that depends only on its present state

does not depend on the system’s past or future

does not depend on how the system arrived at the present state, only on the characteristics of the present state

change in state function/property in going from one state to another is independent on the pathway taken between the two states

e.g. energy

Chemical energy

system: part of the universe in which one focuses attention

surroundings: everything else in the universe

exothermic: evolution of heat; energy flows out of the system

endothermic: absorption of heat; energy flows into the system

energy gained by the surroundings must be equal to the energy lost by the system

in an exothermic reaction, some of the potential energy stored in the chemical bonds is being converted to thermal energy (random kinetic energy) via heat

thermodynamics: study of energy and its interconversions

first law of thermodynamics: the energy of the universe is constant

internal energy (E): sum of the kinetic and potential energies of all the “particles” in the system

can be changed by flow of work or heat

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thermodynamic quantities always consist of two parts

number: magnitude of change

sign: system’s point of view

positive

endothermic

energy increasing

negative

exothermic

energy decreasing

work done by a gas (expansion) or work done to a gas (compression)

example: gas confined to cylindrical container with movable piston

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F is the force acting on piston of area A

magnitude (size) of work required to expand a gas ΔV against a pressure P

pressure defined as force per unit area

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work defined as force (F) over distance (Δh)

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volume equals area times height

change in volume: final volume - initial volume

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sign of the work

gas (system) is expanding

system is doing work on the surroundings

from system’s point of view, sign should be negative

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expanding: ΔV and w must have have opposite signs

w is negative because work flows out of system

ΔV is positive because volume is increasing

compressed: ΔV is negative (volume decreases) and w is positive (work flows into systeM)

Enthalpy and calorimetry

Enthalpy

enthalpy (H):

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E: internal energy of system

P: pressure of system

V: volume of system

enthalpy is a state function (because internal energy, pressure, volume are state functions)

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at constant pressure (where only PV work is allowed), the change in enthalpy ΔH of the system is equal to the energy flow as heat

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enthalpy of products is greater than reactants: ΔH is positive (endothermic)

enthalpy of products is less than reactants: ΔH is negative (exothermic)

Calorimetry

calorimeter: device used experimentally to determine the heat associated with a chemical reaction

calorimetry: science of measuring heat

heat capacity (C): measure of response to being heated

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specific heat capacity (J/℃ × g or J/K × g): heat capacity per gram of substance

molar heat capacity (J/℃ × mol or J/K × mol): heat capacity per mole of substance

constant-pressure calorimetry: pressure remains constant during the process

used in determining the changes in enthalpy (heats of reactions) for reactions occurring in solution

e.g. with coffee-cup calorimeter

two nested Styrofoam cups with a cover

stirrer and thermometer inserted

example: 50.0 mL of 1.0 M HCl at 25.0℃ with 50.0 mL of 1.0 M NaOH at 25℃ in calorimeter; temperature increased to 31.9℃

net ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l)

energy released as heat → increased random motion of solution components → increased temperature

quantity of energy released can be determined from temperature increase, mass of solution, specific heat capacity

specific heat capacity of water: 4.18 J/℃ × g

4.18 J of energy is required to raise the temperature of 1 g of water by 1℃

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(4.18 J/℃ × g)(1.0 × 10² g)(6.9℃)

2.9 × 10³ J

example: calculating moles of H⁺ ions consumed in preceding experiment

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2.9 × 10³ J heat released when 5.0 × 10⁻² mol H⁺ ions reacted

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magnitude of enthalpy change per mole: 58 kJ/mol

heat is evolved, so ΔH = -58 kJ/mol

extensive property: depends directly on amount of substance

e.g. twice as much solution results in twice as much heat

intensive property: not related to the amount of substance

e.g. temperature

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s: specific heat capacity

m: mass

ΔT: change in temperature

when ΔT is positive, q is positive

constant-volume calorimetry: volume remains constant during the process

e.g. with bomb calorimeter

weighted reactants inside rigid steel container (bomb) and ignited

energy change determined by measuring increase in temperature of water and other calorimeter parts

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change in volume (ΔV) is zero so work is also zero (-PΔV)

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example: measure energy of combustion of octane (C₈H₁₈); 0.5260 g octane placed in bomb calorimeter with heat capacity of 11.3 kJ/℃; temperature increase 2.25℃

energy released = temperature increase × energy required to change temperature by 1℃

ΔT × heat capacity of calorimeter

2.25℃ × 11.3 kJ/℃ = 25.4 kJ

number of moles of octane

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energy released per mole

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reaction exothermic so ΔE is negative

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no work done so ΔE is equal to heat, so q = -5.50 × 10³ kJ/mol

substance

specific heat capacity (J/℃ × g)

H₂O(l)

4.18

H₂O(s)

2.03

Al(s)

0.89

Fe(s)

0.45

Hg(l)

0.14

C(s)

0.71

There are no rows in this table

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Hess’s law

Hess’s law: in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps

example: oxidation of nitrogen to produce nitrogen dioxide

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Characteristics of enthalpy changes

characteristics of ΔH for a reaction

if a reaction is reversed, the sign of ΔH is also reversed

the magnitude of ΔH is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of ΔH is multiplied by the same integer

explanation of rules:

first rule: sign of ΔH indicates direction of heat flow at constant pressure

second rule: ΔH is extensive, depending on amount of substances reacting

calculations

typically require that several reactions be manipulated and combined to give the reaction of interest

work backward from the required reaction, using the reactants and products to decide how to manipulate the other given reactions

reverse any reactions as needed to give the required reactants and products

multiply reactions to give the correct numbers of reactants and products

Standard enthalpies of formation

standard enthalpy of formation (

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): change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states

standard state (degree symbol in thermodynamic function)

for a compound

gaseous substance: pressure of 1 atm

pure substance in condensed state (liquid or solid): pure liquid or solid

substance present in solution: concentration of 1 M

for an element: form in which the element exists at 1 atm and 25℃

example: calculate standard enthalpy change for combustion of methane

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any convenient pathway from reactants to products can be used (Hess’s law)

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reaction (a)

methane being taken apart is the reverse of the formula reaction for methane

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reversing a reaction means keeping magnitude and changing sign

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reaction (b)

oxygen is an element in its standard state, so no change needed

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for reactions (c) and (d), use the elements formed in reactions (a) and (b) to form products

reaction (c)

formation reaction for carbon dioxide

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reaction (d)

formation reaction for water

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2 moles of water required for balanced equation

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change in enthalpy: sum of all ΔH values (including signs)

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the enthalpy change for a given reaction can be calculated by subtracting the enthalpies of formation of the reactants from the enthalpies of formation of the products. Multiply the enthalpies of formation by integers as required by the balanced equation

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Σ: take the sum of the terms

n_p and n_r: moles of product or reactant

elements are not included in the calculation because elements require no change in form

enthalpy calculations

when a reaction is reversed, the magnitude of ΔH remains the same, but its sign changes

when the balanced equation for a reaction is multiplied by an integer, the value of ΔH for that reaction must be multiplied by the same integer

the change in enthalpy for a given reaction can be calculated from the enthalpies of formation of the reactants and products

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elements in their standard states are not included in the ΔH_reaction calculations

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for an element in its standard state is 0

Present sources of energy

woody plants, coal, petroleum, natural gas hold vast amounts of energy that originally came from the sun

plants store energy through photosynthesis that can be claimed by burning the plants themselves or the decay products that have been converted over millions of years to fossil fuels

Petroleum and natural gas

most likely formed from the remains of marine organisms that lived approximately 500 million years ago

natural gas: mostly methane, but also significant amounts of ethane, propane, butane

usually associated with petroleum deposits

tremendous reserves deep in shale deposits

shale is very impermeable and gas does not flow out into wells on its own

hydraulic fracturing/fracking: injecting slurry of water, sand, chemical additives under pressure through well to produce fractures

environmental concerns

potential contamination of ground water

risks to air quality

possible mishandling of wastes

petroleum: thick, dark liquid composed mostly of hydrocarbons

hydrocarbon: compound that contains carbon and hydrogen

composition varies

hydrocarbons with chains of 5-25+ carbons

must be separated into fractions by boiling to be used efficiently

lighter molecules (lowest boiling points) boiled off, leaving heavier ones behind

history

increased demand for lamp oil during Industrial Revolution

first oil well: EdwinDrank (1859)

petroleum refined into kerosene for lamps; gasoline initially discarded

rise of electric lights reduced kerosene use but gasoline became essential to cars

William Burton invented pyrolytic (high-temperature) cracking to increase gasoline yield

tetraethyl lead was added to gasoline to prevent engine knocking, but lead gasoline caused environmental pollution

lead phased out; required extensive and expensive modifications of engines and gasoline refining process

Coal

coal: remains of plants that were buried and subjected to high pressure and heat over long periods of time

plant materials rich in cellulose

complex molecule

CH₂O

molar mass: 500000 g/mol

plants and trees died and buried; chemical changes reduced oxygen and hydrogen content of cellulose molecules

matures through four stages

each stage has a higher carbon-to-oxygen ratio and carbon-to-hydrogen ratio (relative carbon content gradually increases)

lignite

subbituminous

bituminous

anthracite

energy available from combustion increases as carbon content increases

anthracite is the most valuable coal

lignite is the least valuable

coal currently furnishes ~23% of energy in United States

as supply of petroleum dwindles, share of energy supply from coal likely to increase

drawbacks

coal is expensive and dangerous to mine underground

burning of coal (especially high-sulfur) yields air pollutants (e.g. sulfur dioxide, which can lead to acid rain)

carbon dioxide (from when burned) has significant effects

Effects of carbon dioxide on climate

earth receives a tremendous quantity of energy from the sun

Nature of energy

Chemical energy

Enthalpy and calorimetry

Enthalpy

Calorimetry

Hess’s law

Characteristics of enthalpy changes

Standard enthalpies of formation

Present sources of energy

Petroleum and natural gas

Coal

Effects of carbon dioxide on climate

New energy sources

Coal conversion

Hydrogen as fuel

Other energy alternatives

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