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energy: the capacity to do work or to produce heatlaw of conservation of energy: energy can be converted from one form to another but can neither be created nor destroyedpotential energy: energy due to position or compositione.g. water behind a dam, attractive and reactive forceskinetic energy: energy due to the motion of the objectdepends on mass and velocity
energy can be converted from one form to anotherenergy is transferrable in two waysheat: transfer of energy between two objects due to a temperature differencenot a substance contained by an objectfrictional heating: kinetic energy transferred to surface as heatwork: force acting over a distancepathway: specific conditions dictating the way energy transfer is divided between work and heattotal energy transferred will be constantwork and heat are dependent on pathwaystate function/state property: property of the system that depends only on its present statedoes not depend on the system’s past or futuredoes not depend on how the system arrived at the present state, only on the characteristics of the present statechange in state function/property in going from one state to another is independent on the pathway taken between the two statese.g. energy
system: part of the universe in which one focuses attentionsurroundings: everything else in the universeexothermic: evolution of heat; energy flows out of the systemendothermic: absorption of heat; energy flows into the systemenergy gained by the surroundings must be equal to the energy lost by the systemin an exothermic reaction, some of the potential energy stored in the chemical bonds is being converted to thermal energy (random kinetic energy) via heatthermodynamics: study of energy and its interconversionsfirst law of thermodynamics: the energy of the universe is constantinternal energy (E): sum of the kinetic and potential energies of all the “particles” in the systemcan be changed by flow of work or heat
thermodynamic quantities always consist of two partsnumber: magnitude of changesign: system’s point of viewpositiveendothermicenergy increasingnegativeexothermicenergy decreasingwork done by a gas (expansion) or work done to a gas (compression)example: gas confined to cylindrical container with movable pistonF is the force acting on piston of area Amagnitude (size) of work required to expand a gas ΔV against a pressure Ppressure defined as force per unit area
work defined as force (F) over distance (Δh)
volume equals area times heightchange in volume: final volume - initial volume
sign of the workgas (system) is expandingsystem is doing work on the surroundingsfrom system’s point of view, sign should be negative
expanding: ΔV and w must have have opposite signsw is negative because work flows out of systemΔV is positive because volume is increasingcompressed: ΔV is negative (volume decreases) and w is positive (work flows into systeM)
enthalpy (H): 
E: internal energy of systemP: pressure of systemV: volume of systementhalpy is a state function (because internal energy, pressure, volume are state functions)
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
enthalpy of products is greater than reactants: ΔH is positive (endothermic)enthalpy of products is less than reactants: ΔH is negative (exothermic)
calorimeter: device used experimentally to determine the heat associated with a chemical reactioncalorimetry: science of measuring heatheat capacity (C): measure of response to being heated
specific heat capacity (J/℃ × g or J/K × g): heat capacity per gram of substancemolar heat capacity (J/℃ × mol or J/K × mol): heat capacity per mole of substanceconstant-pressure calorimetry: pressure remains constant during the processused in determining the changes in enthalpy (heats of reactions) for reactions occurring in solutione.g. with coffee-cup calorimetertwo nested Styrofoam cups with a coverstirrer and thermometer insertedexample: 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 temperaturequantity of energy released can be determined from temperature increase, mass of solution, specific heat capacityspecific heat capacity of water: 4.18 J/℃ × g4.18 J of energy is required to raise the temperature of 1 g of water by 1℃
(4.18 J/℃ × g)(1.0 × 10² g)(6.9℃)2.9 × 10³ Jexample: calculating moles of H⁺ ions consumed in preceding experiment
2.9 × 10³ J heat released when 5.0 × 10⁻² mol H⁺ ions reacted
magnitude of enthalpy change per mole: 58 kJ/molheat is evolved, so ΔH = -58 kJ/mol extensive property: depends directly on amount of substancee.g. twice as much solution results in twice as much heatintensive property: not related to the amount of substancee.g. temperatures: specific heat capacitym: massΔT: change in temperaturewhen ΔT is positive, q is positiveconstant-volume calorimetry: volume remains constant during the processe.g. with bomb calorimeterweighted reactants inside rigid steel container (bomb) and ignitedenergy change determined by measuring increase in temperature of water and other calorimeter partschange in volume (ΔV) is zero so work is also zero (-PΔV)
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 calorimeter2.25℃ × 11.3 kJ/℃ = 25.4 kJnumber of moles of octane
energy released per mole
reaction exothermic so ΔE is negative
no work done so ΔE is equal to heat, so q = -5.50 × 10³ kJ/mol
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 stepsexample: oxidation of nitrogen to produce nitrogen dioxide
characteristics of ΔH for a reactionif a reaction is reversed, the sign of ΔH is also reversedthe 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 integerexplanation of rules:first rule: sign of ΔH indicates direction of heat flow at constant pressuresecond rule: ΔH is extensive, depending on amount of substances reactingcalculationstypically require that several reactions be manipulated and combined to give the reaction of interestwork backward from the required reaction, using the reactants and products to decide how to manipulate the other given reactions
AP Chemistry
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6. Thermochemistry
Nature of energy
Chemical energy
Enthalpy and calorimetry
Enthalpy
Calorimetry
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
Hess’s law
Characteristics of enthalpy changes
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