Chapter 5: Thermochemisty

Chapter 5: thermochemistry

thermodynamics – study of energy and its transformations
thermochemistry – relationships between chemical reactions and energy changes

5.1 The Nature of Energy

w = Fd
heat – energy transferred from hotter object to colder one
energy – capacity to do work

5.1.1 Kinetic and Potential Energy

kinetic energy – energy of motion
potential energy – stored energy
electron has potential energy when near a proton
charged particles have potential energy through attractive and repulsive electrostatic forces
chemical energy – potential energy stored in arrangement of atoms
thermal energy – kinetic energy of molecules

5.1.2 Energy Units

SI unit for energy = joule, J
1J = 1 kg*m²/s²
1cal = 4.184 J
1 Cal = 1000cal = 1kcal

5.1.3 Systems and Surroundings

system – specific amount of matter defined for study
surroundings – everything outside of system

5.2 The First Law of Thermodynamics

first law of thermodynamics – conservation of energy

5.2.1 Internal Energy

total energy of system – sum of all kinetic and potential energy
internal energy – total energy of a system

5.2.2 Relating D E to Heat and Work

system can exchange energy in two ways : heat and work
D E = q + w
both heat added and work done on system increases internal energy
both heat lost and work done by system on surrounding lower internal energy

Sign Conversion for q:

Sign D E = q + w

q>0; heat transferred from surroundings to system

q>0 and w>0; D E>0

q>0 and w<0; sign depends on magnitudes of q and w

q<0 and w>0; sign depends on magnitudes of q and w

q<0 and w<0; D E<0

q<0; heat transferred from system to surroundings

Sign conversion for w:

w>0; work done by surroundings on system

w<0; work done by system on surroundings

5.2.3 Endothermic and Exothermic Processes

endothermic – absorption of heat by system
exothermic – evolution of heat

5.2.4 State Functions

state function – property of system determined by specifying its conditions
value of state function does not depend on particular history of sample only its present conditions
E is a state function - D E depends only on initial and final states

5.3 Enthalpy

enthalpy – heat absorbed or released under constant pressure

state function

change in enthalpy equals heat gained or lost by system when process occurs under constant pressure
+D H – system gains heat, endothermic
-D H – system releases heat, exothermic

5.4 Enthalpies of Reaction

D H = H(products) – H(reactants)
enthalpy of reaction – energy change in a reaction
thermochemical equations – balanced equations that show enthalpy change
guidelines for using thermochemical equations and enthalpy diagrams

1) enthalpy is an extensive property
magnitude of D H directly proportional to amount of reactant consumed in process
2) enthalpy change for reaction equal in magnitude but opposite in sign to D H for reverse reaction
3) enthalpy change for a reaction depends on state of reactants and products

5.5 Calorimetry

calorimetry – measurement of heat flow
calorimeter – measures heat flow

5.5.1 Heat Capacity and Specific Heat

heat capacity – amount of heat required to raise temperature by 1K or 1° C
molar heat capacity – heat capacity of 1mol of substance
specific heat – heat capacity of 1g of substance

5.5.2 Constant Pressure Calorimetry

dilute aqueous solutions have specific heats that are the same as water
temperature increase (+D T) – reaction is exothermic (-q_rxn)

5.5.3 Bomb Calorimeter (constant volume calorimeter)

bomb calorimeter – used to study combustion in reactions
C_calorimeter – heat capacity of calorimeter
Because constant volume – heat transferred corresponds to energy change rather than enthalpy change

5.6 Hess’s Law

Hess’s Law – if a reaction is carried out in a series of steps, D H for the reaction will be equal to the sum of the enthalpy changes for the individual steps
Always get same value of D H for overall reaction, regardless of steps

5.7 Enthalpies of Formation

enthalpy of formation – enthalpy change from formation of a compound (D H_f)
standard state – 1atm and 298K (25° C)
standard enthalpy – enthalpy change when reactants and products in standard states (D H° )
standard enthalpy of formation – change in enthalpy for reaction that forms 1mol of compound from elements, with all substances in standard states (D H_f°)
standard enthalpy of formation of most stable form of any element is zero

5.7.1 Using Enthalpies of Formation to Calculate Enthalpies of Reaction

n, m stoichiometric coefficients of reaction
first term – formation of reaction of products
second term – reverse formation reactions of reactants

5.8 Foods and Fuels

fuel value – energy releases when 1g of material is combusted

5.8.1 Foods

body uses chemical energy from foods for:
maintaining body temperature, drive muscles, construct and repair tissues
fats serve as body’s energy reserve
1) insoluble in water
2) produce more energy per gram
average fuel value for fat – 38kJ/g (9 kcal/g)
average fuel value of proteins and carbohydrates – 17kJ/g (4 kcal/g)

5.8.2 Fuels

the greater then percentage of carbon and hydrogen in a fuel, the higher the fuel value
fossil fuels – coal, petroleum, natural gas
natural gas – gaseous hydrocarbons, mostly methane, small amounts of ethane, propane, and butane
petroleum – mostly hydrocarbons; rest composed of compounds containing sulfur, nitrogen, or oxygen
coal – hydrocarbons of high molecular weight, and compounds of sulfur, oxygen, and nitrogen
coal gasification
coal + steam ® complex mixture ® mixture if CH₄, H₂, and CO (syngas)
syngas – easier to transport, less air pollution

5.8.3 Other Energy Sources

nuclear and solar energy

Sign Conversion for q:	Sign D E = q + w
q>0; heat transferred from surroundings to system	q>0 and w>0; D E>0 q>0 and w<0; sign depends on magnitudes of q and w q<0 and w>0; sign depends on magnitudes of q and w q<0 and w<0; D E<0
q<0; heat transferred from system to surroundings
Sign conversion for w:
w>0; work done by surroundings on system
w<0; work done by system on surroundings