Grade 12 Chemistry Notes
Chapter Five
Thermodynamics – the study of energy and its transformations
Thermochemistry – relationships between chemical changes and energy changes that involve heat
Energy
Heat – energy used to change the temperature of an object – (transfer energy)
Work – F*d – (transfer energy)
Ek – the energy of motion
Ep – energy of position
Chemical energy of fuel due to the potential energy in the atom arrangement
One calorie is 4.184 Joules
A system is the portion we study
Everything else is the surroundings
Reactants and products are the system, container and beyond would be the surrounding
Open system where matter and energy can be exchanged with the surroundings (uncovered boiling
water)
Closed system where energy can be exchanged but not matter (piston and cylinder with gas creating
energy)
Isolated system where energy and matter cannot be exchanged (thermos)
Transferring Energy
Work is done when we lift against the force of gravity
Heat is transferred and can be potential energy. Transfers from a hot to cold object
The First Law of Thermodynamics – energy is conserved, not created or destroyed but transferred
Always note that an energy increase in a system is energy lost for the surroundings
Internal energy – sum of all kinetic and potential energies of every part of the system
Has a number, unit and a sign for increase / decrease
Products are final state and reactants are initial state
Concerned with change in energy
∆𝐸 = 𝐸𝑓𝑖𝑛𝑎𝑙 − 𝐸𝑖𝑛𝑡𝑖𝑡𝑖𝑎𝑙
Positive when final > initial
Negative change means system lost energy to surroundings (exothermic)
Positive change means system gained energy (endothermic)
Relate E to Heat and Work
The change in internal energy is the sum of the heat added or liberated, q, and the work done on or by
the system, w.
∆𝐸 = 𝑞 + 𝑤
Head added or work done increases the energy
For q; + means system gained heat, - means system loses heat
For w; + means word done on the system, - means work done by the system
For E; + means net gain of energy by system, - means net loss of energy by system
State functions
A property of a system that is determined by specifying the system’s condition, or state. Does not care
about process to get to a certain state, only the end point and start point.
State functions are not q and w, but E is a state function.
,Enthalpy
In relation to the earths atmospheric pressure
Enthalpy denoted by H, is defined as the internal energy plus the product of the pressure P, and volume,
V, of the system:
𝐻 = 𝐸 + 𝑃𝑉
P and V are state functions. Enthalpy is a state function due to this
P-V Work
The only kind of work produced by chemical or physical changes open to the atmosphere is the
mechanical work associated with a change in volume
An increase in volume means the system does P-V work
Gas produced from a reaction can do work on its surroundings (massless piston)
The work involved in the expansion or compression of gases is called pressure volume work
When pressure is constant in a process the sign and magnitude of the P-V work are given by
𝑤 = −𝑃∆𝑉
To expand (increase volume) the system is doing work (negative sign mandatory)
Negative indicates work done by system on surrounding
Where P is pressure and V is change in volume
Pressure is always positive or zero
When volume expands, volume is also positive, but work is negative
When gas expands, the system does work on the surroundings, indicated by negative w
When gas compresses, surroundings do work, indicated by positive w, smaller volume occupation
𝑤 = 𝑎𝑡𝑚 ∗ 𝐿
1-L atm = 101.3 J
Enthalpy Change
Denoted by ∆𝐻,
∆𝐻 = ∆(𝐸 + 𝑃𝑉)
(constant pressure)
For systems that do no motion/displacement work:
∆𝐻 = 𝑞𝑝
Subscript P indicates constant pressure
Enthalpy is just heat (q) when system does no work and does not change volume
Ex. Heating a beaker of water, no movement, no gas therefor no V change
Endothermic - 𝑞𝑝 = positive (gain heat)
Exothermic - 𝑞𝑝 = negative (lose heat)
H is a state function
The relationship between ∆𝐻 and 𝑞𝑝 has special limitations that only P-V work is involved, and the
pressure is constant.
Enthalpies of Reaction
Because final – initial the enthalpy change for a chemical reaction is
∆𝐻 = 𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − 𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠
The enthalpy change that accompanies a reaction is called either enthalpy of reaction or the heat of
reaction (∆𝐻𝑟𝑥𝑛 )
When giving a numerical value, specify the reaction involved
Negative ∆𝐻 means exothermic
1. Enthalpy is an extensive property – magnitude of ∆𝐻 is proportional to the amount of reactant
consumed in the process
When 1 mol of methane is burned it is -890kJ, when 2 mol is burned it is 2* -890kJ
, 2. The enthalpy change for a reaction is equal in magnitude, but opposite sign, to ∆𝑯 for the
reverse reaction
When reversing a reaction, we reverse the roles of the products and reactants. An exothermic reaction
will turn endothermic.
𝐶𝑂2(𝑔) + 2𝐻2 𝑂(𝑙) → 𝐶𝐻4(𝑔) + 2𝑂2(𝑔) , ∆𝐻 = +890𝑘𝐽 ENDO
(reverse the roles for #1 to lose energy)
3. The enthalpy change for a reaction depends on the states of the reactants and products. Put E
If the products in the equation for #1 were H20 gas, then ∆𝐻𝑟𝑥𝑛 would be -802kJ instead of -890kJ into L to
Less heat would be available for transfer to surroundings because enthalpy of H2O gas is greater than
G state
liquid.
change
Change from liquid to gas is endothermic for water and then that will have less exothermic energy.
For questions, change given mass to moles and then use that value to multiply by the enthalpy for the
new enthalpy
Liquid to gas absorbs energy to the system
State the states of matter
The sign and magnitude of enthalpy is valuable
H2O liquid to gas requires +88kJ/mol
Calorimetry – measure of heat flow
Calorimeter – device used to measure heat flow
Heat Capacity and Specific Heat
Heat Capacity – temp change experienced by an object when it absorbs a certain amount of heat
Heat required to raise temp by 1 K or 1 C
Molar heat capacity – heat capacity of one mol of substance
𝑞 = 𝑚𝑜𝑙 ∗ 𝑚𝑜𝑙𝑎𝑟 ℎ𝑒𝑎𝑡 𝑐𝑎𝑝𝑐𝑖𝑡𝑦 ∗ ∆𝑇
𝑞𝑢𝑎𝑛𝑡𝑖𝑡𝑦 𝑜𝑓 ℎ𝑒𝑎𝑡 𝑡𝑟𝑎𝑛𝑠𝑓𝑒𝑟𝑟𝑒𝑑
Specific Heat - 𝐶𝑠 , (𝑔
𝑜𝑓 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒)∗(𝑡𝑒𝑚𝑝 𝑐𝑎𝑛𝑔𝑒)
𝑞
𝐶𝑠 =
𝑚 ∗ ∆𝑇
𝐽
𝐽=𝑔∗ ∗ ∆℃
𝑔°𝐶
When a sample absorbs heat, its temp increases.
Rearrange for q
Constant pressure calorimeter, ∆𝐻 = 𝑞𝑝 at constant p and no change in volume
The reactants and products of the reaction are the system.
Assume calorimeter is an isolated system
Exothermic heat transferred from the reaction to the water, which we measure.
Or we measure heat from water to reaction
Monitoring temperature allows us to see heat flow
The heat gained or lost by the system, 𝑞𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 , is equal in magnitude but opposite in sign to heat
absorbed or released
𝑞𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = −𝑞𝑟𝑥𝑛(𝑠𝑦𝑠𝑡𝑒𝑚)
𝑞𝑠𝑜𝑙 = 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 ℎ𝑒𝑎𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ∗ 𝑔𝑟𝑎𝑚𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ∗ ∆𝑇 = −𝑞𝑟𝑥𝑛
Temp increase (∆𝑇 > 0) means the reaction is exothermic (𝑞𝑟𝑥𝑛 < 0)
Specific heat of water – 4.18J/g-K
Density gives a 1:1 ratio of mass to volume
∆𝐻𝑠𝑢𝑟 = −∆𝐻𝑟𝑥𝑛
Chapter Five
Thermodynamics – the study of energy and its transformations
Thermochemistry – relationships between chemical changes and energy changes that involve heat
Energy
Heat – energy used to change the temperature of an object – (transfer energy)
Work – F*d – (transfer energy)
Ek – the energy of motion
Ep – energy of position
Chemical energy of fuel due to the potential energy in the atom arrangement
One calorie is 4.184 Joules
A system is the portion we study
Everything else is the surroundings
Reactants and products are the system, container and beyond would be the surrounding
Open system where matter and energy can be exchanged with the surroundings (uncovered boiling
water)
Closed system where energy can be exchanged but not matter (piston and cylinder with gas creating
energy)
Isolated system where energy and matter cannot be exchanged (thermos)
Transferring Energy
Work is done when we lift against the force of gravity
Heat is transferred and can be potential energy. Transfers from a hot to cold object
The First Law of Thermodynamics – energy is conserved, not created or destroyed but transferred
Always note that an energy increase in a system is energy lost for the surroundings
Internal energy – sum of all kinetic and potential energies of every part of the system
Has a number, unit and a sign for increase / decrease
Products are final state and reactants are initial state
Concerned with change in energy
∆𝐸 = 𝐸𝑓𝑖𝑛𝑎𝑙 − 𝐸𝑖𝑛𝑡𝑖𝑡𝑖𝑎𝑙
Positive when final > initial
Negative change means system lost energy to surroundings (exothermic)
Positive change means system gained energy (endothermic)
Relate E to Heat and Work
The change in internal energy is the sum of the heat added or liberated, q, and the work done on or by
the system, w.
∆𝐸 = 𝑞 + 𝑤
Head added or work done increases the energy
For q; + means system gained heat, - means system loses heat
For w; + means word done on the system, - means work done by the system
For E; + means net gain of energy by system, - means net loss of energy by system
State functions
A property of a system that is determined by specifying the system’s condition, or state. Does not care
about process to get to a certain state, only the end point and start point.
State functions are not q and w, but E is a state function.
,Enthalpy
In relation to the earths atmospheric pressure
Enthalpy denoted by H, is defined as the internal energy plus the product of the pressure P, and volume,
V, of the system:
𝐻 = 𝐸 + 𝑃𝑉
P and V are state functions. Enthalpy is a state function due to this
P-V Work
The only kind of work produced by chemical or physical changes open to the atmosphere is the
mechanical work associated with a change in volume
An increase in volume means the system does P-V work
Gas produced from a reaction can do work on its surroundings (massless piston)
The work involved in the expansion or compression of gases is called pressure volume work
When pressure is constant in a process the sign and magnitude of the P-V work are given by
𝑤 = −𝑃∆𝑉
To expand (increase volume) the system is doing work (negative sign mandatory)
Negative indicates work done by system on surrounding
Where P is pressure and V is change in volume
Pressure is always positive or zero
When volume expands, volume is also positive, but work is negative
When gas expands, the system does work on the surroundings, indicated by negative w
When gas compresses, surroundings do work, indicated by positive w, smaller volume occupation
𝑤 = 𝑎𝑡𝑚 ∗ 𝐿
1-L atm = 101.3 J
Enthalpy Change
Denoted by ∆𝐻,
∆𝐻 = ∆(𝐸 + 𝑃𝑉)
(constant pressure)
For systems that do no motion/displacement work:
∆𝐻 = 𝑞𝑝
Subscript P indicates constant pressure
Enthalpy is just heat (q) when system does no work and does not change volume
Ex. Heating a beaker of water, no movement, no gas therefor no V change
Endothermic - 𝑞𝑝 = positive (gain heat)
Exothermic - 𝑞𝑝 = negative (lose heat)
H is a state function
The relationship between ∆𝐻 and 𝑞𝑝 has special limitations that only P-V work is involved, and the
pressure is constant.
Enthalpies of Reaction
Because final – initial the enthalpy change for a chemical reaction is
∆𝐻 = 𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − 𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠
The enthalpy change that accompanies a reaction is called either enthalpy of reaction or the heat of
reaction (∆𝐻𝑟𝑥𝑛 )
When giving a numerical value, specify the reaction involved
Negative ∆𝐻 means exothermic
1. Enthalpy is an extensive property – magnitude of ∆𝐻 is proportional to the amount of reactant
consumed in the process
When 1 mol of methane is burned it is -890kJ, when 2 mol is burned it is 2* -890kJ
, 2. The enthalpy change for a reaction is equal in magnitude, but opposite sign, to ∆𝑯 for the
reverse reaction
When reversing a reaction, we reverse the roles of the products and reactants. An exothermic reaction
will turn endothermic.
𝐶𝑂2(𝑔) + 2𝐻2 𝑂(𝑙) → 𝐶𝐻4(𝑔) + 2𝑂2(𝑔) , ∆𝐻 = +890𝑘𝐽 ENDO
(reverse the roles for #1 to lose energy)
3. The enthalpy change for a reaction depends on the states of the reactants and products. Put E
If the products in the equation for #1 were H20 gas, then ∆𝐻𝑟𝑥𝑛 would be -802kJ instead of -890kJ into L to
Less heat would be available for transfer to surroundings because enthalpy of H2O gas is greater than
G state
liquid.
change
Change from liquid to gas is endothermic for water and then that will have less exothermic energy.
For questions, change given mass to moles and then use that value to multiply by the enthalpy for the
new enthalpy
Liquid to gas absorbs energy to the system
State the states of matter
The sign and magnitude of enthalpy is valuable
H2O liquid to gas requires +88kJ/mol
Calorimetry – measure of heat flow
Calorimeter – device used to measure heat flow
Heat Capacity and Specific Heat
Heat Capacity – temp change experienced by an object when it absorbs a certain amount of heat
Heat required to raise temp by 1 K or 1 C
Molar heat capacity – heat capacity of one mol of substance
𝑞 = 𝑚𝑜𝑙 ∗ 𝑚𝑜𝑙𝑎𝑟 ℎ𝑒𝑎𝑡 𝑐𝑎𝑝𝑐𝑖𝑡𝑦 ∗ ∆𝑇
𝑞𝑢𝑎𝑛𝑡𝑖𝑡𝑦 𝑜𝑓 ℎ𝑒𝑎𝑡 𝑡𝑟𝑎𝑛𝑠𝑓𝑒𝑟𝑟𝑒𝑑
Specific Heat - 𝐶𝑠 , (𝑔
𝑜𝑓 𝑠𝑢𝑏𝑠𝑡𝑎𝑛𝑐𝑒)∗(𝑡𝑒𝑚𝑝 𝑐𝑎𝑛𝑔𝑒)
𝑞
𝐶𝑠 =
𝑚 ∗ ∆𝑇
𝐽
𝐽=𝑔∗ ∗ ∆℃
𝑔°𝐶
When a sample absorbs heat, its temp increases.
Rearrange for q
Constant pressure calorimeter, ∆𝐻 = 𝑞𝑝 at constant p and no change in volume
The reactants and products of the reaction are the system.
Assume calorimeter is an isolated system
Exothermic heat transferred from the reaction to the water, which we measure.
Or we measure heat from water to reaction
Monitoring temperature allows us to see heat flow
The heat gained or lost by the system, 𝑞𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 , is equal in magnitude but opposite in sign to heat
absorbed or released
𝑞𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = −𝑞𝑟𝑥𝑛(𝑠𝑦𝑠𝑡𝑒𝑚)
𝑞𝑠𝑜𝑙 = 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 ℎ𝑒𝑎𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ∗ 𝑔𝑟𝑎𝑚𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ∗ ∆𝑇 = −𝑞𝑟𝑥𝑛
Temp increase (∆𝑇 > 0) means the reaction is exothermic (𝑞𝑟𝑥𝑛 < 0)
Specific heat of water – 4.18J/g-K
Density gives a 1:1 ratio of mass to volume
∆𝐻𝑠𝑢𝑟 = −∆𝐻𝑟𝑥𝑛
