Thermodynamics
X Unit 9
Energy: Basic Principles
Thermodynamics – the study of energy
changes
 Energy – the ability to do work or
produce heat


Kinetic Energy – energy of motion

Potential Energy – energy due to position or
composition
Law of Conservation of Energy

A.k.a. First Law of Thermodynamics

Energy can be converted from one form to
another but can’t be created or destroyed

This means the total energy of the
universe is CONSTANT!
Heat vs. Temperature

Temperature – measure of the random
motion of a substance


Temperature is proportional to kinetic
energy (it is a measure of the average
kinetic energy in a substance)
Heat (q) – flow of energy due to a
temperature difference
System vs. Surroundings

A system is the part of the universe we
are studying.
The surroundings
are everything else
outside
of the system.

Direction of Heat Flow

Heat transfer occurs when two objects are
at two different temperatures.

Eventually the two objects reach the same
temperature

At this point, we say that the system
has reached equilibrium.
Thermal Equilibrium

Heat transfer
always occurs with
heat flowing from
the HOT object to
the COLD object.
Exothermic vs. Endothermic
 Exothermic
process  heat is
transferred from the system to the
surroundings
 Heat is lost from the system
(temperature in system decreases)
 Endothermic
process  heat transferred
from the surroundings to the system
 Heat is added to the system
(temperature in system increases)
Units of Energy

Joule (J) is the SI unit of energy & heat
One kilojoule (kJ) = 1000 joules (J)

calorie (cal) = heat required to raise the
temperature of 1.00 g of water by 1 °C
1 calorie = 4.184 J
Units of Energy

Food is measured in Calories (also known
as kilocalories) instead of calories
1 Cal = 1 kcal = 1000 calories
Units of Energy
3800 cal = __________ Cal = _________ J
Units of Energy
The label on a cereal box indicates that 1
serving provides 250 Cal. What is the energy
in kJ?
Heat Transfer
Direction and sign of heat flow – MEMORIZE!
ENDOTHERMIC: heat is added to the
system & the temperature increases (+q)
EXOTHERMIC: heat is lost from the system
(added to the surroundings) & the
temperature in the system decreases (-q)
Specific Heat
(Specific Heat Capacity)

Specific Heat (C) - The quantity of heat
required to raise the temperature of one
gram of a substance by 1 °C

Units:
 J/(g·°C) or J/(g·K)
 cal/(g·°C) or cal/(g·K)
Examples of Specific Heat
At the beach, which gets hotter, the sand or
the water?
Higher specific heat means the substance
takes longer to heat up & cool down!
Examples of Specific Heat

Specific heat (C)= the heat required to
raise the temperature of 1 gram of a
substance by 1 °C
Cwater = 4.184 J/(g°C)
Csand = 0.664 J/(g°C)
Calculating Changes in Thermal E
q = mCΔT
q = heat (J)
m = mass (g)
C = specific heat capacity, J/(g°C)
ΔT = change in temperature, Tfinal – Tinitial
(°C or K)
***All units must match up!!!***
Example
q = mCΔt
How much heat in J is given off by a 75.0 g
sample of pure aluminum when it cools from
84.0°C to 46.7°C? The specific heat of
aluminum is 0.899 J/(g°C).
Example
q = mCΔt
What is the specific heat of benzene if 3450 J
of heat are added to a 150.0 g sample of
benzene and its temperature increases from
22.5 °C to 35.8 °C?
Enthalpy

Enthalpy (H)
 The heat content of a reaction
(chemical energy)

ΔH = change in enthalpy
 The amount of energy absorbed or
lost by a system during a chemical
process
 ΔH
= ΔHfinal - ΔHinitial
Representation of Enthalpy as a Graph
Two Ways to Designate
Thermochemical Equations
Endothermic:
a) H2
(g) + I2 (s)  2 HI (g)
b) H2
(g) + I2 (s) + 53.0 kJ  2 HI (g)
ΔH = 53.0 kJ
Two Ways to Designate
Thermochemical Equations
Exothermic:
a)
½ CH4 (g) + O2 (g)  ½ CO2 (g) + H2O (l)
ΔH = -445.2 kJ
b)
½ CH4(g) + O2(g)  ½ CO2(g) + H2O(l) + 445 .2 kJ
Two Ways to Designate
Thermochemical Equations
Note the meaning of the sign in ΔH in the
equations above!!
Endothermic: ΔH = +
Exothermic: ΔH = -