IB Physics HL Topic 3: Thermal Physics

Temperature

- The property that determines the direction of thermal energy transfer between two objects.
- A measure of the average random kinetic energy of the particles of a substance.

State the relation between the Kelvin and Celsius scales of temperature

T(K) = t(°C) + 273

Thermal Equilibrium

When they are at the same temperature so that there is no transfer of thermal energy between them

Internal energy of a substance

The total potential energy and random kinetic energy of the molecules of the substance
Internal = KE+PE

Potential energy

Due to intermolecular forces

Kinetic Energy

- from random motion of molecules
- translational KE from moving in certain direction
- rotational KE from rotating about one or more axes

Thermal Energy (Heat)

Is the non-mechanical transfer of energy between a system and its surroundings

Mole

The basic SI unit for an amount of a substance that contains the same number of atoms as 0.012 kg of carbon-12 (12C)

Molar mass

- The mass of one mole of a substance.
- Mass of 6.02 ⨉10²³ molecules of a substance

Mass number

Nucleon number
Number of nucleons (protons + neutrons) in nucleus

Avogadro constant

The number of atoms in 0.012 kg of 12C ( = 6.02 x 1023), i.e. in 1 mol

Thermal capacity

The energy required to raise the temperature of a substance by 1K

Q =CΔT
Q = Energy transferred in [J],
C = Thermal heat capacity [JK^-1],
ΔT = Temperature change [K]

Specific heat capacity

The energy required per unit mass to raise the temperature of a substance by 1K

Q =mcΔT
Q = energy transferred [J]
m = Mass [kg]
c = The specific heat capacity [JK^-1],
ΔT = Temperature change [K]

Fusion

The change of phase from solid to liquid

Evaporation

The change of phase from liquid to gas

Specific Latent Heat

Energy per unit mass absorbed or released during a phase change

Q=mL
Q = Heat change [J], [Nm]
m = mass [kg]
L = Specific latent heat [Jkg^-1]

Ideal gas

a gas that is
- made up of tiny spherical particles
- Elastic in molecule collisions with another and the container they are in
- No intermolecular force
- Molecules move in a random motion at a constant velocity
- follows the ideal gas equation of state (PV = nRT) for all values of P, V, and T

Real gas

A gas that does not follow the ideal gas equation of state for all values of P, V, and T
a real gas is similar to ideal gas when
- density is low
- Volume is large
- Pressure is low
Temperature is high

Work done by an ideal gas

W = PΔV

W = work done by the gas [J]
P = constant pressure of the gas [Pa]
ΔV = volume increase of the gas [m^3]

Explain the macroscopic behaviour of an ideal gas in terms of a molecular model.

PRESSURE AND VOLUME:
•if volume is reduced, particles hit the walls more often since the walls are closer together.
•force exerted by the molecules equal to rate of change of the momentum
•this increaseS if collides more frequent, resulting in an incease pressure

PRESSURE AND TEMPERATURE:
•increase in temperature increases the speed of the molecules
•When the molecules hit the walls, their change of momentum will be greater and they will hit the walls more often
•result in greater rate of change of momentum and hence a larger force
•results in an increase in pressure
DOING WORK ON A GAS:
•When push piston of a pump it collides with the molecules giving them energy.
•doing work on the gas
•increase in kinetic energy results in an increase in temperature and pressure

GAS DOES WORK:
•when a gas expands pushes on surrounding air
•in pushing, gas does work which requires energy
•this energy comes form the kinetic energy of the molecules, resulting in a reduction in temperature

Absolute zero of temperature

Temperature at which a gas would exert no pressure

First Law of Thermodynamics

(U = ΔU + W)
The thermal energy transferred to a system from its surroundings is equal to the work done by the system plus the change in internal energy of the system.
(same as the principle of conservation of energy)

Q = the energy transferred to the gas from its surroundings [J]
ΔU = the increase in the internal energy of the gas [J]
W = the work done by the gas [J]

Second Law of Thermodynamics

The overall entropy of the universe is increasing.
(OR - All natural processes increase the entropy of the universe.)

(NOTE: The second law implies that thermal energy cannot spontaneously transfer from a region of low temperature to a region of high temperature.)

Isochoric (Isovolumetric)

a process that occurs at constant pressure (ΔP = 0)

Isothermal

a process that occurs at constant temperature (ΔT = 0)

Isobaric

a process that occurs at constant pressure (ΔP = 0)

Adiabatic

a process that occurs without the exchange of thermal energy (Q = 0)

Entropy

a system property that expresses the degree of disorder in the system

State that temperature determines the direction of thermal energy transfer between two objects.

- Thermal energy naturally flows from hot to cold
- Reach the same temperature, when no more transfer of thermal energy, THERMAL EQUILIBRIUM

Distinguish between microscopic and macroscopic.

MICROSCOPIC:
inside the system to see how its component parts interact
MACROSCOPIC:
consider the system as a whole to see how it interacts with its surrounding

Distinguish between the macroscopic and microscopic concept of temperature

MACROSCOPIC:
Hotness or coldness of an object as measured by a thermometer
MICROSCOPIC:
average kinetic energy per molecule of the molecules in the substance

Explain in terms of molecular behaviour why temperature does not change during a phase change.

- energy given does not increase KE as it increases there PE
- Intermolecular bonds being broken require energy
- when substance forms bonds, releases energy

Explain the physical differences between the solid, liquid and gaseous phases in terms of molecular structure and particle motion.

SOLID:
•fixed shape and volume
•molecules held in position by strong force (bonds) between atoms
•vibrate around a mean (average) position
•lowest internal energy
•lowest temperature
•low compression/expansion
LIQUID:
•no fixed shape but fixed volume
•weaker forces (some bonds broken), this keeps molecules close
•more energy
•vibrating but not in completely fixed positing, free to move around each other
•low compression/expansion
GAS:
•very weak forces (all bonds broken), molecules essentially independent
•high energy
•atoms completely free to move at high speed, occasionally collide
•gas fills container
•highest temperature
•high compression/expansion
•No fixed shape or volume, no force between molecules

Total KE of gas

KE =nRT(2/3)

n = Number of mole [mol]
R = Molar gas constant
T = Temperature [K]

KE Average of gas

KEavg =kT (3/2)

k =Boltzmann constant 1.38×10^-2 [JK^-1]
T = Temperature [K]