thermal physics Flashcards
definition of specific heat capacity
energy required to raise the temperature of 1kg of a substance by 1 Kelvin
units of specific heat capacity
J kg-1 K-1
experiment to calculate specific heat capacity
Continuous flow calorimeter:
1)Set up the experiment with a heating element heating the water and let the water flow at a steady rate until the water out is at a constant temperature.
2)Record the flow rate of the water and duration of the experiment to find the mass of the water as well as the temperature difference (∆θ) between the water in and water out.
3)Record the current and p.d to find the energy transferred to the water (E=IVt+H where H is the heat lost to the surroundings.
4)Repeat the experiment only changing the p.d of the power supply and flow rate (mass) so ∆θ remains constant. You should now have an equation for each experiment.
5)Rearrange equations to get
c=Q₂-Q₁ / (m₂-m₁)∆θ as the H’s cancel.
6) Use Q=IVt to find Q₁ and Q₂ to find the specific heat capacity.
specific latent heat
the energy required to change the state of 1 kg of a substance without raising its temperature
what happens during a change of state
- there is no change in temperature
- kinetic energies of particles are constant
- the potential energies of particles are changing
definition of specific latent heat of vapourisation
the energy required to change the state of 1 kg of a substance from liquid to gas without changing its temperature
definition of specific latent heat of fusion
energy required to change the state of 1 kg of substance from solid to liquid without changing its temperature
Units of specific latent heat
Jkg-1
definition of absolute zero
temperature at which there is zero kinetic energy per particle
absolute zero in degrees celsius
-273 degrees
converting between Kelvin and Celsius temperature scales
Kelvin to Celsius: subtract 273
Celsius to Kelvin: add 273
pressure volume graph
- pressure is proportional to 1/volume so an increased temperature or mass of gas causes the curve to move right and up.
- use pV=nRT to determine changes to graphs
pressure-temperature graph
- temperature is in Kelvin
- pressure is proportional to T
- decreasing the volume or increasing the mass of gas caused a steeper gradient
- use pV=nRT to determine changes to graphs
volume-temperature graph
- temperature is in Kelvin
- V is proportional to T
- decreasing the pressure or increasing the mass of gas caused a steeper gradient
- use pV=nRT to determine changes to graphs
why might energy put into a material not equal the specific heat capacity calculated
energy lost to surroundings means specific heat capacity may not be constant over the temperature range used
define terms N, m and crms in pV = ⅓Nm(crms2).
N - number of particles(molecules)
m - mass of individual particles (molecules)
crms - square root of mean square speed
average kinetic energy per particle of a gas at temperature
=3/2kT = 1/2m(Crms²)
k - boltzman constant - 1.38x10^-23
Total kinetic energy of gas at temperature T
= N x 1.5kT = N x 0.5mCrms². Where N is the number of particles/molecules in gas
k - boltzman constant - 1.38x10^-23
assumptions of ideal gases
- All molecules/atoms are identical
-Molecules/atoms are in random motion = range of speeds, no preferred direction. - Gas contains a large number of molecules
- The volume of gas molecules is negligible compared to the volume occupied by the gas
- No forces act between molecules except during collisions - Collisions are elastic
- Collisions are of negligible duration compared to the time between collisions
what is meant by random motion in context of particles
Particles have no preferred direction, With a range of speeds
what is meant by all collisions being elastic
a collision in which momentum and kinetic energy are conserved
Why is the average velocity of particles in a container zero?
- Velocity is a vector
- Random velocity means no preferred direction of movement so velocities cancel
by considering motion of particles, explain how they exert pressure on the wall of a container
- molecules/particles have momentum
- momentum changes at wall due to collisions lead to a force being exerted on the wall as F=∆mv / t
in terms of kinetic theory, why does pressure decrease when gas is removed from a container
- less gas so fewer molecules
- so rate of change of momentum is less
- pressure is proportional to the number of molecules per unit volume so pressure decreases
In terms of kinetic theory, what effect does reducing temperature of a gas have on the pressure
- Pressure decreases
- Since molecular collisions with wall are less frequent and rate of ∆momentum is less
how does increasing temperature affect the motion of gas particles
increase in temperature causes an increase in the mean kinetic energy so mean square speed increases
quantities that increase when temperature of a constant mass of gas at a constant volume is increased
- pressure
- average KE per particle
- internal energy
Why are mean square speeds of two different gas molecules at the same temperature not the same
- The mean kinetic energy is the same for both gases but the masses of the atoms is different.
- Mean square speed is proportional to 1/mass
- Lighter mass atom/molecules must have a higher mean square speed in order to have the same average kinetic energy.
Root mean square speed
The square root of the mean square speed of the particles in the gas ∴ the average speed.
derivation of ideal gas equation from kinetic theory
Particle, of mass m, travelling with velocity v in a box of length a, width b and height c. Travelling back and forwards along length a and hitting end of area bc.
- Change in momentum of a molecule hitting the end wall ∆p = - 2mv
- from Newton’s 2nd Law : F = ∆p/∆t - Time between collisions of a molecule hitting the end wall ∆t = 2a / v
- Force, F, on the molecule hitting the end wall F= -mv² / a
- from Newton’s 3rd Law : force on molecule is equal and opposite to force on box
- Force, F, on wall from one molecule hitting the end wall; F = mv² / a
- Pressure = Force / Area
- Area of end wall = bc
- Pressure, p, from one molecule hitting end wall; p = mv² / abc
- Volume V = abc;
- Pressure, p, from one molecule hitting end wall; p = mv² / V
- Pressure, p, from N molecules hitting end wall; p = N mv² / V
- Particles have a range of speeds. Average of speeds squared = (crms)²
- Pressure, p, from N molecules with a range of speeds hitting end wall: p = Nm(crms)² / V
- On average in a box, 1/3 of molecules will move in each of the 3 directions.
- Ideal gas equation pV = 1/3 Nm(crms)²
ideal gas equation
pV=nRT for n moles
pV=NkT for N molecules
By combining the 3 gas law equations
(pV/T=constant)
molecular mass
sum of the masses of all the atoms that make up a molecule = mass of one molecule = A
molar mass
the mass of one mole of a gas. One mole of gas contains Avogadro’s number of particles so N=nNa. Volume of one mole of gas = 24dm³
average molecular kinetic energy derivation
pV=nRT=1/3 Nm(Crms)²
∴0.5m(Crms)² = 1.5nRT/N (multiplying by 1.5)
0.5m(Crms)²=1.5kT (Nk=nR)
0.5m(Crms)²=1.5RT/Na (k=R/Na)
internal energy
the sum of kinetic and potential energies. for ideal gases it is the sum of the kinetic energy as they have no potential energy due to no IM forces
kinetic energy of particles
randomly distributed kinetic energy = number of particles x average kinetic energy per particle
potential energy
intermolecular forces increased by: - transfer of thermal energy (heat)
- doing work on a substance
thermal equilibrium
thermal energy will continue to be transferred until objects are in thermal equilibrium
1st law of thermodynamics
change in internal energy = total energy transferred due to work and heat ∆u = ∆Q + ∆W
temperature definition
average energy per particle
boyle’s law pV = constant
- the volume of a fixed mass of gas at constant temperature is inversely proportional to the pressure of the gas
- if the volume of a gas is reduced, particles will be closer together so collisions between themselves and the container will be more frequent
- therefore the pressure increases
charles’ law V/T = constant
- at constant pressure, the volume of a gas is directly proportional to its absolute temperature
- when a gas is heated, its particles gain kinetic energy so constant pressure means the gas particles move quicker and further apart increasing the volume
the pressure law p/T = constant
- at constant volume, the pressure of a gas is directly proportional to its absolute temperature
- as a gas is heated, its particles gain kinetic energy thus if the volume of the container remains the same, gas particles will collide more frequently with themselves and the walls of the container, increasing the pressure of the gas
experiment for Boyle’s law
- put oil in a tube of fixed dimensions which traps a pocket of air
- use a tyre pump to increase the pressure on the oil
- as pressure increases, more oil is pushed up the tube so the air compresses (reduced volume)
- measure the volume of air at room temperature and increase the pressure recording volume and pressure at regular intervals
- a graph of p against 1/v should be a straight line as pV = constant
experiment for Charles’ law
- Seal a capillary tube at the bottom and place a drop of concentrated sulfuric acid halfway up the tube trapping a column of air
- fill a beaker with it inside with near boiling water and the length of the trapped air column should increase. Water cooling will cause it to decrease
- record the temperature and air column length in regular intervals as the water cools and repeat twice more at the same temperature intervals and average your results
- plot a graph of length of column against temperature which will give a straight line showing Charles’ law
gas laws vs kinetic theory
the gas laws are empirical - they are based on observations and evidence of how gases respond to change in their environment and can be proven
kinetic theory is theoretical - it is based on assumptions and derivations from knowledge and theories we already had
brownian motiion
- the random motion of particles suspended in a fluid resulting from their collisions with each other and other atoms
- this can be seen when smoke moves in a zigzag, random motion due to its collisions with smaller lighter air particles travelling at high speeds
- this is evidence that the air is made up of tiny atoms moving quickly hence gained acceptance for the concept of the atom
