Thermal

    Cards (41)

    • The internal energy of a body is equal to the sum of all of the kinetic energies and potential energies of all its particles.
    • The kinetic and potential energies of a body are randomly distributed.
    • The internal energy of a system can be increased in two ways: Do work on the system to transfer energy to it, or increase the temperature of the system.
    • When the state of a substance is changed, its internal energy also changes, because the potential energy of the system changes, while the kinetic energy of the system is kept constant.
    • The temperature of water increases up until 100°C, after which the energy gained through heating the water is no longer used to increase the temperature, but is used to break bonds between water molecules so it can change state to water vapour, and the potential energy is increased.
    • The specific heat capacity of a substance is the amount of energy required to increase the temperature of 1 kg of a substance by 1 °C/1 K, without changing its state.
    • The specific latent heat of a substance is the amount of energy required to change the state of 1 kg of material, without changing its temperature.
    • There are two types of specific latent heat: the specific latent heat of fusion (when solid changes to liquid) and specific latent heat of vaporisation (when liquid changes to gas).
    • The kinetic energy of a gas molecule is directly proportional to temperature (in Kelvin).
    • The speed the molecules will be travelling at can be found using pythagoras’ theorem, where u, v, and w are the components of the molecule’s velocity in the x, y and z directions.
    • The mean square speed of the gas molecules can be defined as , and multiplied by N, the number of particles in the gas, to estimate the sum of the molecules’ speeds.
    • The total pressure of a gas is the sum of all the individual pressures caused by each molecule.
    • An ideal gas follows the gas laws perfectly, meaning that there is no other interaction other than perfectly elastic collisions between the gas molecules, which shows that no intermolecular forces act between molecules.
    • The gas laws were discovered by a number of scientists and later explained by the development of the kinetic theory model, which wasn’t accepted at first.
    • The equation F = u2 can be simplified because the volume (V) of a cube is equal to the cube’s volume (V).
    • The internal energy of an ideal gas is equal to the sum of the kinetic energies of all of its particles.
    • Absolute zero (-273°C), also known as 0 K, is the lowest possible temperature, and is the temperature at which particles have no kinetic energy and the volume and pressure of a gas are zero.
    • One mole of a substance is equal to atoms/molecules, so you can convert between the number of moles (n) and the number of molecules (N) by multiplying the number of moles by the Avogadro constant (NA).
    • Charles’ Law states that when pressure is constant, volume is directly proportional to absolute temperature.
    • The Pressure Law states that when volume is constant, pressure is directly proportional to absolute temperature.
    • Boyle’s Law states that when temperature is constant, pressure and volume are inversely proportional.
    • The molar gas constant (R) is 8.31 J mol-1 K-1.
    • The absolute scale of temperature is the kelvin scale, all equations in thermal physics use temperature measured in kelvin (K).
    • The ideal gas equation can be simplified by using the Boltzmann constant (k), Vp = NA N RT.
    • A change of 1 K is equal to a change of 1°C, and to convert between the two you can use the formula: K = C + 273.
    • Molar mass is the mass (in grams) of one mole of a substance and can be found by finding the relative molecular mass, which is (approximately) equal to the sum of the nucleons in a molecule of the substance.
    • The gas laws describe the experimental relationship between pressure, volume, and temperature for a fixed mass of gas.
    • The duration of collisions is negligible in comparison to time between collisions.
    • Boyle’s law states that pressure is inversely proportional to volume at constant temperature, for example, if you increase the volume of a fixed mass of gas, its molecules will move further apart so collisions will be less frequent therefore pressure decreases.
    • The Pressure Law states that pressure is directly proportional to temperature at constant volume, when the temperature of a gas is increased, its molecules gain kinetic energy meaning they will move more quickly, as volume is constant the frequency of collisions between molecules and their container increases and they collide at higher speeds therefore pressure is increased.
    • Work is done on a gas to change its volume when it is at constant pressure, this is usually done through the transfer of thermal energy.
    • The motion of molecules in random, and they experience perfectly elastic collisions.
    • Charles’s law states that volume is directly proportional to temperature at constant pressure, when the temperature of a gas is increased, its molecules gain kinetic energy meaning they will move more quickly, as volume is constant the frequency of collisions between molecules and their container increases and they collide at higher speeds therefore pressure is increased.
    • The motion of the molecules follows Newton’s laws.
    • The value of work done can be calculated using the formula: Work done = pΔV
    • Brownian motion is the random motion of larger particles in a fluid caused by collisions with surrounding particles, and can be observed through looking at smoke particles under a microscope.
    • A simple molecular model can be used to explain each of the gas laws: Boyle’s law, Charles’s law, and the Pressure Law.
    • The kinetic theory model equation relates several features of a fixed mass of gas, including its pressure, volume and mean kinetic energy.
    • Brownian motion contributed to the evidence for the existence of atoms and molecules.
    • The molecules move in straight lines between collisions.