4. Thermal equilibrium Thermal equilibrium occurs when all parts of the system are at the same temperature. There is no exchange of thermal energy/please do not mention heat. (This is how a thermometer works) Potential energy Potential energy of the molecules arises from the forces (bonds/ because of intermolecular forces) between them. Kinetic energy Kinetic energy of the molecules arises from the translational, rotational, and vibrational motion of the particles. Heat Heat energy is the thermal energy transferred from one body or system of higher temperature to another of lower temperature. Internal energy Internal energy of a system is the sum of the potential and kinetic energy of the particles making up the system. (= thermal energy of the system) Temperature Temperature is a measure of the average kinetic energy of the molecules of the system. If two substances have the same temperatures, their molecules have the same average kinetic energy.

5. Relative atomic mass Relative atomic mass is the mass of an atom in units of 1/12 of the mass of a carbon-12 atom. mole The mole is the amount of substance that contains the same number of atoms/molecules as 0.012 kg of carbon-12. Molar Mass Molar Mass is the mass of one mole of a substance (kg/mol). 1 mole of a gas at STP occupies 22.4 l (dm 3 ) and contains 6.02 x 10 23 molecules/mol.

6. Heat/Thermal Capacity Heat/Thermal Capacity is the amount of thermal energy needed to raise the temperature of a substance/object by one degree Kelvin. Specific heat capacity Specific heat capacity is the quantity of thermal energy required to raise the temperature of one kilogram of a substance by one degree Kelvin. amount of thermal energy needed to increase the temperature of m kg of a substance with specific heat capacity c by ΔT amount of thermal energy released when the temperature of m kg of a substance with specific heat capacity c decreases by ΔT homogeneous substance: C = mc

7. Latent heat Latent heat is the thermal energy that a substance/body absorbs or releases during a phase change at constant temperature. L = Q at const. temp. unit: J Specific latent heat Specific latent heat is the thermal energy required for a unit mass of a substance to undergo a phase change. If electrical energy is converted into increase of internal energy of the system, then: Q added = electrical energy = Pt = IVt = Q abosorbed P – power, I – current, V – voltage, t - time

8. Methods of finding HEAT CAPACITIES and SPECIFIC HEAT CAPACITIES Two methods Two methods – Electrical method – Method of mixtures In all following calculations it is assumed no energy is lost from the system

9. heater (placed in object or substance)A V variable power supply Electrical energy input = thermal energy gained by object: V I t = C  T Heat capacity of an object : Specific heat capacity of a substance: Electrical method Sources of experimental error: ▪ loss of thermal energy from apparatus. ▪ The container for the substance and the heater will also be warmed up. ▪ It will take some time for the energy to be shared uniformly through the substance. Electrical energy input = thermal energy gained by substance: V I t = m c  T

10. Liquid A (m A, T A, unknown c A ) mixed with liquid B (m B, T B, known c B ) A B Final temp T T A > T B Energy lost by hot substance = Energy gained by colder liquid m A c A (T A – T)= m B c B (T – T B ) m A c A (T A – T)= m B c B (T – T B ) Sources of experimental error: ▪ loss of thermal energy from apparatus, particularly while the liquids are being transferred ▪ The changes of temperature of the container have to be taken into consideration for more accurate result Method of mixtures

11. Liquid A (m A, T A, unknown c A ) mixed with liquid B (m B, T B, known c B ) B Final temp T T A > T B Method of mixtures (calorimetar included) Object A with unknown C is heated to a known temperature (usually by immersing in boiling water for a period of time) Then it is transferred to a known mass of liquid in a calorimeter A Energy lost by block = Energy gained by liquid and calorimeter –Thermal capacity of calorimeter is known C A (T A – T)= m B c B (T – T B ) + C c (T – T c ) Calorimeter

12. heater (placed in water) A V variable power supply Electrical method for measuring the specific latent heat of vaporization of water V I t =  m L Sources of experimental error: ▪ loss of thermal energy from apparatus ▪ some water vapour will be lost before and after timing

13. Method of mixture for measuring the specific latent heat of fusion of water Ice (at 0 0 C) is added to warm water and the temperature of the resulting mix is measured

14. 4 Phases (States) of Matter solid, liquid, gas and plasma; ordinary matter – only three phases CharacteristicSolidLiquidGas Volume and shape definite volume and definite shape definite volume but its shape can change – it takes the shape of their containers. neither definite volume nor definite shape Compressibility Almost IncompressibleVery Slightly CompressibleHighly Compressible Bonds = intermolecular forces characterized by high density and the molecules are held in fixed position by strong bonds. Molecules vibrate around a mean (equilibrium) position. density is lower and molecules are further apart without fixed positions Molecules experience little resistance to motion and move freely about. There are still strong forces between the molecules but they are free to move around each other. the forces between molecules are very weak – molecules are essentially independent of one another but they do occasionally collide Comparative Density High Low Kinetic Energy Vibrational Vibrational, rotational, some translational Mostly translational, higher rotational and vibrational Potential Energy HighHigherHighest (ideal gas – zero) Mean molecular Separation r 0 ( size of the particle) > r 0 10 - 100 r 0

15. Phase transition Phase transition is the transformation of a thermodynamic system from one phase to another.

16. Melting point Melting point of a solid is the temperature at which it changes state from solid to liquid. Once at the melting point, any additional heat supplied does not increase the temperature. Instead is used to overcome the forces between the solid molecules, thus increasing potential energy. ◌ At the melting point the solid and liquid phase exist in equilibrium. freezing While freezing, particles lose potential energy until thermal energy of the system is unable to support distance between particles and is overcome by the attraction force between them. Kinetic energy changes form from vibrational, rotational and part translational to merely vibrational. Potential energy decreases (It is negative!!! = attraction: intermolecular forces become stronger). boiling While boiling, substance gains enough potential energy to break free from inter-particle forces. Similar to evaporation, the only difference being that energy is supplied from external source so there is no decrease in temperature. condensing While condensing, the energy changes are opposite to that of boiling. Adding thermal energy to solid increases (vibrational) KE of the molecules, and as temperature is a measure of average KE, temperature increases. Particles will eventually gain enough thermal energy to break from fixed positions. Heat, once absorbed as energy, contributes to the overall internal energy of the object

17. abrupt change The distinguishing characteristic of a phase transition is an abrupt change in one or more physical properties, in particular the heat capacity, and the strength of intermolecular forces. During a phase change, the thermal energy added or released is used to change (increase/decrease) the potential energy of the particles to either overcome or succumb to the inter-molecular force that pulls particles together. The intermolecular forces between molecules are decreasing. The heat is providing enough energy for the molecules to overcome these attractive forces. In the process, the average kinetic energy will not change, so temperature will not change.

18. Evaporation Evaporation is a change of phase from the liquid state to the gaseous state that occurs at a temperature below the boiling point. Evaporation causes cooling. A liquid at a particular temperature has a range of particle energies, so at any instant, a small fraction of the particles will have KE considerably greater than the average value. If these particles are near the surface of the liquid, they will have enough KE to overcome the attractive forces of the neighbouring particles and escape from the liquid as a gas. The escape of the higher-energy particles will lower the average kinetic energy and thus lower the temperature. rate of evaporation The rate of evaporation is the number of molecules escaping the liquid per second. Evaporation can be increased by increasing temperature/more particles have a higher KE Increasing surface area/more particles closer to the surface Increasing air flow above the surface (gives the particles somewhere to go to)/ decreasing the pressure of the air above liquid

19. Kinetic Model of an Ideal Gas Gas pressure is the force gas molecules exert due to their collisions with a wall of container, per unit area. Assumptions of the kinetic model of an ideal gas. PV = nRT = NkT n – number of moles N – number of particles, R – universal gas constant k – Boltzmann constant, T – temperature in Kelvins Gases consist of tiny hard spheres/particles called atoms or molecules. The total number of molecules in any sample of a gas is extremely large. The molecules are in constant random motion. The range of the intermolecular forces is small compared to the average separation of the molecules The size of the particles is relatively small compared with the distance between them No forces act between particles except when they collide, and hence particles move in straight lines. Between collisions the molecules obey Newton’s Laws of motion. Collisions of short duration occur between molecules and the walls of the container and the collisions are perfectly elastic (no loss of kinetic energy).

20. Application of the "Kinetic Molecular Theory" to the Gas Laws Macroscopic behavior of an ideal gas in terms of a molecular model. Kinetic Model of an Ideal Gas

21. low T medium Thigh T Temperaturea measure Temperature is a measure of the average random kinetic energy of the molecules of an ideal gas. Increase in temperature is equivalent of an increase in average kinetic energy (greater average speed). This leads to more collisions and collisions with greater impulse. Thus resulting in higher pressure.  Just as temperature was a measure of the random kinetic energy of molecules for solids and liquids, so it is for an ideal gas. Decrease in volume results in a smaller space for gas particles to move, and thus a greater frequency of collisions. This results in an increase in pressure. Also, depending on the speed at which the volume decreases, particles colliding with the moving container wall may bounce back at greater speeds. This would lead to an increase in average kinetic energy and thus an increase in temperature. An increase in volume would have an opposite effect.

22. Application of the "Kinetic Molecular Theory" to the Gas Laws Macroscopic behavior of an ideal gas in terms of a molecular model. Kinetic Model of an Ideal Gas low T medium Thigh T Temperaturea measure Temperature is a measure of the average random kinetic energy of the molecules of an ideal gas. Increase in temperature is equivalent of an increase in average kinetic energy (greater average speed). This leads to more collisions and collisions with greater impulse. Thus resulting in higher pressure.  Just as temperature was a measure of the random kinetic energy of molecules for solids and liquids, so it is for an ideal gas.

23. Pressure Law (Gay-Lussac’s Law) Effect of a pressure increase at a constant volume Macroscopically: at constant volume the pressure of a gas is proportional to its temperature: PV = nRT → P = (const) T PV = nRT → P = (const) T Isochors Isochors: lines of constant volume (a closed jar, or aerosol can, thrown into a fire will explode due to increase in gas pressure inside). Straight line projected back to 0K (-273 0 C) at zero pressure. It implies that if the gas could be cooled to -273 0 C it would have zero pressure, and at lower temperature a negative pressure, which makes no sense, of course. The concept of absolute zero temperature was first deduced from experiments with gases. Zero pressure would mean still world. n is constant

24. Microscopically: ∎ As T increases, KE of molecules increase ∎ That implies greater change in momentum when they hit the wall of the container ∎ Thus microscopic force from each molecule on the wall will be greater ∎ As the molecules are moving faster on average they will hit the wall more often ∎ The total force will increase, therefore the pressure will increase low T medium Thigh T low p medium phigh p Pressure Law (Gay-Lussac’s Law) Effect of a pressure increase at a constant volume n is constant

25. The Charles’s law Effect of a volume increase at a constant pressure Macroscopically: at constant pressure, volume of a gas is proportional to its temperature: PV = nRT → V = (const) T PV = nRT → V = (const) T Isobars Isobars: lines of constant pressure V vs. T relationship is linear, with an intercept of absolute zero on the Kelvin scale. The slope depends on the pressure. The dashed lines on the graph are to represent the fact that you can not achieve absolute zero, because the volume would attain its lowest possible value – ZERO n is constant A football inflated inside and then taken outdoors on a winter day shrinks slightly. A slightly underinflated rubber life raft left in bright sunlight swells up The plunger on a turkey syringe thermometer pops out when the turkey is done

26. Effect of a volume increase at a constant pressure Microscopically: ∎ An increase in temperature means an increase in the average kinetic energy of the gas molecules, thus an increase in speed ∎ There will be more collisions per unit time, furthermore, the momentum of each collision increases (molecules strike the wall harder) ∎ Therefore, there would be an increase in pressure ∎ If we allow the volume to change to maintain constant pressure, the volume will increase with increasing temperature small V high p medium V medium p large V low p

27. Boyle-Marriott’s Law Effect of a pressure decrease at a constant temperature Macroscopically: at constant temperature the pressure of a gas is inversely proportional to its volume: PV = nRT → P = (const)/V PV = nRT → P = (const)/V isotherm Each line on the graph is a line of constant temperature, and is called an isotherm (meaning literally –constant temperature). In a plot of P versus 1/V, we see that the isotherms are straight lines with constant slopes. Doubling the pressure on a gas halves its volume, as long as the temperature of the gas and the amount of gas aren't changed. The bubbles exhaled by a scuba diver grow as the approach the surface of the ocean. (The pressure exerted by the weight of the water decreases with depth, so the volume of the bubbles increases as they rise.) Deep sea fish die when brought to the surface. (The pressure decreases as the fish are brought to the surface, so the volume of gases in their bodies increases, and pops bladders, cells, and membranes). Pushing in the plunger of a plugged-up syringe decreases the volume of air trapped under the plunger.

28. Boyle-Marriott’s Law Effect of a pressure decrease at a constant temperature Microscopically: ∎ Constant T means that the average KE of the gas molecules remains constant ∎ This means that the average speed of the molecules, v, remains unchanged ∎ If the average speed remains unchanged, but the volume increases, this means that there will be fewer collisions with the container walls over a given time ∎ Therefore, the pressure will decrease