Thermal energy/heat: The non-mechanical transfer of energy between a system and its surroundings due to a difference in temperature. (It is a process! Hence no such thing as “thermal energy in a body” just as there is no “work in a body”)

To increase precision of measurement of temperature use a more sensitive thermometer with a finer graduated scale, or take values at smaller intervals of time.

Thermal Equilibrium: Two objects placed in contact will eventually come to the same temperature; there is no net heat transfer.

Absolute temperature:

Internal energy: The total energy of all the molecules in a substance; sum of PE (intermolecular forces) and KE (random translational and rotational/vibrational motion) of molecules.

Temperature: A measure of the average KE of individual molecules.

Mole: One mole of any substance is equal to the amount of substance that contains the same number of atoms in 0.012kg of carbon-12.

Molar mass: the mass of one mole of a substance

Avogadro constant: The number of atoms in 0.012kg of carbon-12 =

Specific heat capacity c: The energy required to raise a unit mass of a substance by 1 K.

Thermal capacity C: The energy required to raise the temperature of an object by 1 K.

(The amount of heat required to change the temperature of a material is proportional to the mass and to the temperature change)

Two solid objects undergo the same temperature change. A student states the change in internal energy of the two subjects would be the same. Briefly discuss this statement:

Q=mCT

Change in internal energy depends on the mass and specific heat capacity of the two objects as well, hence change in internal energy will be the same if product of mass and specific heat capacity for the two objects is the same as well.

Moving Particle Theory:

- All matter composed of extremely small particles

- All particles are in constant motion

- All collisions are elastic (KE conserved)

- Mutual attractive force exists between particles

Solid:

- Molecules are close together, held in fixed positions by strong IMF – fixed shape and fixed volume, high density

- As temperature increases, vibrations of the molecules in their fixed positions increases

Liquid:

- Molecules are relatively close together, held within the liquid by relatively strong IMF, able to move around each other – fixed volume, takes shape of container, relatively high density

- As temperature increases, vibrations of the molecules increases

Gas:

- Molecules are far apart and move randomly at high speeds, IMF are almost negligible – Gas expands to fill container, very low density

Compare gas and liquid, forces between gas molecules much smaller, gas molecules move randomly at high speeds, while liquid molecules move more slowly and movement restricted to within the liquid.

Phase change: Melting

- As solid is heated, particles vibrate at an increasing rate, KE of the molecules increases. Intermolecular PE remains constant.

- At melting point, particles vibrate with sufficient thermal energy to overcome strong IMF that hold them in fixed positions within solid

- As solid melts more and more particles gain sufficient energy to overcome IMF until phase change is complete

- Heat added does not to go increasing KE of molecules but rather to increase the intermolecular PE of the molecules, breaking the bonds so that the next phase can occur

- Temperature remains constant throughout change in phase

Phase change: Boiling

- As liquid is heated, vibrational, rotational and translational KE of liquid molecules increases. Intermolecular PE remains constant.

- At boiling point, particles gain sufficient energy to overcome IMF holding them within the liquid, escape into gaseous state.

- Process continues until all liquid particles are changed into gas particles

- Heat added does not to go increasing KE of molecules but rather to increase the intermolecular PE of the molecules, breaking the bonds so that the next phase can occur

- Temperature remains constant throughout change in phase

- Further heating results in higher temperature due to increase in KE of the gaseous molecules and larger translational motion

Phase change: Condensation

- Gaseous molecules lose energy through thermal energy transfer

- Temperature of gas molecules is increasingly lowered until reaches boiling point

- Further energy loss from this point on results in reduction in intermolecular potential energy

- Condensation occurs and phase change from gas to liquid

Evaporation: Process that occurs at the surface of a liquid at any temperature

- Molecules in a liquid at any temperature have a range of kinetic energies

- Molecules collide randomly with one another

- Molecules with higher average KE are able to escape from surface of liquid into air

- Overall net loss in average KE of molecules, hence temperature falls

- Thus evaporation results in cooling.

Factors affecting rate of evaporation:

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-

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- Presence of wind

- : Concept of vapour pressure, in a close system molecules will evaporate until vapour pressure is equal to pressure of surrounding medium (setting up an equilibrium). Thus lowering atmospheric pressure has effect of bringing melting and boiling points closer together. In an area of less pressure, evaporation happens faster because there is less exertion on the surface keeping the molecules from launching themselves.

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Maxwell-Boltzmann Distribution:

- Due to molecular collisions, most molecules move at an average speed

- Few molecules have speeds greater or less than this average speed

Latent Heat: The thermal energy required for a material to undergo a phase change, the energy given out during freezing or condensing.

Specific Latent Heat: The amount of energy per unit mass absorbed or released during a phase change.

Specific latent heat of fusion: Heat required to change a unit mass of material from solid to liquid without a change in temperature.

Specific latent heat of vaporization: Heat required to change a unit mass of a material from liquid to vapour without a change in temperature.

Heater Experiment to find Latent Heat of Vaporization: Repeating experiment using different power allows effective elimination of heat loss and calculation of Lv to be more accurate. If heat loss not accounted for, calculated Lv will be higher than actual value, because only a portion of heat supplied used to boil the fluid, rest lost to the surroundings, thus more heat than required needed to boil the liquid.

Finding heat capacity by observing how temperature of a sample changes with time when heated at constant power:

Hence plotting graph of temperature against time:

Coefficient of volume expansion of water =

Kinetic Theory of Ideal Gas:

(when moving particle theory applied to gases)

- Molecules are in rapid, random motion

Kinetic energy is not constant
Momentum is not constant

- Collisions between molecules and between molecules and walls of container are elastic

Force exist between molecules and container (action reaction)

- No intermolecular attractive forces (IMF of repulsion only act during collisions), hence ideal gases cannot be liquefied.

Intermolecular potential energy of the molecules is constant

- Volume of gas molecules negligible compared to volume of the container

- Time of collision much less than time between collisions

- Newton’s laws apply

The average translational KE of molecules in random motion in an ideal gas is directly proportional to the absolute temperature of the gas.

The internal energy of a fixed mass of an ideal gas is dependent only on the temperature of the system/ average KE of the particles (no intermolecular forces hence PE=0).

Therefore if temperature constant, internal energy dependent on pressure and volume only.

Average is specified as not all the molecules have the same speed, and the speed of the molecules changes each time they collide, thus speed of the individual molecules is always changing.

Gas pressure: When molecule collides with wall of container travels in opposite direction; its momentum changes. Newton’s Second Law, therefore walls exerts force on molecule, and meaning molecule has exerted equal and opposite force on the wall. Collisions between wall and molecules over a period of time means there is effectively a constant force on the wall from the gas.

Thus Pressure = Force exerted per unit area

PV = nRT

R = Universal gas constant 8.31 JK-1mol-1

Temperature in Kelvins.

At standard temperature and pressure: T = 273K, P = 1atm, Volume of 1 mole of an ideal gas = 22.4dm3.

Pressure law – Constant volume,

(When temperature increases, molecules have more average KE, experience greater change of momentum when collide with walls, hence force from each molecule is greater. Also, rate of collisions increases. Thus total force on wall increases and pressure increases)

Charles’ Law – Constant pressure,

Higher temperature, faster moving molecules, collide wall with more force, rate of collisions increases. To maintain constant pressure, rate at which collision take place on a unit area must decrease, hence area must increase, so that average force per unit area maintained. Thus gas occupies greater volume.

Boyle’s Law – Constant temperature,

Constant temperature, molecules move at constant average speed, force on wall remains constant. When volume increases, rate of collisions between molecules and wall decreases, hence average total force on wall decreases and pressure decreases.

If a system and its surroundings are at different temperatures, and the system undergoes a process, the energy transferred by non-mechanical means, to or from the system is referred to as thermal energy or heat Q.

For ideal gas heated under two different situations:

At constant pressure, and at constant volume, which has higher thermal capacity? (assuming same amount of heat supplied)

At constant volume, all the thermal energy supplied goes to increasing the internal energy of the gas, increase in temperature will thus be greater. Since C=Q/T, thermal capacity will be less.

When a gas is compressed rapidly by a piston, its temperature rises because the molecules of gas gain energy from the moving piston. (not because of collisions with one another or walls of container, or reduced distance between molecules)

First Law of Thermodynamics:

The increase in internal energy of a closed system is equal to the heat added to the system minus the work done by the system (Law of conservation of energy).

+-U = Change in internal energy of systemIncrease (in temperature)Decrease (in temperature)Q = Heat supplied to systemHeat enters systemHeat leaves systemW = Work done BY gasGas expands (gas does work)Gas contracts (Work done ON gas)

Isothermal Process: Temperature does not change, Q=W

Adiabatic: No heat flow into or out of the system, occurs during rapid expansion or compression of gases, little time for any heat to exit or enter the system. (P/V curve is steeper than that for isothermal process)

Isochoric/Isovolumetric: Volume remains constant,

Isobaric: Pressure remains constant,

For compression of gas at constant temperature: entropy of gas decreases. The gas is at constant temperature and energy is being transferred to the surroundings (to maintain it at constant temp, because as volume decreases temperature decreases, hence energy released to bring temperature up). The same number of moles of gas molecules occupies a reduced volume and therefore the gas becomes more ordered.

Meanwhile the entropy of the surroundings increases!

KNOW HOW TO DRAW GRAPH PROCESSES.

The area under the graph is the work done by the gas for the process. Total area enclosed is the total work done by the gas for that one cycle.

Area beneath the curve is a measure of energy. Useful work done can be estimated by finding area enclosed by the graph.

Question: When Singapore moves for a 3rd to a 1st world country, the entropy of the country and its infrastructure decreases. Explain how this observation is consistent with the second law of thermodynamics: The entropy of the surroundings increases by a greater factor and becomes more disordered. This is because processes such as the building of infrastructure release thermal energy into the surroundings.

Second Law of Thermodynamics:

The overall entropy of the universe is always increasing.

All natural processes increase the entropy of the universe (though not necessarily local entropy).

Heat can flow spontaneously from a hot to a cold object, but it will not flow spontaneously from a cold to a hot object.

Entropy: The measure of disorder of a system.

It is not possible to convert thermal energy fully into work.

When an ice crystal forms from liquid water, the entropy of the water decreases, by reference to 2nd law discuss the entropy change:

The process of freezing results in the release of latent heat. Release of heat increases entropy of surroundings, more than decrease in entropy when water freezes.

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Physics (Notes on Chapters) Notiz am THERMODYNAMICS, erstellt von ibukunadeleye66 am 15/01/2014.