Thermal Physics Flashcards
Temperature
Measure of the average KE of the particles in a substance (so temp ∝ KE).
Addition of energy to a substance increases the KE of the individual molecules and, therefore, its temperature (difference in heat and temp lies in the transfer bit).
Units: °C, °F, °K
Instrument: thermometer
Construction of a thermometer
Substance must vary linearly with the addition of energy.
Fixed points or reference marks must be present:
Ice point: point of equilibrium between ice and water (0°C, 32°F, 273K)
Steam point: point of equilibrium between steam and water (100°C, 212°F, 373K)
Calibrations must be equally spaced between degrees.
Conversions
Celsius to Kelvin:
K = °C + 273
Celsius to Fahrenheit:
TF = 9/5TC + 32
TC = 5/9(TF - 32)
Heat
Energy transferred between objects because of temp difference.
Always flows from hotter body to colder.
Form of energy (thermal energy).
Thermal equilibrium
State in which two objects in physical contact with each other achieve the same temp.
At this point, there’s no transfer of heat between the two objects.
States of matter
Solid: Molecules tightly packed, vibrate about fixed positions.
Liquid: Loosely packed, glide/slide around one another.
Gas: Far apart, move randomly.
Methods of heat transfer
Conduction, convection, and radiation.
Conduction
Heat transfer by direct contact of objects: occurs mostly in solids because the molecules are tightly packed together.
Sometimes also occurs in liquids.
Conduction involves a medium of transfer (solid molecules).
Ex. transfer of heat from electric stove to pan, spoon in hot coffee.
Convection
Transfer of heat by the movement of air/water molecules: occurs mostly in liquids and gases.
As air/water gets warm, the molecules at the bottom become lighter and rise and get replaced by the colder ones (convection current).
Requires a medium to occur (liquid/gas molecules).
Ex. heating water up, heating up/cooling of the room, land/sea breeze.
Radiation
Transfer of heat by electromagnetic waves through the vacuum of space.
No medium of transfer required (no contact).
Ex. drying clothes under the sun, warming your hands near the fire.
Phase change
Change from one state of matter to another.
Evaporation - Liquid to gas (ex. boiling of water)
Condensation - Gas to liquid (ex. steam room)
Sublimation - Solid to gas (ex. dry ice)
Melting - Solid to liquid (ex. ice cube melting)
Freezing - Liquid to solid (ex. water to ice cube)
Deposition - Gas to solid (ex. water vapor to snow)
Law of Conservation of Energy
States that energy can neither be created nor destroyed, but can be changed from one form to another.
Gas Laws
Explain the behavior of gases:
- Boyle’s Law
- Charles’s Law
- Pressure (Gay-Lussac’s) Law
Deal with Kelvin (Celsius has negative values: we use Kelvin to avoid getting negative volume/pressure values [impossible]).
All about pressure, temp, and volume: for each of these laws, one of these must remain constant (looks at the relationship between the other two).
Boyle’s Law
At constant temp, pressure is inversely proportional to volume.
Pressure is determined by the molecules hitting the walls, so when they’re given more space, there’s less hitting the walls, and so the pressure decreases.
Formula (the factor that’s constant doesn’t play a role in the formula): P ∝ 1/v
PV = constant
P1V1 = P2V2
(main formula)
*See doc for graphs
Charles’s Law
At constant pressure, volume ∝ temp.
Heat gives the molecules more [kinetic] energy, they hit the walls faster, and the volume increases.
Formula: V ∝ T
V/T = constant
V1/T1 = V2/T2
(main formula)
*See doc for graphs
*-273°C is the coldest Celsius temp (hypothetical [hence the dashed line])
*a ∝ b
a∝1/b → inversely proportional to b/directly proportional to 1/b
Pressure (Gay-Lussac’s) Law
At constant vol, temp ∝ pressure.
A higher temperature means greater kinetic energy for the particles (thermal energy is a form of kinetic energy because thermal energy comes from the movement of the particles, and kinetic energy is the energy of movement), which means more molecules hit the walls, which makes for increased pressure.
Formula: P ∝ T
P/T = constant
P1/T1 = P2/T2
(main formula)
- See doc for graphs
- Remember: Constant doesn’t play a role in the formula (also, everything must be in Kelvin, and pressures and volumes must be in consistent units [also, gas laws only apply to the ideal—so, nonexistent–gas, though there are certain things you can do to get real gases close to idal])
Metal rod as a conductor; air as an insulator
A metal rod is a good conductor of heat: molecules are closely packed (easy heat transfer [because they are in contact]).
Air is a very good insulator/poor conductor; it stops the heat from being transferred from the inside to the outside.
Winter clothing is fluffy because of air (same with styrofoam [another great insulator]: the material is made such that there are air molecules trapped within).
Difference: A conductor allows current to flow easily through it–an Insulator doesn’t allow current to flow through it.
Determining heat transfer
Either add the two and divide by two, or subtract the colder body from the hotter body.
Specific heat capacity (c)
The energy required to change the temp of 1kg (unit) of a substance by 1°C or 1K (we can interchange here because: there’s a 10-degree difference between 10°C and 20°C; there’s a 10-degree difference between 283K and 293K, and the each is equal to its counterpart, so it doesn’t matter which scale we use)—hence the ‘specific’ (specific to 1kg).
Every substance has its own specific heat capacity (if the bonds between the atoms are stronger, it will require more energy to heat up the substance). Water has a very high SHC (reason for being a universal cooling agent [if something has a high SHC, it would be good for cooling hot things down because it takes a while for it to get heated up as it absorbs the energy]).
Formula: c = Q/mΔT c = SHC Q = heat/thermal energy m = mass T = temp
Unit: J/kg°C or J/g°C (or J/kgK [your mass and SHC must have the same unit: ex. 4.186 J/g°C and 4196 kJ/g°C?])
Calorimetry
The process by which a calorimeter is used to measure the SHC of an unknown substance.
Usually water, whose SHC is well-known (we need an SHC we know to measure one we don’t), is used to obtain the SHC of the unknown substance.
Method 1: Process involves the dropping of the heated unknown substance into water (at room temp)–the heated substance loses energy while water gains energy.
Method 2: The mixing of hot and cold liquids (in a calorimeter), with either of them having an unknown SHC.
Recall:
c = Q/mΔT
∴ Q = mcΔT
For calorimetry, Qw = Qx mcΔT w = mcΔT x The c x is the SHC we’re solving for; everything else is unknown. *Remember: Larger - Smaller
Specific latent heat (L)
Energy added/removed per 1kg of a substance during a phase change of the substance.
L = Q/m L = SLH Q = heat m = mass of substance Unit: J/kg
*Used when there’s no change in temp (not Q = mcΔT): also, remember to look at the states
There are two types of SLH: SLH of fusion (SLH of a substance during melting or freezing) and SLH of vaporization (SLH during boiling or condensation).
What happens during a phase change?
During a phase change, energy added only goes into changing the bonds: that’s why ex. melting points and boiling points are single values (energy added/removed only affects the PE [and not the KE] of the molecules).
KE is the energy of moving particles (define temperature); PE is the energy involved in the rearrangement of molecules.
For heating curve, temp (°C) and energy (J): -10, 0, 100 (ice [mcΔT], ice + water [melting: mL], water [mcΔT], water + steam [boiling: mL], steam [mcΔT]).
*During the phase changes temp is constant (PE only affected): otherwise, temp increases (KE only affected)
