Thermal Properties of Matter

• Thermal properties of matter are being discussed.

Zeroth Law of Thermodynamics

• If two systems are each in thermal equilibrium with a third system, then they are in thermal equilibrium with each other.
• Thermal equilibrium means there is no spontaneous net heat flow in either direction.
• This law is the basis for temperature measurements.

Thermal Energy, Temperature & Heat

• Temperature is a measure of the average kinetic energy of individual particles in a substance.
• Atoms in an object are in constant motion.
• When a horseshoe is hot, the particles move quickly.
• When a horseshoe has cooled, its particles move more slowly.

Temperature

• Temperature can be measured with a thermometer.
• Temperature can be measured in Kelvin, Celsius, and Fahrenheit.
• Absolute zero temperature is the point at which particles stop moving and equals 0°K.
• Air molecules collide with a thermometer transferring energy inside.
• Cold air molecules move slower resulting mercury not rising.
• Warm air molecules move rapidly causing mercury to rise.

Temperature Conversions

• Fahrenheit to Celsius conversion: F = 1.8C + 32
• Celsius to Fahrenheit conversion: C = (F-32)/1.8

SI Unit for Temperature

• The standard unit of measurement is Kelvin.
• Kelvin is Celsius + 273, 10C = 283K.
• Celsius is Kelvin - 273, 10K = -263C.

Temperature and Kinetic Energy

• Temperature relates to the random motion of molecules.
• In the simplest case of an ideal gas, temperature is proportional to the average kinetic energy of molecular translational motion.
• Temperature is more complicated in solids and liquids where molecules are more constrained and have potential energy.
• The warmth felt when touching a hot surface is the kinetic energy transferred by molecules in the surface to molecules in your fingers.
• Temperature is not a measure of the total kinetic energy of all the molecules in a substance.
• Two liters of boiling water have twice as much kinetic energy as one liter.
• The temperatures are the same because the average kinetic energy of molecules in each case is the same.
• There is more molecular kinetic energy in a bucketful of warm water than in a small cupful of higher-temperature water.

Thermal Energy Relationships

• Thermal energy depends on temperature, mass, and type of substance.
• As temperature increases, so does thermal energy (because the kinetic energy of the particles increases).
• Even if the temperature doesn't change, thermal energy is higher in a more massive substance because it is a total measure of energy.
• A beaker with a larger mass of water has more thermal energy even when both have the same temperature.

Heat

• Heat flows from higher-temperature substances into lower-temperature substances when two substances of different temperatures are in contact.
• Heat is the flow of thermal energy from one object to another.
• Heat always flows from warmer to cooler objects.
• The direction of spontaneous energy transfer is always from a warmer to a cooler substance.
• Heat refers to the energy that transfers from one object to another because of a temperature difference between objects.
• A large bowl of warm water has more thermal energy than a red-hot thumbtack because the water has more total molecular kinetic energy.
• If a tack is immersed in water, heat flows from the hot tack to the cooler water.
• Heat flows based on temperature differences, which means average molecular kinetic energy differences.
• Heat does not flow on its own from a cooler substance into a hotter substance.

Thermal Equilibrium

• When a thermometer is in contact with a substance, heat flows between them until they have the same temperature.
• Objects in contact that reach the same temperature are in thermal equilibrium.
• When objects are in thermal equilibrium, no heat flows between them.
• Water seeks a common level with pressures at equal elevations the same.
• A thermometer and its surroundings reach a common temperature with the average kinetic energy per particle the same.
• If a flame is used to add heat to 1 liter of water, and the water temperature rises by 2°C, then adding the same quantity of heat to 2 liters of water, the temperature will rise by 1°C.

Measuring Heat

• The amount of heat transferred can be determined by measuring the temperature change of a known mass of a substance that absorbs the heat.
• Heat is energy transferred from one substance to another by a temperature difference.
• When a substance absorbs heat, the resulting temperature change depends on more than just the mass of the substance.
• To quantify heat, the mass and kind of substance affected must be specified.
• Although the same quantity of heat is added to both containers, the temperature of the container with less water increases more.
• A unit of heat is defined as the heat necessary to produce a standard temperature change for a specified mass of material.
• The calorie is a commonly used unit for heat.
• A calorie is defined as the amount of heat required to raise the temperature of 1 gram of water by 1°C.

Specific Heat Capacity

• The capacity of a substance to store heat depends on its chemical composition.
• Different substances behave differently when subjected to different temperature changes.
• Food like boiled onions stay hot, and apple pie fillings burn tongues because of a high heat capacity.
• An aluminum covering can be peeled with bare fingers from a hot oven because of a low specific heart capacity,
• An aluminum pan has a relatively low heat capacity, whereas the food has a high heat capacity.
• Different substances have different capacities for storing internal energy, or heat.
• Water on a stove might take 15 minutes to boil.
• The same mass of iron might only take 2 minutes.
• For silver, it may take less than 1 minute.
• A material requires a specific amount of heat to raise the temperature of a given mass a specified number of degrees.
• The specific heat capacity of a material is the quantity of heat required to raise the temperature of 1 gram by 1 degree.
• A gram of water requires 4.186 J/g°C to raise the temperature 1°C.
• Aluminum has a specific heat capacity of 0.900J/g°C.
• Clay has a specific heat capacity of 1.4 J/g°C.
• Copper has a specific heat capacity of 0.386 J/g°C.
• Lead has a specific heat capacity of 0.128 J/g°C.
• Olive oil has a specific heat capacity of 1.97 J/g°C.
• Silver has a specific heat capacity of 0.23 J/g°C.
• Steel has a specific heat capacity of 0.448 J/g°C.
• Specific heat capacity is like thermal inertia signifying the resistance of a substance to change in its temperature.
• A gram of water requires 1 calorie of energy to raise the temperature 1°C.
• It takes about one eighth as much energy to raise the temperature of a gram of iron by the same amount.
• Water has a higher specific heat capacity than sand.
• Water is much slower to warm and cool.

Specific Heat

• Land heats up and cools down faster than water.
• Specific heat is the amount of heat required to raise the temperature of 1 kg of a material by one degree (C or K).
• C water = 4184 J / kg C.
• C sand = 664 J / kg C.
• Land heats up and cools quickly because of lower specific heat capacity.
• Water has a high specific heat.
• Water molecules form strong bonds that require more heat energy to break.
• Metals have weak bonds and require less energy.

Heat Transfer

• The specific heat of aluminum is 920 J/(kg·K).
• The specific heat of water is 4184 J/(kg·K).
• Aluminum has a higher specific heat than copper and will take longer to cool down.
• The energy given off by one system and the energy received by the system are equal by the Law of Conservation of Thermal Energy.
• Heat is denoted by the letter Q.
• Energy is constantly in motion whenever there is a temperature difference between interacting bodies.
• Transfer of energy will stop once the temperature of the objects attains equilibrium.
• Heat Q ∝ ΔT means an increase in temperature needs an increase in energy.
• Heat Q ∝ m ΔT means the amount of heat needed to raise the temperature of any object is proportional to the mass of the object.
• The greater the mass of the object is, the greater the amount of heat needed to raise its temperature.
• Q = mcΔT , where c is the specific heat of the substance and m is the mass of the substance.
• Q = m × ΔT × c where heat is measured in Joules, mass in kilograms, temperature in Kelvin/Celcius and specific heat in J/kg·K
• ∆T = Tf - Tᵢ, -Q = heat loss and + Q = heat gain.
• Adding or extraction of heat from a body may raise or lower temperature without changing physical state.
• The amount of heat an object receives d depends on mass, change of temperature and a coefficient called specific heat capacity, c.
• ΔΤ = Thigh - Tlow and Q = mcΔT.

Heat Transfer Sample Problems

• A 200 g steel ball at 20.0°C is placed in boiling water. The thermal energy absorbed by the ball is calculated as shown:
  • m = 200 g = 0.200 kg
  • Tᵢ = 20.0°C
  • Tf = 100°C
  • Csteel = 0.108 kcal/kg °C
  • Q = mcΔT = (0.200 kg) (0.108 kcal/kg °C) (100°C – 20.0°C) = 1.73 kcal
• A hot 200 g steel ball at 60.0°C is cooled to −4.00°C. How much energy is lost by the ball?
• Q = mcΔT = (0.200 kg) (0.108 kcal/kg °C) (60.0°C – (-4.00°C)) = 1.38 kcal
• A 32-g silver spoon cools from 60°C to 20°C. How much heat is lost by the spoon?
  • Q = mcΔT, m = 32 g, ΔT = 20°C - 60°C = – 40°C, and csilver = 0.23 J/g°C
  • Q = (32 g) (-40°C) (0.23 J/g°C) = -290 J
• Heat is required to warm 230 g of water from 12.0°C to 90.0°C?
  • Q = mcAT, m = 230 g, AT = 90°C - 12°C = 78°C and c = 4.186 J/g°C
• Q = (230 g) (78°C) (4.186 J/g°C) = 75,100 J