Chapter 10 Heat 10-1 Temperature and Thermal Equilibrium

TEMPERATURE - a measure of the average Kinetic Energy of the molecules. Greater the temperature – greater the average Kinetic Energy.

Measure Temperature using Thermometers Alcohol filled Mercury filled Temperature Scales 1) Fahrenheit 2) Celsius and 3) Kelvin Scales Celsius – metric unit for temperature Kelvin – SI unit for temperature – used by scientists and engineers (no negative numbers)

Graph of an Ideal Gas – As Temp decreases so does its volume Graph of an Ideal Gas – As Temp decreases so does its volume. At Volume = 0 Temp equals -273 0C = 0 K

Temperature Conversions Celsius to Fahrenheit TF = 9/5 TC + 32.0 Fahrenheit to Celsius TC = 5/9 (TF – 32.0) Kelvin (K) = TC + 273.15 Kelvin temp = Celsius temp + 273.15 Celsius to Kelvin = add 273 Kelvin to Celsius = subtract 273 Temperatures in the Fahrenheit or Celsius scales can have positive, negative values or zero values. Kelvin temperatures are always positive. In 2003, researchers at MIT found new record of lowest temp in Kelvins 450 pK (0.45 nK) http://mashable.com/2013/01/06/absolute-zero/

Question: A healthy person has a temperature of 98. 6 ºF Question: A healthy person has a temperature of 98.6 ºF. What would this reading be on the Celsius scale? Answer = 37 0C Example: Draw a picture representing molecular motion of three identical molecules at these two temperatures

Thermal Equilibrium- When temperatures are equal and there is an even exchange of heat energy Tcan = 11º C Twater = 11º C

Thermal equilibrium – basis of how thermometers work. if the temperature of a substance increases, so does its’ volume. (thermal expansion) One exception – water Most substances contract when colder (loss of heat energy) Expansion Bridges

HW: Read p. 357-364 HW: Practice 10 A p. 363 (1,2,3,5) and Do problems p. 364 1-3, 4 a and c only, and #5 and Concept Development wks due Fri

10-2 Defining Heat Heat – the transfer of energy because of a difference in temperature. Always flows from high kinetic energy to low kinetic energy. (Hot to cold)

Heat content depends on the mass and the type of material

Rate of transfer depends on temperature difference: The greater temperature difference the greater the energy transfer Twater = 20º C Tcan = 15º C Twater = 35º C Tcan = 5º C Where would the greater energy transfer take place and which way would the energy transfer? Ice = 0 ºC Juice = 20 ºC Ice = 0 ºC Juice = 25 ºC B. has a bigger temperature difference and therefore greater energy transfer. Energy transfers from hot to cold: Juice to Ice

of work raised the temperature by 10C. James Prescott Joule Known for the 1st law of Thermodynamics “Conservation of Energy” Determined that heat is a form energy from the simple device below. He found that the same amount of work done always produced the same amount of heat. Heat = work For 1 kg of water, 4186 joule of work raised the temperature by 10C. Sealed device - Stirring adds energy

Heat has Units of Energy = Joules (J) Symbol used for heat (energy) = Q or q The heat transferred to or from an object = the change in it’s internal energy (∆ U) Q = ∆ U

Or in units of calories = 1 cal/g 0C 10-3 Specific Heat Capacity The ability of any material to retain heat energy is called that material’s Heat Capacity. Specific Heat Capacity (Cp) amount of energy needed to raise the temperature of 1 kg of a substance by 10C . Equation: Cp = __Q___ m x ∆ T Cp = specific heat capacity (at constant pressure) Q = heat released or absorbed m = mass of substance ∆T = change in temperature Specific Heat of water = 4.186 J / g 0C Or in units of calories = 1 cal/g 0C

Specific Heat values for some common substances Do Substances that heat up slowly have high or low specific heat values? High

A calorimeter is used to measure the heat released or absorbed Based on the principle that: Heat lost by one substance = the heat gained by another -QA = QB Law of Conservation of Energy – total energy is conserved Calorimeter

Qwater = mwater x Cp water x Twater Qwater = 5442 J Cp Ni = __-Q__ Sample Problem 100.0 g of nickel at 150 °C was placed in 1.00 L of water (1000g) at 25.0 °C. The final temperature of the water was 26.3°C. What is the specific heat of nickel? Heat (Q) gained by the water = the heat (Q) lost by the metal (nickel). +Q = - Q Qwater = mwater x Cp water x Twater Qwater = 5442 J Cp Ni = __-Q__ mmetal x Tmetal -5442/100g/-123.70C = 0.434 J/g0C

How much heat (energy) in Joules would this be? Example: If you do 5,000 Joules of work, how much heat is that in calories? 5,000 J ● ( 1 cal) = (4.186 J) = 1,200 cal Example: How much heat does it take to heat 54g of H2O from 25°C to 65°C ? Q = m • Cp • ∆T = 54g • 1cal/g°C • (65°C - 25°C) = 54g • 1cal/g°C • 40°C = 2,160 cal or 2,200 cal How much heat (energy) in Joules would this be? 2,200 cal (4.186 J) = 9209 J 1 cal

Units of Heat Energy 1 kcal = 4.186 x 103 J Upper case C Calories = Nutritional Science One calorie is the energy it takes to heat one gram of water by one degree Celsius 1 kcal = 4.186 x 103 J 1 calorie = 4.186 J Upper case C 1 Cal = 1000 cal 1 Cal = 4.186 x 103 J

Internal Energy vs. Heat Internal energy (U)- kinetic energy, potential energy, and all other energies in the molecules of a substance. Unit: Joule Heat (Q) is energy in transit Unit: Joule or calorie An object never has “heat” or “work” only internal energy (heat is transferred and work is done)

Water has a very high specific heat value (4186 J /kg x 0C) or 4.186 J / g oC or in units of cal = 1 cal /g oC Water also has a high “Heat of Vaporization” - phase change from liquid to gas (horizontal line D) (2.26 x 106 J/ kg) Heat of vaporization = Amount of energy needed to change phase from liquid to gas)

Question: Why does a steam burn hurt so much? Vaporization (evaporation)– endothermic process (energy absorbing) – body cools off by sweating. Condensation– (exothermic) Question: What has more energy a gas or a liquid? Gas - when steam hits skin condenses back to liquid releasing all the energy it absorbed.

10-4 Energy is transferred 3 ways

1) Radiation –energy transferred as electromagnetic waves. No matter is transferred. the only form of energy transfer that can occur in vacuum Energy from our sun is transferred by radiation.

2) Thermal Conduction heat transfer from molecule to molecule 2) Thermal Conduction heat transfer from molecule to molecule. Thermal conductivity of any material is dependent on two things: i. Motion of free electrons ii. Molecular vibrations Usually occurs between a solid and a solid – but can occur between a solid (earth’s surface) and the air (gas) above it. The rate that energy is transferred as heat depends on the: Density and atomic structure of the material, length, cross sectional area and the temp difference. Substances that transfer thermal energy quickly are called thermal conductors Thermal Insulators conduct energy slowly (trap in the heat) Copper pans

Thermal Conductivities of some Common Substances Conductivity W/(m • K) Diamond 42,907 Copper 6,832 Aluminum 4,205 Brass 1,925 Lead 615 Steel 238 Nickel 138 Concrete 21 Glass 12.55 Water Wood (Oak) 2.51 Wool 0.84 Goose Down 0.42 Air and Most Gases Foam Polystyrene 0.17

3) Convection – the movement of energy (heat) from one place to another. Occurs in liquids and gases Caused by differences in density. Heat rises – cool sinks (more dense)