thermodynamics: energy, heat, temperature, and phase changes chapter 16
TRANSCRIPT
Thermodynamics: Energy, Heat, Temperature, and
Phase Changes
Chapter 16
16.1 Energy A. Energy –
“the capacity to do work or cause the flow of heat”
(work = force x distance)
16.1 Energy
1. Kinetic Energy “energy due to motion”
KE = 1/2 mv2
Ex. The rock actually falling on Wiley Coyote.
2. Potential Energy Energy due to position or arrangement
Ex. The rock actually above his head…levitating there
16.1 Energy B. Law of conservation of energy
Energy is not created or destroyed, only transformed Most of the time true except for nuclear reactions
C. Chemical Potential Energy Energy due to chemical bonding
Attractions and repulsions due to ionic and covalent bonding
16.1 Energy D. Heat
“q”- energy transfer between a system and its surroundings caused by difference in temp
Flows High T Low T stops when system and surrounding the same T
K.E. transfer from system to surroundings
What was temperature again???
16.1 Energy Measuring Heat
Temperature is used to monitor the flow of heat in and out of a system
16.1 Energy 1. Units of Heat
calorie (cal) “Quantity of heat that will raise 1.0 g of water 1.0 oC” Based on water (common substance) so it’s easy to
calculate
Joule (J) Unit of energy used to measure all forms of energy, not
just heat SI Unit of heat (our favorite one in Chem) Ex. 60 Watt Light bulb used 60 J/s of energy
Calorie (Cal): typically known as the food calorie 1 kcal = 1000 calories
Conversions 1 calorie = 4.184 joules 1 Calorie = 1000 calories
16.1 Energy
Practice: How many joules of heat are there in 325
calories? 400 Calories?
16.1 Energy
F. Specific Heat 1. Definition
Amount of energy required to raise 1 gram of substance by 1 degree Celsius
2. Units -
16.1 Energy
CgJ
16.1 EnergyTable I: Specific Heats of Common Substances at 298 K (25 ˚C)
Substance Specific heat J/(g˚C)
Water (liquid) 4.184
Water (ice) 2.03
Water (steam) 2.01
Ethanol 2.44
Aluminum 0.987
Granite 0.803
Iron 0.449
Lead 0.129
Silver 0.235
Gold 0.129
3. Applications of specific heat Ex. Pot of water on the stove
How fast does the pot heat up? The water?
Why is water so special?
16.1 Energy
Calculating the amount of heat evolved or absorbed
Endothermic vs. Exothermic Endothermic= energy put INTO the system Exothermic= energy is released FROM the
system
16.1 Energy
q = m x c x ΔT
q = heat absorbed or released in joules or calories
m = mass of the sample in grams c = specific heat of the substance in joules
or calories/g °C ΔT = Tf – Ti change in temperature in
Celsius
16.1 Energy
Practice: If the temperature of 34.4 g of ethanol increases from 25.0 ˚C to 78.8 ˚C, how much heat has been absorbed by the ethanol?
16.1 Energy
A. Measuring Heat using a calorimeter Calorimeter- Device used to measure heat,
based on the law of Conservation of Energy Energy gained by one substance had to be lost by
another
16.2 Heat in Chemical Reactions and Processes
B. Determining the specific heat
heat lost = - heat gained
qlost = - qgained
16.2 Heat in Chemical Reactions and Processes
Coffee Cup Calorimeter
Thermometer
Why a Styrofoam Cup?
1) Good insulator
2) Won’t absorb heat for itself
-Heat lost by hot solids is gained by water in cup
-From mass + temp change of water, you can calculate a quantity of heat H2O
16.2 Heat in Chemical Reactions and Processes
Practice: A piece of metal with a mass of 4.68 grams absorbs 256 J of heat when its temperature increases by 182 ˚C. What is the specific heat of the metal?
16.2 Heat in Chemical Reactions and Processes
Chemical energy and the universe System
Whatever we are studying (usually a chemical reaction in a beaker/vessel)
Surroundings Everything else, such as room conditions, etc.
Universe Contains both system and surrounding
16.2 Heat in Chemical Reactions and Processes
Enthalpy (H) Term that includes heat tranfers (q) and also PV
work that is done by a system We usually simplify that H = q
Enthalpy of a reaction Hreaction = Hproducts - Hreactants
16.2 Heat in Chemical Reactions and Processes
F. Transition State diagrams Exothermic Endothermic
“energy is released” “energy is absorbed”
16.2 Heat in Chemical Reactions and Processes
A. Exothermic reaction
Ex. 4 Fe (s) + 3 O2 (g) 2 Fe2O3 (s) + 1625 kJ
16.3 Thermochemical Equations
B. Endothermic reaction
Ex. NH4NO3 (s) + 27 kJ NH4+1 (aq) + NO3
-(aq)
16.3 Thermochemical Equations
C. Heating Curve / Cooling Curve
16.3 Thermochemical Equations
A. Standard Heat of Formation Amount of energy required to form a compound
directly from its elements in the gaseous phase
16.4 Calculating Enthalpy Change
A. Spontaneous process Reaction that can occur readily on its own based
on the favorable energy changes that happen for both reactants and products Does NOT mean it will occur fast!
B. Enthalpy (H) Energy difference between reactants and
products Exothermic reactions are favorable for Spontaneous
reactions (H = -)
16.5 Reaction Spontaneity
C. Entropy (S) Amount of disorder of atoms and molecules Entropy of the universe is always increasing More disorder is favorable for spontaneous
reactions
D. Gibbs Free Energy (G) Predicts whether a reaction will be spontaneous Negative Gibbs Free Energy = Spontaneous G = H - TS
16.5 Reaction Spontaneity