Unformatted text preview: Energy is...
The ability to do work. Conserved. made of heat and work. a state function. independent of the path, or how you get from point A to B. Work is a force acting over a distance. Heat is energy transferred between objects because of temperature difference. 2 Chapter 6
Energy Thermodynamics 1 The universe
is divided into two halves. the system and the surroundings. The system is the part you are concerned with. The surroundings are the rest. Exothermic reactions release energy to the surroundings. Endothermic reactions absorb energy from the surroundings. 3 4 CH 4 + 2O 2 CO 2 + 2H 2 O + Heat Potential energy CH 4 + 2O 2 Heat CO 2 + 2 H 2 O N 2 + O 2 + heat 2NO Direction
Every energy measurement has three parts. 1. A unit ( Joules of calories). 2. A number how many. 3. and a sign to tell direction. negative  exothermic positivepositive endothermic 2NO Potential energy Heat N2 + O2
6 5 1 Surroundings System Surroundings System Energy Energy E <0
7 8 E >0 Same rules for heat and work
Heat given off is negative. Heat absorbed is positive. Work done by system on surroundings is positive. Work done on system by surroundings is negative. ThermodynamicsThermodynamics The study of energy and the changes it undergoes. 9 10 First Law of Thermodynamics
The energy of the universe is constant. Law of conservation of energy. q = heat w = work E = q + w Take the systems point of view to decide signs. Energy is state function Heat and work are not What is work?
Work is a force acting over a distance. w= F x d P = F/area d = V/area w= (P x area) x (V/area)= PV P Work can be calculated by multiplying pressure by the change in volume at constant pressure. units of liter x atm = Latm L11 12 Work needs a sign
If the volume of a gas increases, the system has done work on the surroundings. work is negative w =  PV Expanding work is negative. Contracting, surroundings do work on the system w is positive. 1 L atm = 101.3 J 2 Examples
What amount of work is done when 15 L of gas is expanded to 25 L at 2.4 atm pressure? If 2.36 J of heat are absorbed by the gas above, what is the change in energy? How much heat would it take to change the gas without changing the internal energy of the gas? 13 14 Enthalpy
abbreviated H H = E + PV (that's the definition) H = E + PV at constant pressure. H = E + PV P the heat at constant pressure qp can be calculated from E = qp + w = qp  PV qp = E + P V = H Calorimetry
Measuring heat. Use a calorimeter. Two kinds Constant pressure calorimeter (called a coffee cup calorimeter) heat capacity for a material, C is calculated C= heat absorbed/ T = H/ T specific heat capacity = C/mass Q = Cm T 15 16 Calorimetry
molar heat capacity = C/moles heat = specific heat x m x T heat = molar heat x moles x T Make the units work and you've done the problem right. A coffee cup calorimeter measures H. An insulated cup, full of water. The specific heat of water is 1 cal/gC Heat of reaction= H = C x mass x T Examples
The specific heat of graphite is 0.71 J/gC. Calculate the energy needed to raise the temperature of 75 kg of graphite from 294 K to 348 K. A 46.2 g sample of copper is heated to 95.4C and then placed in a calorimeter containing 75.0 g of water at 19.6C. The final temperature of both the water and the copper is 21.8C. What is the specific heat of copper? 17 18 Calorimetry
Constant volume calorimeter is called a bomb calorimeter. Material is put in a container with pure oxygen. Wires are used to start the combustion. The container is put into a container of water. The heat capacity of the calorimeter is known and/or tested. Since V = 0, PV = 0, E = q P 3 Bomb Calorimeter
thermometer stirrer full of water ignition wire Steel bomb sample 19 20 Properties
intensive properties not related to the amount of substance. density, specific heat, temperature. Extensive property  does depend on the amount of stuff. Heat capacity, mass, heat from a reaction. Hess's Law
Enthalpy is a state function. It is independent of the path. We can add equations to come up with the desired final product, and add the H Two rules If the reaction is reversed the sign of H is changed If the reaction is multiplied, so is H 22 21 Standard Enthalpy
O2 2NO 112 kJ 180 kJ NO2 68 kJ N2 2O2 23 24 The enthalpy change for a reaction at standard conditions (25C, 1 atm , 1 M solutions) Symbol H When using Hess's Law, work by adding the equations up to make it look like the answer. The other parts will cancel out. H (kJ) 4 Example
Given 5 C 2 H 2 (g) + O 2 (g) 2CO 2 (g) + H 2 O( l ) 2 H= 1300. kJ C(s) + O 2 (g) CO 2 (g) H= 394 kJ 1 H 2 (g) + O 2 (g) H 2 O(l) 2 H= 286 kJ calculate H for this reaction 2C(s) + H 2 (g) C 2 H 2 (g) 25 26 Example
Given O 2 (g) + H 2 (g) 2OH(g) H= +77.9kJ O 2 (g) 2O(g) H= +495 kJ H 2 (g) 2H(g) H= +435.9kJ Calculate H for this reaction O(g) + H(g) OH(g) Standard Enthalpies of Formation
Hess's Law is much more useful if you know lots of reactions. Made a table of standard heats of formation. formation. The amount of heat needed to for 1 mole of a compound from its elements in their standard states. Standard states are 1 atm, 1M and 25C For an element it is 0 There is a table in Appendix 4 (pg A22) 27 28 Standard Enthalpies of Formation
Need to be able to write the equations. What is the equation that would give you the heat of formation of NO2 ? N2 (g) + O2 (g) NO2 (g) Have to make one mole to meet the definition. Write the equation that would give you the heat of formation of methanol, CH3OH. Since we can manipulate the equations
We can use heats of formation to figure out the heat of reaction. Lets do it with this equation. C2H5OH +3O2(g) 2CO2 + 3H2O which leads us to this rule. Since we can manipulate the equations
We can use heats of formation to figure out the heat of reaction. Lets do it with this equation. C2H5OH +3O2(g) 2CO2 + 3H2O which leads us to this rule. ( H o products)  ( H o reactants) = H o f f 29 30 5 What is the enthalpy change for the reaction 2C2H6+ 5O2(g) 4CO(g) + 6H2O(l)? Heats of formation are 84.7kJ/mole for C2H6 110.5 kJ/mole for CO(g) 241.8 kJ/mole for H2O(g) 285.8 kJ/mole for H2O(l).) 31 6 ...
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This note was uploaded on 09/28/2009 for the course CHEM 102 taught by Professor Freeman during the Spring '08 term at South Carolina.
 Spring '08
 FREEMAN

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