1. Chapter 5 Thermochemistry -relationship between chemical reactions and energy changes energy- capacity to do work or transfer heat work- energy used to cause an object to move against a force heat- the energy used to cause the temp of an object to increase

2. kinetic energy- energy of motion E k = ½ mv 2 m= mass (kg)v= velocity (m/s) ex- Which has more KE? -car moving at 55mph -tractor trailer moving at 55mph potential energy- stored energy, energy of position -as PE energy increases, KE decreases

3. Electrostatic potential energy -arises from the interactions between charged particles -energy is proportional to the electrical charges on the two interacting objects E el = kQ 1 Q 2 d k= proportionality constant= 8.99x10 9 J∙m/C 2 C= Coulomb (unit of electrical charge) Q 1 and Q 2 = electrical charges (≈ 1.60 x 10 -19 C) d= distance (m)

4. -when Q 1 and Q 2 have the same sign, the particles repel each other -E el is positive and PE decreases -when Q 1 and Q 2 have opposite signs, the particles attract each other -E el is negative and PE increases **FIGURE 5.3 page 161**

5. Units of Energy Joule (J)  SI unit for energy/heat -use kJ often b/c J is small 1J = 1kg∙m 2 /s 2 calorie (cal)  amount of energy needed to raise the temp of 1g of water 1°C 1cal = 4.184J 0.2890cal = 1J Calorie (Cal) = 1000 cal

6. system- what is being studied surroundings- everything else but the system universe- system and surroundings together

7. Systems may be: open- matter and energy can be exchanged with surroundings ex- boiling pot of water with no lid closed- can exchange energy but not matter with surroundings ex- page 162 fig 5.4 isolated- energy or matter cannot be exchanged with surroundings ex- thermos

8. Transferring Energy work- causing the motion of an object against a force heat- causing a temp change force- any push or pull on an object work (w) = F · d F = m ∙ g g = force of gravity = 9.8m/s 2 work = m ∙ g ∙ d *work will be in J b/c kg∙m 2 /s 2

9. First Law of Thermodynamics -energy is conserved Internal Energy (E) -sum of all the KE and PE of the components of a system -concerned with the change in energy (∆E) ∆E = E final – E initial initial = reactantsfinal = products +∆E = system has gained energy -∆E = system has lost energy

10. Relating ∆E to Heat and Work ∆E = q + w q= heatw= work done For q:+ if system gains heat - if system loses heat For w:+ if work done on system - if work done by system For ∆E:+ if net gain of energy by system - if net loss of energy by system *if volume is constant, then w= 0 b/c w= -P∆V

11. endothermic- system absorbs heat from surroundings ex- melting of ice exothermic- system loses heat to surroundings ex- burning gasoline

12. State functions- properties that are determined by the state of the system, regardless of how that condition was achieved. Potential energy of hiker 1 and hiker 2 is the same even though they took different paths. ex- energy, pressure, volume, temperature

13. Enthalpy (H) -heat released or absorbed in a reaction at constant pressure H = E + PV P= pressureV= volume ∆H = ∆E+ P∆V *if ∆H is + system has gained heat *if ∆H is – system has released heat

14. Enthalpies of Reaction -also known as the heat of a reaction (∆H rxn ) -you need the chemical reaction to solve ex- CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(ℓ) ∆H= -890kJ -∆H is exothermic +∆H is endothermic *if reaction is reversed, the sign on ∆H is reversed

15. Enthalpy Changes/Molar Heats fusion s→ℓ+∆H (endo) solidificationℓ → s-∆H(exo) vaporizationℓ → g+∆H(endo) condensationg → ℓ -∆H(exo) solutions → aq ±∆H combustionburning -∆H(exo)

16. Calorimetry -measurement of heat flow calorimeter- used to measure heat flow -insulated container heat capacity- amount of heat required to raise the temp of an object 1°C (J/°C or J/K) specific heat (C) - amount of heat needed to raise the temp of 1g of a substance 1°C (J/g·°C or J/g·K) molar heat capacity- heat capacity of one mole of a substance (J/mol·°C or J/mol·K)

17. Specific heat of water = 4.18 J/g∙K q = (m)(C)(ΔT) q = heat (J or cal) m = mass (g) C = specific heat capacity (J/g∙K) *will be given to you if not solving for ΔT = temp change T final – T initial (K) m = q/C∆T C = q/m∆T ∆T = q/mC

18. Constant-Pressure Calorimetry No heat enters or leaves! page 177 fig 5.18

19. BOMB CALORIMETER *used for combustion *pg 179 Constant-Volume Calorimetry No heat enters or leaves!

20. q rxn = -C cal x ∆T *q rxn is – b/c it is combusting C cal = heat capacity of calorimeter -in kJ/°C

21. Hess’s Law -if a reaction is carried out in a series of steps, ∆H of the overall reaction equals the sum of the enthalpy changes for each step *∆H is a state function so it will be the same whether the reaction takes place in one step or a series of steps

22. Standard enthalpy of formation (∆H f °) -change in enthalpy for the reaction that forms one mole of the compound from its elements with all substances in their standard states -in kJ/mol page 184 *Standard states = 101.3kPa, 298K

23. Calculating Enthalpies of Reaction ∆H° rxn = ∑n∆H° f (products) - ∑n∆H° f (reactants) *n = moles from balanced equation *values found in App. C page 1059 or page 184