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

kinetic energy- energy of motion

Ek = ½ mv2

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

Electrostatic potential energy

-arises from the interactions between charged particles

-energy is proportional to the electrical charges on the two interacting objects

Eel = kQ1Q2

d

k= proportionality constant= 8.99x109J∙m/C2

C= Coulomb (unit of electrical charge)

Q1 and Q2 = electrical charges (≈ 1.60 x 10-19C)

d= distance (m)

-when Q1 and Q2 have the same sign, the particles repel each other

-Eel is positive and PE decreases

-when Q1 and Q2 have opposite signs, the particles attract each other

-Eel is negative and PE increases

**FIGURE 5.3 page 161**

Units of Energy

Joule (J) SI unit for energy/heat

-use kJ often b/c J is small

1J = 1kg∙m2/s2

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

system- what is being studied

surroundings- everything else but the system

universe- system and surroundings together

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

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/s2

work = m ∙ g ∙ d

*work will be in J b/c kg∙m2/s2

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 = Efinal – Einitial

initial = reactants final = products

+∆E = system has gained energy

-∆E = system has lost energy

Relating ∆E to Heat and Work

∆E = q + w

q= heat w= 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

endothermic- system absorbs heat from surroundings

ex- melting of ice

exothermic- system loses heat to surroundings

ex- burning gasoline

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

Enthalpy (H)

-heat released or absorbed in a reaction at constant pressure

H = E + PV

P= pressure V= volume

∆H = ∆E+ P∆V

*if ∆H is + system has gained heat

*if ∆H is – system has released heat

Enthalpies of Reaction

-also known as the heat of a reaction (∆Hrxn)

-you need the chemical reaction to solve

ex-CH4(g) + 2O2(g) CO2(g) + 2H2O(ℓ) ∆H= -890kJ

-∆H is exothermic

+∆H is endothermic

*if reaction is reversed, the sign on ∆H is reversed

Enthalpy Changes/Molar Heats

fusion s→ℓ +∆H (endo)

solidificationℓ → s -∆H (exo)

vaporization ℓ → g +∆H (endo)

condensationg → ℓ -∆H (exo)

solution s → aq ±∆H

combustion burning -∆H (exo)

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)

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 Tfinal – Tinitial (K)

m = q/C∆T

C = q/m∆T

∆T = q/mC

Constant-Pressure Calorimetry

No heat enters or leaves!page 177 fig 5.18

BOMB

CALORIMETER

*used for

combustion

*pg 179

Constant-Volume Calorimetry

No heat enters or leaves!

qrxn = -Ccal x ∆T

*qrxn is – b/c it is combusting

Ccal = heat capacity of calorimeter

-in kJ/°C

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

Standard enthalpy of formation (∆Hf°)

-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

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