Fuels •  In 2007 the United States consumed 1.06 × 1020 J of fuel. •  Most from petroleum and natural gas. •  Remainder from coal, nuclear, and hydroelectric. •  Fossil fuels are not renewable.

• Fuel value = energy released when 1 g of substance is burned. • Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g.

Foods • Fuel value = energy released when 1 g of substance is burned. • 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal. • Energy in our bodies comes from carbohydrates and fats (mostly). • Intestines: carbohydrates converted into glucose: C6H12O6 + 6O2 → 6CO2 + 6H2O, ΔH = -2816 kJ • Fats break down as follows: 2C57H110O6 + 163O2 → 114CO2 + 110H2O, ΔH = -75,520 kJ • Fats: contain more energy; are not water soluble, so are good for energy storage.

More About Fuels - Section 6.9

100,000 kWh/person/year in the US

Kinetic Energy and Potential Energy •  Kinetic energy is the energy of motion:

•  KE = (1/2) mv2

•  Potential energy is the energy an object possesses by virtue of its position. •  Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill. Units of Energy •  SI Unit for energy is the joule, J:

KE = (1/2)mv2 = (1/2) (2kg) (1 m/s)2 = 1kg m 2 s-2 = 1J

We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal

Systems and Surroundings •  System: part of the universe we are interested in. •  Surroundings: the rest of the universe.

Transferring Energy: Work and Heat • Force is a push or pull on an object. • Work is the product of force applied to an object over a distance: w = F x d • Energy is the work done to move an object against a force. • Heat is the transfer of energy between two objects. • Energy is the capacity to do work or transfer heat.

Relating ΔE to Heat and Work •  Energy cannot be created or destroyed. •  Energy of (system + surroundings) is constant. •  Any energy transferred from a system must be transferred to the surroundings (and vice versa). •  From the first law of thermodynamics:

when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system: ΔE = q + w

Exothermic and Endothermic Processes - Section 6.5 • Endothermic: absorbs heat from the surroundings. • Exothermic: transfers heat to the surroundings. • An endothermic reaction feels cold. • An exothermic reaction feels hot.

State Functions • State function: depends only on the initial and final states of system, not on how the internal energy is used.

•  Chemical reactions can absorb or release heat. •  However, they also have the ability to do work. •  For example, when a gas is produced, then the gas produced can be used to push a piston, thus doing

work. Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)

•  The work performed by the above reaction is called pressure-volume work. •  When the pressure is constant, •  Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.

H = E + pV •  Enthalpy is a state function. •  If the process occurs at constant pressure,

•  Then ΔΗ = qp

ΔΗ = Δ(E + pV)

= ΔΕ + pΔV

•  Since we know that w = -pΔV •  We can write: ΔH = ΔE + PΔV

= qp + w - w = qp •  When ΔH, is positive, the system gains heat from the surroundings. •  When ΔH, is negative, the surroundings gain heat from the system.

•  For a reaction: ΔH = Hfinal - Hinitial = Hproducts - Hreactants

•  Enthalpy is an extensive property (magnitude ΔH is directly proportional to amount):

•  CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔH = -802 kJ •  2CH4(g) + 4O2(g) → 2CO2(g) + 4H2O(g) ΔH = -1604 kJ

•  When we reverse a reaction, we change the sign of ΔH: CO2(g) + 2H2O(g) → CH4(g) + 2O2(g) ΔH = +802 kJ

•  Change in enthalpy depends on state: H2O(g) → H2O(l) ΔH = -88 kJ

•  For problem-solving, heat evolved (exothermic reaction) can be thought of as a product. Heat absorbed (endothermic reaction) can be thought of as a reactant.

•  We can generate conversion factors involving ΔH. •  For example, the reaction:

H2(g) + Cl2(g) 2 HCl(g) ΔH = –184.6 kJ can be used to write:

–184.6 kJ ———— 1 mol H2

–184.6 kJ ———— 1 mol Cl2

–184.6 kJ ———— 2 mol HCl

What is the enthalpy change associated with the formation of 5.67 mol HCl(g) in this reaction?

Heat Capacity and Specific Heat •  Calorimetry = measurement of heat flow. •  Calorimeter = apparatus that measures heat flow. •  Heat capacity = the amount of energy required to raise the temperature of an object (by one degree). •  Molar heat capacity = heat capacity of 1 mol of a substance. •  Specific heat = specific heat capacity = heat capacity of 1 g of a substance.

Constant Pressure Calorimetry •  Atmospheric pressure is constant!

• Place an empty iron pot weighing 5 lb on the burner of a stove. • Place an iron pot weighing 1 lb and containing 4 lb water on a second identical burner (same total mass). • Turn on both burners. Wait five minutes. • Which pot handle can you grab with your bare hand? • Iron has a lower specific heat than does water. It takes less heat to “warm up” iron than it does water.

ΔHrxn = qrxn / moles

•  Reaction carried out under constant volume. •  Use a bomb calorimeter. •  Usually study combustion.

Bomb Calorimetry (Constant Volume Calorimetry)

•  Hess’s law: if a reaction is carried out in a number of steps, ΔH for the overall reaction is the sum of ΔH for each individual step.

•  For example: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) ΔH = -802 kJ

2H2O(g) → 2H2O(l) ΔH = -88.0 kJ CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = -890 kJ

Note that: ΔH1 = ΔH2 + ΔH3