Energy Part II Conservation of Energy First Law of Thermodynamics Enthalpy of Reaction

Law Of Conservation Of Energy Energy cannot be created or destroyed but can be transformed from one form of energy to another Also known as the first law of thermodynamics

Internal Energy is Conserved 1st Law of Thermodynamics: For an isolated system the internal energy (E) is constant: Δ E = Ef - Ei = 0 Δ E = Eproduct - Ereactant = 0 We can’t measure the internal energy of anything, so we measure the changes in energy E is a state function E = work + heat

Calorimeters A calorimeter is used to measure heat transfer. can be made with a coffee cup and a thermometer. indicates the heat lost by a sample indicates the heat gained by water.

State Function A property whose value depends only on the present state of the system, not on the method or mechanism used to arrive at that state Position is a state function: both train and car travel to the same locations although their paths vary The actual distance traveled does vary with path New York Los Angeles

ΔE= q + w internal energy changes are state functions

Work and Pistons Pressure = force/area. force=P x area Work = P x (area x d) Work = -P×ΔV In expansion, ΔV>0, and is exothermic Work is done by the system in expansion W= force x d

How does work relate to reactions? Work = Force · Distance is most often due to the expansion or contraction of a system due to changing moles of gas. Gases push against the atmospheric pressure , so Psystem = -Patm w = -Patm×ΔV W=work done by system Work = Force X distance

H = E + W; W= work done by system on surrounding Enthalpy is the internal energy of the system plus the work done by the system on surroundings (H) E= is the internal energy and it is the amount of heat absorbed or release by a system plus the work done by the surroundings H = E + W; W= work done by system on surrounding W = -P×ΔV (Isobaric system) ∆H = ∆E -P∆V (1) E= Q + W ; W=work down by surroundings ∆E = ∆Q + ∆PV ∆Q = q ∆E = q + P∆V (2) Plugging (2) in equation (1) ∆H= q + P ∆V -P∆V AH = q Isobaric: Constant pressure ΔP =0

∆H- Change in Enthalpy The difference in enthalpy is the maximum amount of thermal energy derivable from a thermodynamic process in which the pressure is held constant. Most processes involve transfers of thermal energy at a constant atmospheric pressure ∆H= q Such changes can be easily measured and the heat evolved or absorbed is called the change in enthalpy and given the symbol ∆H

The Sign Convention Endothermic systems require energy to be added to the system, thus the q is (+) Exothermic reactions release energy to the surroundings. Their q is (-) Energy changes are measured from the point of view of the system

Entropy Change Entropy (S)-disorder or randomness in a system or universe ∆S =Change in disorder or randomness For a change, as the system proceeds from… order disorder Entropy increases, system looses energy disorder order Entropy decreases, system gains energy

Energy changes States Changes of Matter The heat of fusion is the amount of heat released when 1 gram of liquid freezes (at its freezing point). is the amount of heat needed to melt 1 gram of a solid (at its melting point). for water (at 0 °C) is 80. cal 1 g water

Sublimation Sublimation occurs when particles change directly from solid to a gas. is typical of dry ice, which sublimes at -78 C. takes place in frost-free refrigerators. is used to prepare freeze-dried foods for long-term storage.

Evaporation and Condensation Water evaporates when molecules on the surface gain sufficient energy to form a gas. condenses when gas molecules lose energy and form a liquid.

Boiling At boiling, all the water molecules acquire enough energy to form a gas. bubbles appear throughout the liquid.

Heat of Vaporization The heat of vaporization is the amount of heat absorbed to vaporize 1 g of a liquid to gas at the boiling point. released when 1 g of a gas condenses to liquid at the boiling point. Boiling Point of Water = 100 °C Heat of Vaporization (water) = 540 cal 1 g water

Heating Curve

- Heat + Heat GAS LIQUID SOLID - Heat - Heat + Heat + Heat Deposition Sublimation + Heat GAS LIQUID Condensation SOLID Crystallization - Heat - Heat + Heat + Heat Liquefaction Vaporization

B Concept Question A solid with a high heat of fusion has: a) high melting point and weak attractive forces between particles b) high melting point and large attractive forces between particles c) low melting point and weak attractive forces between particles d) low melting point and large attractive forces between particles B

Concept Question Why is the enthalpy from liquid to gas usually larger than the enthalpy from solid to liquid in a heating curve?

Chemical Potential Energy Chemical bond: net attractive forces that bind atomic nuclei and electrons together Exothermic reactions form stronger bonds in the product than in the reactant and release energy (-E) Reactions release energy to the surroundings. Their q is (-) Endothermic reactions break stronger bonds than they make and require energy (+E) Requires energy to be added to the system, thus the q is (+) Energy changes are measured from the point of view of the system

Energy changes Endothermic HEAT + Reactants Products Exothermic Heat is the most common energy form observed in chemical reactions. Most processes involve transfers of thermal energy at a constant atmospheric pressure ∆H= Q Endothermic HEAT + Reactants Products the q is (+) Exothermic Reactants Products + HEAT q is (-)

Physical / Chemical Changes: A B + Heat Exothermic X + Heat Y Endothermic AH =q (sys) Exothermic q (sys) (-) AH (-) Endothermic q (sys) (+) AH (+)