Part I (Yep, there’ll be a Part II)

Energy  The capacity to do work or transfer heat  Measured in Joules  Two Types  Kinetic (motion)  Potential (based on position)

1 st Law of Thermodynamics

 aka Law of Conservation of Energy  Energy can be converted from one form to another but cannot be created or destroyed

Heat vs Temperature HeatTemperature  A measure of energy content  What is transferred during a temperature change  Units Joules  Reflects random motion of particles in a substance  Indicates the direction in which heat energy will flow  Units °C or K

Energy, Heat, and Work  ∆E = q + w  Change in energy equals heat plus work  Energy and heat are state functions. Work is not  A state function is independent of the pathway taken to get to that state. (only the beginning and end matter)

Sign Conventions for q, w, and ∆E q+ means system gains heat (endothermic) - means system loses heat (exothermic) w+ means work done on system -means work done by system ∆E+ mean net gain of energy by system -means net loss of energy by system

Work, Pressure, and Volume  w = -P∆V Work done By a gas (through exapansion) To a gas (by compression) ∆V+- w-+

Chemical Energy EndothermicExothermic  Rxns in which energy is absorbed from the surroundings  Energy flows into the system to increase the potential energy of the system  Rxns that give off energy as they progress  Some of the potential energy stored in the chemical bonds is converted to thermal energy (random KE) through heat  Products are more stable (stronger bonds) than reactants

Enthalpy  H = E + PV  H – enthalpy  E – internal energy  P- pressure  V – volume  In systems at constant pressure, where the only work is PV, the change in enthalpy is due only to energy flow as heat  ∆H = heat of rxn

Sample Problem  Indicate the sign of the enthalpy change, ∆H, in each of the following processes carried out under atmospheric pressure, and indicate whether the process is endothermic or exothermic  An ice cube melts  1g of butane is combusted

Enthalpies of Reaction  ∆H = H products - H reactants 1. Enthalpy is an extensive property 2. The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to ∆H for the reverse reaction. 3. The enthalpy change for a reaction depends on the state of the reactants and products

Sample Problem  How much heat is released when 4.50g of methane gas is burned in a constant pressure system? (∆H = -890 kJ)

Calorimetry  Science of measuring heat  Heat Capacity (C)  C= heat absorbed/Temp increase  Specific Heat Capacity  Energy required to raise the temp of 1 gram of a substance 1 °C  Molar Heat Capacity  Energy required to raise the temp of 1 mole of a substance 1 °C

Calorimetry  ∆H = mc∆T or q = mc∆T  m – mass (g)  c – specific heat (J/g·K)  ∆T – change in temperature (K)

Sample Problem  How much heat is needed to warm 250 g of water from 22°C to near its boiling point, 98°C? (The specific heat of water is 4.18 J/g·K)  What is the molar heat capacity?

Constant Pressure Calorimetry Coffee Cup Calorimeter

Sample Problem  When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0°C to 27.5°C. Calculate the enthalpy change for the reaction in kJ/mol HCl, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that its density is 1.0 g/mL, and that its specific heat is 4.18 J/g·K.

Constant Volume Calorimetry Bomb Calorimeter

Hess’s Law  In going from a particular set of reactants to a particular set of products, the change in enthalpy (∆H) is the same whether the reaction takes place in one step or a series of steps

Hess’s Law

Using Hess’s Law 1. Work backward from the final reaction 2. Reverse reactions as needed, being sure to also reverse the sign of ∆H 3. Remember that identical substances found on both sides of the summed equation cancel each other out.

Sample Problem  The enthalpy of reaction for the combustion of C to CO 2 is kJ/mol·C, and the enthalpy for the combustion of CO to CO 2 is kJ/mol·C. Using this data, calculate the enthalpy for the combustion of C to CO.

Standard State  For a compound  Gaseous state  Pressure of 1 atm  Pure liquid or solid  Standard state IS the pure liquid or solid  Substance in a soln  Concentration of 1 M  For an element  The form in which the element exists at 1 atm and 25°C

Standard Enthalpies of Formation (∆H f °)  The change in enthalpy that accompanies the formation of one mole of a compound from its elements with all elements in their standard state

Calculating Enthalpy Change 1. When a rxn is reversed, the magnitude of ∆H remain the same, but its sign changes. 2. When the balanced eqn for a rxn is multiplied by an integer, the value of ∆H must be multiplied by the same integer 3. The change in enthalpy for a rxn can be calculated from the enthalpies of formation of the reactants and products ∆H° rxn =Σ∆H f ° products - Σ∆H f ° reactants 4. Elements in their standard states are not included  For elements in their standard state, ∆H f ° = 0

Sample Problem  For which of the following reactions at 25°C would the enthalpy change represent a standard enthalpy of formation? For those where it does not, what changes would need to be made in the reaction conditions? a) 2Na(s) + ½ O 2 (g)  Na 2 O(s) b) 2K(l) + Cl 2 (g)  2KCl(s) c) C 6 H 12 O 6 (s)  6C (diamond) + 6H 2 (g) + 3O 2 (g)

Sample Problem  Calculate the standard enthalpy change for the combustion of 1 mol of benzene, C 6 H 6.

Sample Problem  Compare the quantity of heat produced by combustion of 1.00g propane, C 3 H 8, to that produced by 1.00 g benzene.