THERMOCHEMISTRY CHAPTER 5

5.1: THE NATURE OF ENERGY Energy: ability to do work WorkWork: energy used to cause an object to move against a force HeatHeat: energy used to raise the temperature of an object 2 basic types of energy: kinetic and potential KE: energy in the form of motion (related to temperature) PE: stored energy due to position relative to other objects Attractions and repulsions between particles

5.1: COULOMB’S LAW Energy Energy is proportional to the electrical charges on the interacting objects Inversely proportional to the distance between them E = kQ 1 Q 2 d Q 1 and Q 2 are charge particles if both are positive or both negative: repel if positive/negative: attract E is the electrostatic potential energy E=0  no attraction  infinite separation (no interactions) E= +  pushing apart, E decreases as particles separate until E = 0 E= -  attracting  E increases when particles are pulled apart

5.1: UNITS OF ENERGY Joule = kgm 2 /s 2 From W=Fd and F=ma 1 calories = 4.184J 1 Calorie = 1kcal (these are food calories) 1Calorie = 4.184kJ

5.1: SYSTEM AND SURROUNDINGS System + Surroundings = UniverseUniverse System: where reaction occurs Surroundings: everything else Closed system: can exchange energy but not matter with the surroundings Piston in diagram Open system: matter and energy exchanged with surroundings Pot of boiling water Isolated system: neither matter nor energy can be exchanged with the surroundings Insulated thermos (not perfect)

5.1: SYSTEM AND SURROUNDINGS 3 components: 1. number: how much energy 2. unit: J, kJ or cal 3. direction of energy flow: Exothermic=negative  heat leaving system Endothermic=positive  heat entering system Always reference the system i.e. take the system point of view!

5.1: TRANSFERRING ENERGY: WORK & HEAT Work: w = Fd Throwing the ball, pushing an object Heat: energy transfer from a hotter object to a colder one Example: burning of natural gas (combustion) Chemical energy stored in the fuel is released System: molecules involved in the reaction Surroundings: everything else Released energy causes the temperature of the system to increase Energy is then transferred from the system to the surroundings

5.2: FIRST LAW OF THERMODYNAMICS Energy is conserved Internal energy: all the energy of the particles Not just the particles, but the nuclei and electrons too Don’t really know the system’s internal energy Use  E  E = E final – E initial Ef > Ei :  E is positive  system gains energy (endothermic) Ef < Ei :  E is negative  system loses energy (exothermic) Remember, take the view of the system:system When energy of system increases the energy of the surroundings decreases

5.2: RELATING  E TO HEAT & WORK A system can gain or lose energy with the surroundings due to heat or worksystem  E = q + w Heat added or work done on a system  internal energy increases

5.2: STATE FUNCTIONS Internal energy (E) is influenced by temperature, pressure and amount of matter Called a state function b/c internal energy only depends on the current state of the system, not the pathway to get therestate function q and w are not state functions Amount of heat released/absorbed or work done depends on what is happening to the system

5.3: ENTHALPY Heat energy (H) H = E + PV P= pressure; V= volume State functions In chemistry, work is really only related to motion caused by changes in volumechemistry When H 2 gas is produced, to maintain constant pressure of the piston, the volume increases, raising the piston  Called P-V work or Pressure-volume work: w = -P  V If volume increases,  V is positive and w is negative--> work is done by the system (P is always positive or zero) If volume decreases,  V is negative and w is positive → work is done on the system by the surroundings

5.3: ENTHALPY Heat energy (H) H = E + PV P= pressure; V= volume State functions In chemistry, work is really only related to motion caused by changes in volume When H2 gas is produced, to maintain constant pressure of the piston, the volume increases, raising the piston  Called P-V work or Pressure-volume work: w = -P  V If volume increases, dV is positive and w is negative--> work is done by the system (P is always positive or zero) If volume decreases, dV is negative and w is positive → work is done on the system by the surroundings

5.3: ENTHALPY At constant pressure a change in enthalpy  H = d(E + PV)  Rewritten:  H =  E + P  V (assuming constant pressure)   E = q + w and w = -P  V  SO all combined:  H = q + w – w  OR  H = q  Long story short: change in enthalpy is heat gained or lost at constant pressure Long story short Also easier to calculate q rather than  H

5.4: ENTHALPIES OF REACTION For a chemical reaction  H = H products – H reactants Heat of reaction aka enthalpy of reaction:  H rxn Thermochemical equation lists the reaction and enthalpy change: 2H 2 (g) + O 2 (g) → 2H 2 O(g)  H = kJ  Negative  H = exothermic reaction (H f <H r : “extra” energy of the reactants is released)extra Enthalpy diagram shows changes in enthalpy

5.4: ENTHALPIES OF REACTION Enthalpy guidelines  Enthalpy is an extensive property  H is proportional to amount of reactant used  If 1 mol of CH 4 and 2 mol O 2 produces 890 kJ of energy, then 2 mol of CH 4 and 4 mol O 2 produces twice as much energy  Enthalpy for a reaction is equal in size and opposite in sign of the reverse reaction Enthalpy CH 4 + O 2 → CO 2 + H 2 O  H = -890kJ CO 2 + H 2 O → CH 4 + O 2  H = +890kJ  Enthalpy for a reaction depends on the states of the reactants and products 2H 2 O(l) → 2H 2 O(g)  H = +88kJ

5.5: CALORIMETRY Calorimetry: measurement of heat flow Heat capacity: amount of heat required to raise the temp of an object by 1K Molar heat capacity: h c for one mole of a substance (C m ) Specific heat capacity: h c for 1g of a substance (C s )  Can vary slightly with temperature  Water C s = 4.184J/gK If a sample absorbs heat → temp increases  q = mC  T

5.5: CONSTANT PRESSURE CALORIMETRY Calorimeter: used to measure enthalpy  “Coffee cup” calorimeter “Coffee cup” calorimeter 2 styrofoam cups with lid Not fully sealed so open to constant pressure of the atmosphere Mix 2 solutions – any change in enthalpy will cause the temp of the water to raise or lower  Heat from the reaction is transferred to the water  Measure initial and final temperature of water to calculate q For dilute solutions, assume C=4.184J/gK  Heat “lost” by the reaction is “gained” by the water q soln = -q rxn

5.5: CONSTANT PRESSURE CALORIMETRY Bomb calorimeter  Constant volume calorimetry Sealed/insulated container Water absorbs heat released by combustion reaction Must know the heat capacity of the calorimeter  q rxn = -C cal  T More accurate than coffee cup calorimeters Heat transferred is really  E not  H, but really no difference

5.6: HESS'S LAW Hess's Law: if a reaction is carried out in a series of steps,  H for the overall reaction equals the sum of the enthalpy changes for the individual steps  Number of steps does not mattersteps  State function Even if paths are different, will get same final resultdifferent Can use a “few” individual steps to determine the enthalpy changes for a huge number of reactions  Useful for reactions that are difficult to measure directly

5.7: ENTHALPIES OF FORMATION Enthalpy of formation: enthalpy change associated with building a compound from its constituent elements (  H f )  Because enthalpy depends on temp, pressure and state, need to define a standard state – standard set of conditions to make comparison easier 1atm and 298K Standard enthalpy of formationStandard enthalpy of formation:  H f for one mole   H f °  List the most stable for of the element (diatomic elements) O 2 is more stable than O so O 2 is used

5.7: CALCULATING ENTHALPIES OF REACTION Add the enthalpy of formation for each part of a reaction to get the enthalpy of the reactionenthalpy of the reaction Reactants: split into constituent elements   H = -  H f Products: add constituent elements  H f values are listed in Appendix C Multiply  H f by any coefficients needed for balancing

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