Thermodynamics 101Thermodynamics 101  First Law of Thermodynamics  Energy is conserved in a reaction (it cannot be created or destroyed)---sound familiar???  Math representation: Δ E total = Δ E sys + Δ E surr = 0  Δ = “change in”  ΔΕ = positive (+), energy gained by system  ΔΕ = negative (-), energy lost by system  Total energy = sum of the energy of each part in a chemical reaction

Mg+ 2HCl  MgCl 2 + H 2

Exothermic  Temperature increase (--isolated system)  Heat is released to surroundings (--open/closed system)  q = - value  Chemical  Thermal Energy

Endothermic  Temperature decrease (--isolated system)  All energy going into reaction, not into surroundings  Heat absorbed by system, surroundings have to put energy into reaction  q = + value  Thermal  Chemical Energy

Heat of ReactionHeat of Reaction  Amount of heat exchange happening between the system and its surroundings for a chemical reaction.  Temperature remains constant  Usually reactions happen at constant volume or constant pressure

How does work factor into heat of reaction?  W = -P Δ V  If volume is constant ( Δ V), P Δ V = 0 and no other work sooooo  If pressure (P) is constant so volume can change, work is being done soooo

Work in terms of energy change  System DOES work POSITIVE work value for system, system is LOSING energy  System has work on ON it----NEGATIVE work value for system, system is GAINING energy

Enthalpy (H)Enthalpy (H)  Measures 2 things in a chemical reaction: 1)Energy change 2)Amount of work done to or by chemical reaction  2 types of chemical reactions: 1) Exothermic —heat released to the surroundings, getting rid of heat, - ΔΗ 2) Endothermic —heat absorbed from surroundings, bringing heat in, + ΔΗ ** Enthalpy of reaction —heat from a chemical reaction which is given off or absorbed, units = kJ/mol  Enthalpy of reaction  Heat from a chemical reaction which is given off or absorbed  At constant pressure  Units = kJ/mol

Enthalpy (H) cont.Enthalpy (H) cont.  Most chemical reactions happen at constant pressure (atmospheric pressure)—open container  Temperature and pressure are constant  Only work is through pressure/volume  Sum of reaction’s internal energy + pressure/volume of system  H = U + PV  Δ H = Δ U + P Δ V

Properties of EnthalpyProperties of Enthalpy  Extensive Property  Dependent on amount of substance used  State Function  Only deals with current condition  Focus on initial and final states  Enthalpy changes are unique  Each condition has specific enthalpy value SO enthalpy change ( Δ H) also has specific value

Example 1Example 1  CH 4 + 2O 2  CO 2 + 2H 2 O Δ H = kJ

Example 2Example 2  2HgO  2Hg + O 2 Δ H = kJ  HgO  Hg + ½ O 2 Δ H = kJ

More EnthalpyMore Enthalpy  The reverse of a chemical reaction will have an EQUAL but OPPOSITE enthalpy change  HgO  Hg + ½ O 2 Δ H = kJ  Hg + ½ O 2  HgO Δ H = kJ  SOOO-----total Δ H = 0

Example 1:Example 1:  Based on the following:  2Ag 2 S + 2H 2 O  4Ag + 2H 2 S + O 2 Δ H = kJ  Find the Δ H for the reaction below:  Ag + ½ H 2 S + ¼ O 2  ½ Ag 2 S + ½ H 2 O Δ H = ?

Example 2:Example 2:  Write a chemical equation for ice melting at 0°C through heat absorption of 334 kJ per gram.

Stoichiometry ReturnsStoichiometry Returns

Example 1:Example 1:  H 2 + Cl 2  2HCl Δ H = kJ

Example 2:Example 2:  Calculate the Δ H for the following reaction when 12.8 grams of hydrogen gas combine with excess chlorine gas to produce hydrochloric acid.  H 2 + Cl 2  2HCl Δ H = kJ

Example 3:Example 3:  Pentaborane (B 5 H 9 ) burns to produce B 2 O 3 and water vapor. The Δ H for this reaction is kJ/mol at 298°K. What is the Δ H with the consumption of mol B 5 H 9 ?  2B 5 H O 2  5B 2 O 3 + 9H 2 O

Calorimetry  Experimentally “measuring” heat transfer for a chemical reaction or chemical compound  Calorimeter  Instrument used to determine the heat transfer of a chemical reaction  Determines how much energy is in food  Observing temperature change within water around a reaction container ** assume a closed system, isolated container  No matter, no heat/energy lost  Constant volume

Specific Heat CapacitySpecific Heat Capacity  Amount of heat required to increase the temperature of 1g of a chemical substance by 1°C  Units--- J/g  °K  Unique to each chemical substance  Al (s) = 0.901J/g  °K  H 2 O (l) = 4.18 J/g  °K

q = sm Δ T

Example 1Example 1  How much heat is needed to raise the temperature of a 500g iron bar from 25° to 50°C ?

“Coffee Cup” calorimeter“Coffee Cup” calorimeter  Styrofoam cup with known water mass in calorimeter  Assume no heat loss on walls  Initial water temp and then chemical placed inside  Final temperature recorded  Any temperature increase has to be from the heat lost by the substance SOOO  All the heat lost from the chemical reaction or substance is transferred to H 2 O in calorimeter

“Coffee Cup” calorimeter (cont.)  q chemical = -q water

Example 2: Using the following data, determine the metal’s specific heat.  Metal mass = 25.0g Water mass = 20.0g  Temperature of large water sample = 95°C  Initial temperature in calorimeter = 24.5°C  Final temperature in calorimeter = 47.2°C  Specific heat of water = 1.00 cal/g  °C OR J/g  °K (KNOW!!!!)

 Δ q rxn  Heat gained/lost in experiment with calorimeter  Δ H rxn  Heat gained/lost in terms of the balanced chemical equation

Example 3:Example 3:  A 50.0 ml sample of 0.250M HCl and 50.0 ml sample of 0.250M NaOH react in a cofee cup calorimeter. The temperature increases from 19.50°C to 21.21°C. Calculate the Δ H for this reaction.

Homework  pp #25, 27, 33-35