Chemistry Unit 6 Thermodynamics

Unit 6 : Thermodynamics ENDOTHERMIC AND EXOTHERMIC PROCESSES ● In a constant pressure system, the net energy change is the Enthalpy (ΔH°) ○ Enthalpy: the heat (q) exchanged at constant pressure ● Takes energy to break bonds and attractive forces ○ Energetically unfavorable ○ +ΔH° ● Frees energy to form bonds and attractive forces ○ Energetically favorable ○ -ΔH° ● Endothermic processes ○ Energy required to break bonds ＞ Energy released when bonds form ○ +ΔH° ○ Heat flows from the surroundings to the system ○ Work is done on the system (+) ● Exothermic processes ○ Energy required to break bonds ＜ Energy released when bonds form ○ -ΔH° ○ Heat flows from the system to the surroundings ○ Work is done by the system (-) ● 1st Law of Thermodynamics: energy cannot be created nor destroyed ● Surroundings: everything that is external to the system ○ If the temperature of the surroundings increases, the system released heat, and the reaction is exothermic ○ If the temperature of the surroundings decreases, the system absorbed heat, and the reaction is endothermic ENERGY DIAGRAMS ● Endothermic Reaction ○ ○ ○ ○ ○ PE of products is higher than reactants Products are less stable Energetically unfavorable process Ep ＞ Er = energy is drawn from the surroundings into the system Enthalpy (ΔH°) is positive ● Exothermic Reaction ○ PE of products is lower than the reactants ○ The products are more stable ○ Energetically favorable process ○ Ep ＜ Er = energy is released from the system into the surroundings ○ Enthalpy (ΔH°) is negative ● ΔH° = Ep - Er = -kJ/mol HEAT TRANSFER AND THERMAL EQUILIBRIUM ● Temperature: a measure of average kinetic energy ○ In high temperature systems, molecules have a higher kinetic energy ○ In low temperature systems, molecules have a lower kinetic energy ● When a high temperature system comes into contact with a low temperature system, molecules collide ○ The molecules with higher kinetic energy transfer energy to molecules with lower kinetic energy ○ The molecules with higher kinetic energy slow down and the molecules with lower kinetic energy speed up ○ Energy is transferred from the system at the higher temperature to the system at the lower temperature ● Thermal equilibrium occurs when the two systems reach the same temperature and same average kinetic energy HEAT CAPACITY AND CALORIMETRY ● Heat capacity: the amount of heat (J or kJ) required to raise the temperature of a substance by one degree (C or K) ○ Specific Heat Capacity ○ Molar Heat Capacity ● As  Cp,  heat (q) that can be absorbed before  temperature by 1 ℃ or 1 K ● Cp of H2O = 4.186 J/g℃ ● Calorimetry: method used to measure heat changes in an experiment ○ Heat of a reaction can be calculated by the equation q = mcΔT ■ q = (mass)(specific heat capacity)(temperature change) ○ Heat change in a reaction is equal, but opposite in sign to the heat gained in a reaction ■ q(lost) = -q(gained) ENERGY OF PHASE CHANGES ● Moving from a solid, to a liquid, to a gas, requires energy ○ Moves from low PE to high PE ○ Breaks attractions between molecules ○ Melting and vaporization ● Moving from a gas, to a liquid, to a solid, releases energy ○ Moves from high PE to low PE ○ Forms attractions between molecules ○ Freezing and condensation ● Molar enthalpy of fusion (ΔH°fusion) is the energy required to melt a substance ○ Energy released when a substance freezes is equal, but opposite in sign to energy required to melt a substance (-ΔH°fusion) ● Molar enthalpy of vaporization (ΔH°vaporization) is the energy required to vaporize a substance ○ Energy released when a substance condenses is equal, but opposite in sign to energy required to vaporize a substance (-ΔH°vaporization) INTRODUCTION TO ENTHALPY OF REACTION ● Enthalpy change of a reaction ○ gives the amount of heat energy released (for negative values) or absorbed (for positive values) by a chemical reaction at constant pressure ○ Usually represented by ΔHrxn ○ Represents the difference (Δ) in enthalpy (H) between the reactants and products ● ΔH values given always correspond to the molar coefficients for the reaction as written in the balanced chemical equation ● Negative ΔH = exothermic = heat can be treated as a product ● Positive ΔH = endothermic = heat can be treated as a reactant BOND ENTHALPIES ● In a chemical reaction: ○ Bonds may be broken, which requires an input of energy (endothermic) ○ New bonds may be formed, which releases energy (exothermic) ● Bond energy: the energy stored in a bond ○ If negative, represents the energy released when a bond forms ○ If positive, represents the energy required to break a bond ● ΔHrxn = Σ(BE of bonds broken) + Σ(BE of bonds formed) ENTHALPY OF FORMATION ● Standard enthalpy of formation (ΔH°f) = the change in enthalpy when a substance forms ○ Can be used to estimate the overall enthalpy change of a reaction (ΔHrxn) ● “°” means the value is measured in the standard state ○ Gases = 1 atm of pressure ○ Liquids/solids = most common, stable form at 1 atm and 25°C ○ Solutions = 1 M concentration ● For a compound, ΔH°f is the enthalpy change when 1 mole of a compound forms its elements in their standard states ● For a pure element in its standard state ΔH°f = 0 ● ΔHrxn = ΣΔH°f (products) - ΣΔH°f (reactants) HESS’S LAW ● When a chemical reaction can be broken down into a series of smaller steps, the overall enthalpy change of the reaction (ΔHrxn) is the sum of the ΔH values of each smaller step ○ The ΔH is the same whether the reaction takes place in one step or a series of steps ● If a chemical reaction is reverse, the magnitude of ΔHrxn remains the same, but has the opposite sign ● If a chemical reaction is multiplied by a factor, the ΔHrxn is also multiplied by that factor