Thermochemistry

Thermochemistry Energy Transformations

Definitions Thermochemistry – • The study of energy changes that occur during chemical reactions and changes in state. • Energy changes occur as either heat transfer or work, or a combination of both. Heat – • Represented by q • The change in heat is represented by ∆H • The energy that transfers from one object to another because of a temperature difference between them. • Always flows from a warmer object to a cooler object. System – • A part of the universe on which you focus your attention Surroundings – • In thermochem, the region in the immediate vicinity of the system

Endothermic Reaction Endothermic reactions – chemical reaction that absorbs energy to break existing bonds • Heat goes into the reaction (system) from the surroundings • This movement of heat is defined as positive; q has a positive value • The surroundings will feel colder • Temperature of endothermic reactions goes down • Ex: Boiling Water or Melting Ice (absorbing energy)

Endothermic Reaction

Exothermic Reactions • Exothermic reactions – chemical reaction in which energy is released • Heat goes out of the reaction (system) into the surroundings • This movement of heat is defined as negative; q has a negative value • The surroundings will feel hotter • Temperature of exothermic reactions goes up • EX: Condensation (gas to liquid) releasing energy!

Exothermic Reaction

Surroundings System Heat qsys > 0 Endo- and Exothermic Surroundings System Heat qsys < 0 ENDOTHERMIC EXOTHERMIC

Endothermic & Exothermic Reactions • Are the following reactions endothermic or exothermic? • CO + 3H2 CH4 + H2O H= -206kJ • Exothermic (H is negative) • I add magnesium metal to some hydrochloric acid. The temperature goes from 23C to 27 C • Exothermic – temperature goes up • I mix together some vinegar & baking soda. The temperature goes from 28C to 23C • Endothermic – temperature goes down

Chemistry Happens in Moles • An equation that includes energy is called a thermochemical equation CH4 + 2 O2→ CO2 + 2H2O ΔH = -802.2 kJ • 1 mole of CH4 releases 802.2 kJ of energy STEPS: Convert to moles and then multiply by the energy given over the number of moles of the compound! • If 10.3 g of CH4 are burned completely, how much heat will be produced?

The Work CH4+ 2 O2→ CO2 + 2H2O ΔH = -802.2 kJ 10.3 g CH4 515 kJ 1 mol CH4 802.2 kJ 16.0425 g CH4 1 mol CH4

Thermochemical equations S + O2 SO2H = – 296.9 kJ • If we change the equation, then the H also changes … SO2 S + O2H = + 296.9 kJ • If the reaction is reversed the sign is reversed • Also, if numbers in the equation change, so will the amount of energy produced/absorbed: 2S + 2O2 2SO2H = – 593.8 kJ

Heat and Changes of State • Heat of combustion (∆H)= the heat of reaction for the complete burning of one mole of a substance • Molar heat of fusion (∆Hfus)= the heat absorbed by one mole of a substance in melting from a solid to a liquid at a constant temperature • Molar heat of solidification (∆Hsolid)= heat lost when one mole of a liquid freezes to a solid at a constant temperature (equal to the negative heat of fusion) • Molar heat of vaporization (∆Hvap)= the heat absorbed by one mole of a substance in vaporizing from liquid to a gas • Molar heat of condensation (∆Hcond)= heat released by one mole of a vapor as it condenses

Example(Heat of combustion) • The standard heat of combustion (∆H°rxn) for glucose (C6H12O6) is -2808 kJ/mol. If you eat and burn 71 g of glucose in one day, how much energy are you getting from the glucose? • C6H12O6 + 9 O2→ 6 CO2 + 6 H2O ΔH = -2808 kJ • Step one: convert g of glucose to moles • 70. g glucose 1 mol = 0.28 mol glucose 180.1559 g • Step two: Use (∆H°rxn) to find amount of kJ gained • 0.28 mol glucose x 2808 kJ = 790 kJ gained (+ b/c gained not lost) 1 1 mol

Thermochemistry