Thermochemistry

Thermodynamics Study of energy transformations Thermochemistry is a branch of thermodynamics which describes energy relationships in chemical reactions

Energy Capacity to do work or to transfer heat Mechanical work (w) is the product of force (F) operating on an object and the distance (d) through which it moves W = F x d Energy is required to do work

Heat (Q) Heat is the energy transferred from one object to another due to a difference in temperature

Forms of Energy Kinetic Energy – energy of motion - magnitude depends on the mass of the object and its velocity - E K = ½ m v 2 - both mass and speed determine how work it can do

Potential Energy – stored energy Other forms of energy are simply types of kinetic or potential on an atomic or molecular level

Energy Units Joule (J) 1J = 1 kg m 2 / s 2 A calorie (cal) is the amount of energy required to raise the temp of 1 g of water 1 ºC 1 cal = J

Example A 145 g baseball is thrown with a speed of 25 m/s. Calculate the kinetic energy in Joules. What is the kinetic energy in calories?

Systems Portion we single out for study Surroundings is everything else outside the system When studying energy changes in a chemical reaction, the reactants and products are the system and everything else is the surroundings

Law of Conservation of Energy Energy can be converted from one form to another but cannot be created or destroyed Also called “First Law of Thermodynamics”

Internal Energy Total energy of a system – sum of kinetic and potential energies Cannot determine exact internal energy Can only determine a change in internal energy ΔE = E final – E initial

If ΔE is positive there is a gain in internal energy in the system If ΔE is negative the system lost energy to its surroundings Higher energy systems tend to lose energy and are therefore less stable

Heat and Work Any system can exchange energy with surroundings in two ways – as heat or work Internal energy increases as heat is added to or work is done on a system ΔE = Q + w Q is positive if heat is added to system w is positive if work is done on the system

Heat Changes Exothermic Reactions – when heat is given off by the reaction (-Q) Endothermic Reactions – when heat is used by the reaction (+Q)

Example As a combustion reaction occurs the system loses 550 J of heat to its surroundings and it does 240 J of work in moving a piston. What is the change in its internal energy?

State Function These are systems for whom the value of ΔE does not depend on the previous history of the sample, only on the present condition Energy is a state function Work and heat are not state functions