Thermochemistry

THERMOCHEMISTRY The study of heat released or required by chemical reactions Fuel is burnt to produce energy - combustion (e.g. when fossil fuels are burnt) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) + energy

Energy is the capacity to do work Thermal energy is the energy associated with the random motion of atoms and molecules Chemical energy is the energy stored within the bonds of chemical substances Nuclear energy is the energy stored within the collection of neutrons and protons in the atom Electrical energy is the energy associated with the flow of electrons Potential energy is the energy available by virtue of an object’s position 6.1

Two main general forms of energy Kinetic energy (E K ) = ½ mv 2 Potential energy (E P ) = mgh Energy due to motion Energy due to position (stored energy) Energy is measured in the standard unit of Joules 1 J = 1 kg ∙ m 2 /s 2 mass must be in kg velocity must be in m/s height must be in meters g = acceleration due to gravity must be in m/s 2

First Law of Thermodynamics First Law of Thermodynamics: the total energy of the universe is constant and can neither be created nor destroyed; it can only be transformed. internal energy The internal energy, U or E, of a sample is the sum of all the kinetic and potential energies of all the atoms and molecules in a sample i.e. it is the total energy of all the atoms and molecules in a sample

Total Internal Energy = Kinetic Energy + Potential Energy U t (E t ) = E K + E P Kinetic energy & potential energy are interchangeable Ball thrown upwards slows & loses kinetic energy but gains potential energy The reverse happens as it falls back to the ground

Heat is the transfer of thermal energy between two bodies that are at different temperatures. Energy Changes in Chemical Reactions Temperature is an indirect measurement of the thermal or heat energy. Temperature is NOT Thermal Energy 90 0 C 40 0 C greater thermal energy 6.2 greater temperature

UNITS OF ENERGY S.I. unit of energy is the joule (J) Heat and work ( energy in transit) also measured in joules 1 kJ (kilojoule) = 10 3 J Calorie (cal): 1 cal is the energy needed to raise the temperature of 1g of water by 1 K 1 cal = J 1000 cal = 1 Calorie (food calorie)

The specific heat (C) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius. Heat (q) absorbed or released: q = m C  t M = mass of substance C = specific heat of substance  t = t final - t initial

How much heat is given off when an 869 g iron bar cools from 94 0 C to 5 0 C? C of Fe = J/g 0 C  t = t final – t initial = 5 0 C – 94 0 C = C q = mC  t = 869 g x J/g 0 C x –89 0 C= -34,000 J

openclosedisolated SYSTEM 6.2 SURROUNDINGS In thermodynamics, the world is divided into a system and its surroundings A system is the part of the world we want to study (e.g. a reaction mixture in a flask) The surroundings consist of everything else outside the system

OPEN SYSTEM: can exchange both matter and energy with the surroundings (e.g. open reaction flask, rocket engine) CLOSED SYSTEM: can exchange only energy with the surroundings (matter remains fixed) e.g. a sealed reaction flask ISOLATED SYSTEM: can exchange neither energy nor matter with its surroundings (e.g. a thermos flask)

Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process is any process in which heat has to be supplied to the system from the surroundings. 2H 2 (g) + O 2 (g) 2H 2 O (l) + energy H 2 O (g) H 2 O (l) + energy energy + 2HgO (s) 2Hg (l) + O 2 (g) energy + H 2 O (s) H 2 O (l)

Photosynthesis is an endothermic reaction (requires energy input from sun) Forming Na + and Cl - ions from NaCl is an endothermic process Burning fossil fuels is an exothermic reaction

Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. (comes from Greek for “heat inside”)  H = H (products) – H (reactants)  H = heat given off or absorbed during a reaction at constant pressure H products < H reactants  H < 0 H products > H reactants  H > 0

Thermochemical Equations H 2 O (s) H 2 O (l)  H = 6.01 kJ Is  H negative or positive? System absorbs heat Endothermic  H > kJ are absorbed for every 1 mole of ice that melts at 0 0 C and 1 atm.

Thermochemical Equations CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O (l)  H = kJ Is  H negative or positive? System gives off heat Exothermic  H < kJ are released for every 1 mole of methane that is combusted at 25 0 C and 1 atm.

Work is the transfer of energy that takes place when an object is moved against an opposing force i.e. a system does work when it expands against an external pressure, like the pressure of air Car engine: gasoline burns & produces gases which push out pistons in the engine and transfer energy to the wheels of car Heat and work are 2 equivalent ways of changing the internal energy of a system

+= Change in internal energy Heat either gained or lost by the system Work done either by or on the system  U (E) = q (heat) + w (work) q w q w U U (E) like reserves of a bank: bank accepts deposits or withdrawals in two currencies (q & w) but stores them as common fund, U (E).

First Law of Thermodynamics can be reworded…. - The internal energy of an isolated system is constant – or,  U = q + w

EXPANSION WORKAn important form of work is EXPANSION WORK i.e. the work done when a system changes size and pushes against an external force like air pressure e.g. the work done by hot gases in an engine as they push back the pistons In a system that can’t expand, no work is done (w = 0)  U (E) = q + w when w = 0,  U (E) = q (at constant volume) MOST CHEMICAL REACTIONS WE OBSERVE SIMPLY WASTE HEAT AND DO NO WORK – THEY ARE NOT HOOKED UP TO MACHINES IN THE LAB!

6.7 What if work is done by a chemical reaction? In this example, the chemical reaction is producing hydrogen gas The hydrogen gas is doing work by pushing the piston against the air pressure Work is being done by the system W = F x d How can we translate that here?

6.7 w = F x d The force the reaction is working against is from the air Pressure = Force/Area So, w = Pressure x Area x distance The distance here is the height that the cylinder has to move up against the air The area would be the area of the circle, or piston Area x height of a piston = volume So, w = P x  V Since the work is done by the system, and should be negative, we say: w = - P x  V h Work is in Joules Pressure must be in the standard unit of Pascals Volume must be in the standard unit of m 3

Mathematically, enthalpy is defined as: H = U + PV We also know that  U = q + w, and w = - P  V So,  U = q – P  V Rearranging, we can also say that: q =  U + P  V Also,  H =  U + P  V If no work is done by a chemical reaction, then there is no volume change at a constant pressure, and therefore: q =  U and  H =  U, so:  H = q Telling me that all of the heat lost by the reaction, since no work is done, is equal to the change in enthalpy – This is common sense, since enthalpy is supposed to measure the heat change of a reaction!

A state function is a property whose value does not depend on the path taken to reach that specific value. Potential energy of hiker 1 and hiker 2 is the same even though they took different paths. Change in enthalpy, pressure, volume, temperature, potential energy are all state functions For example: Density is a state function because a substance's density is not affected by how the substance is obtained. Lets say we have a certain amount H 2 O, it does not matter whether that H 2 O is obtained from one's tap, from a well, or from a bottle, because as long as all three have the same states, they will have the same density. When deciding whether a certain property is a state function or not, keep this rule in mind: is this property or value affected by the path or way taken to establish it? If the answer is no, then there is a state function, but if the answer is yes, then there is no state function.

There are four ways to measure the heat of a chemical reaction: –Calorimetry – using a calorimeter to experimentally measure the heat released or absorbed from a reaction –Hess’ Law – using other chemical reactions to algebraically solve for the heat of a reaction –Heats of formation – using a table of potential energies to measure the potential energies of the reactants and products of a reaction, and thereby calculating the  H of a reaction:  H = Σ H prod - Σ H reac –Add the energy it takes to break the reactant’s bonds, and subtract the energy released when the new bonds are made:  H = Energy in – Energy out = Bonds broken – Bonds made How do we measure the heat (  H) of a reaction…?

No heat enters or leaves! A calorimeter is used…. A calorimeter is an insulated device used to capture all of the heat either absorbed or released by a reaction! The reaction is usually surrounded by water….why? Water is stable, and has a high specific heat It changes temperature slowly! q reaction = - q surroundings By measuring the heat that the water absorbs or releases, we can calculate the heat of the reaction! Using calorimetry…..

A g sample of solid quinone (C 6 H 4 O 2 ) is burned in a bomb calorimeter that contains 373 grams of water. The temperature of the calorimeter increases by 3.2°C. Calculate the energy of combustion of quinone per mole. First – write a balanced chemical equation! 1 C 6 H 4 O 2 (s) + 6 O 2 (g) → 6 CO 2 (g) + 2 H 2 O (g) The heat released by the reaction is absorbed by the calorimeter: q = mc  T q reaction = - q calorimeter q = (373 g H 2 O)(4.184 J/g 0 C)(3.2°C) = J gained by calorimeter q reaction = J

This is not the energy of combustion, though! 1 C 6 H 4 O 2 (s) + 6 O 2 (g) → 6 CO 2 (g) + 2 H 2 O (g) The energy of combustion, or  H, is the energy released for the reaction the way it was written! We only used.1964 grams of the chemical! The reaction calls for one mole of the chemical! 1 mole C 6 H 4 O 2 (s) So I set up a ratio: J/.1964 g = X/108 g X = -2,746, J So,  H combustion = -2,700,000 J (2 significant figures  H combustion = -2,700 KJ

The other two methods, Hess’ Law and Heats of Formation, will be discussed in class next time!

Just to clarify…. Molar heat of combustion is the heat change for one mole of a substance burning:  H combust The standard heat of formation is the heat change for a substance being formed from the elements:  H form The molar enthalpy is the  H or change in heat for one mole of a substance – not necessarily for how the reaction is written! Standard enthalpy change is  H o, which means the heat change for a reaction the way it is written under standard conditions. Enthalpy change or  H is the heat change for the reaction the way it is written at any conditions.