In Chapter 6 you are responsible for sections 6.1-6.6, we do not cover 6.7.

This chapter is concerned with energy changes in chemical reactions. Thermochemistry, the study and measurement of heat effects of chemical reactions is a branch of thermodynamics, the science of energy, heat and work.

1 Energy, heat and work (Chang sections 6.1, 6.2 and 6.3, pp. 224-232)

Energy can be in a number of forms including:

Thermal energy (related to the temperature of a system)

Chemical energy (the energy stored in the chemical bonds of a molecule) Potential energy (e.g. gravitational energy)

Radiant energy, the energy contained in electromagnetic waves (light, infra red radiation, UV radiation, etc.).

Energy, heat and work are expressed in the same unit, the joule, symbol J. We have already seen in Chapter 5 that:

work (J) = force(N) x distance over which the force acts (m), or:

w = F.dJ = N.m

The unit of heat energy formerly was the calorie, cal:

1 calorie is the amount of heat needed to raise the temperature of 1 g water by 1 oC. (Raising the temperature by 1 oC is the same as raising the temperature by 1 K).

The relation between the calorie and the joule has to be determined by experiment, it is not one of “nature’s constants”. In fact, the first scientist to try and measure the equivalency of heat and work was no other than ….. Mr. Joule! The exact conversion factor was determined long after Joule’s efforts:

1 cal = 4.184 J

From the definition of the calorie above, we recognize that the calorie is in fact the specific heat of water:

The specific heat, s, is the amount of heat necessary to increase the temperature of 1 g substance by 1 oC (or by 1 K). For instance; the specific heat of water, s(H2O(l)) = 4.184 J/g.K The specific heat of copper metal, s(Cu(s)) = 0.385 J/g.K

The heat capacity, C , of a substance is the amount of heat needed to increase the temperature of a given quantity of substance by 1 oC (or 1 K). C has the units J/K

The heat (symbol q) needed to raise m grams of substance from T1 to T2 is:

q = s.m.(T2 − T1) = s.m.ΔT

The heat needed to raise the temperature of an object with heat capacity C from T1 to T2 is:

q = C(T2 − T1) = CΔT

Examples.

1. How much heat is needed to raise the temperature of 1000 g water from 20 oC to 100 oC? Answer.

q = s.m.ΔT = 4.184(J/K.g)x1000(g)x80(K) = 3.35x103 J or 3.35 kJ Note that for temperature differences ΔT in oC and K is the same!

2. How much heat is released when a piece of copper, Cu, weighing 95 g is cooled from 400 oC to 27 oC? The specific heat of Cu(s) = 0.385 J/K. Answer.

q = s.m.ΔT = 0.385(J/g.K)x95(g)x(−373)(K) = −1.36x104 J or −13.6 kJ Note that in this case q is indeed negative, the heat is given off to the environment (and lost by the system, i.e., the copper)

Heat, work and the first law of thermodynamics

The first law of thermodynamics states that:

different forms of energy can be converted into each other, but the total energy (i.e., the sum of all different forms of energy) must be conserved, it cannot be created or destroyed.

In chemical systems, we are concerned with heat and work as the forms of energy. Then the first law of thermodynamics can be written in equation form as:

ΔE = q + w

Where:q = the heat added to a system

w = the work added to the system

ΔE is the change in energy of the system

In other words, the first law in equation form expresses that when we add heat to a system, or add work to a system (“do work on the system”), the energy of the system must increase by exactly the amount of heat and work added. This is most easily seen if ΔE, q, and w all are expressed in J.

In the first law, the ‘work term” w is often in the form of “pressure-volume work”, it is the work we do (added to the system) when we compress a gas, or the work the system does (lost to the system) when a gas expands...