Chapter 17

Reaction Energy and Reaction Kinetics

Sect. 17-1: Thermochemistry

Thermochemistry – the study of the transfers of energy as heat that accompany chemical reactions and physical changes

Calorimeter – device used to measure the energy absorbed or released as heat in a chemical or physical change

Temperature – a measure of the average KE of the particles in a sample of matter

Joule (J) – the SI unit of heat and energy; kJ is also commonly used

Heat – energy transferred between samples of matter because of a difference in their temperatures

Specific heat – the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius. Units are typically J/(g x°C) or cal/(g x°C) q = Cp x m x ΔT, where q is heat, Cp is

specific heat, m is mass, & ΔT is change in temperature

Example: a 4.0 g sample of glass was heated from 274 K to 314 K, a temperature increase of 40 K and was found to have absorbed 32 J of energy as heat. What is the specific heat of this type of glass?

Heat of reaction – quantity of energy released or absorbed as heat during a chemical reaction (difference between stored energy of reactants and products)

Thermochemical equation – an equation that includes the quantity of energy released or absorbed as heatEx: 2 H2 (g) + O2 (g) 2 H2O (g)+ 483.6kJ

Enthalpy change (Δ H) – the amount of energy absorbed or lost by a system as heat during a process at constant pressure ΔH = Hproducts – Hreactants

Negative for exothermic reactions and positive for endothermic reactions

Important things to remember Coefficients represent # moles & can

be written as fractions if need be Physical state of reactants/products

matters Change in energy is directly

proportional to number of moles reacting

ΔH is usually not significantly influence by changing temperature

Molar heat of formation – the energy released or absorbed as heat when one mole of a compound is formed by combination of its elements

When given for the standard state of that substance it is written as ΔH0

f; the 0 is for standard state and the f for heat of formation

Substances that have a large negative ΔH0

f are very stable

Small negatives or small positive ΔH0f

are relatively unstable and will decompose easily

Large positive ΔH0f are very unstable

and will decompose or react violently

Heat of combustion ΔH0c – the energy

released as heat by the complete combustion of one mole of a substance

Calculating Heats of Reaction

Hess’s Law – the overall enthalpy change in a reaction is equal to the sum of the enthalpy changes for the individual steps in the process If a reaction is reversed, the sign of ΔH is

also reversed Multiply the coefficients as needed

Example

Calculate the heat of reaction for the combustion of nitrogen monoxide gas to form nitrogen dioxide gas. NO + ½ O2 NO2

From table A-14 on pg. 902:

½ N2 + ½ O2 NO ΔH0f = +90.29kJ/mol

½ N2 + O2 NO2 ΔH0f = +33.2 kJ/mol

Heats of formation can be determined by combining the heat of formation and heat of combustion for various substances

Sect 17-2: Driving Force of Reactions

Most spontaneous reactions tend toward products that have a lower energy state than the reactants (exothermic reactions)

However, some endothermic reactions do occur spontaneously

Entropy (S) – a measure of the degree of randomness of the particles in a system Reactants tend towards a less ordered

state of matter (example: ice melting – liquid is less organized than solid)

In general, gases have the highest entropy, then liquids, and then solids

ΔS is postive for an increase in entropy and negative for a decrease in entropy

Free energy

Nature drives processes toward lowest enthalpy and highest entropy, when these are opposite directions, the dominant factor determines the direction

Free energy (G) – the combined enthalpy-entropy function ΔG0 = ΔH0 - TΔS0

Example

For the reaction NH4Cl NH3 + HCl at 298.15 K, ΔH0 = 176 kJ/mol and ΔS0 = 0.285 kJ/(molK). Calculate ΔG0, and tell whether this reaction can proceed in the forward direction at 298.15K.

+ 91 kJ/mol, so it does not occur naturally at this temperature

Sect. 17-3: The Reaction Process

Reaction mechanism – the step-by-step sequence of reactions by which the overall chemical change occurs

Intermediates – species that appear in some steps but not in the net equation

Homogeneous reaction – a reactions whose reactants and products exist in a single phase

Collision theory – set of assumptions regarding collisions and reactions if collision is too gentle, the species

rebound unchanged If colliding species are poorly oriented,

they will not react

Activation energy (Ea) – the minimum energy required to transform the reactants into an activated complex

Activated complex – a transitional structure that results from an effective collision and that persists while old bonds are breaking and new bonds are forming (not the same as intermediate)

(a) activation energy for forward reaction

(b) Activation energy for reverse reaction

(c) Energy change in reaction

http://www.bbc.co.uk/scotland/education/bitesize/higher/img/chemistry/calculations_1/pe_diags/fig10.gif

Sect. 17-4: Reaction Rate

Reaction rate – the change in concentration of reactants per unit time as a reaction proceeds

Chemical kinetics – the area of chemistry that is concerned with reaction rates and reaction mechanisms

Rate-influencing factors

Nature of reactants Surface area Temperature Concentration Presence of catalysts

Catalyst – increases rate of reaction with out being used up

Catalysis – action of a catalyst Homogeneous catalyst – same phase

as reactants/products Heterogeneous catalyst – different

phase as reactants/products

Rate law – an equation that relates reaction rate and concentrations of reactants for a reaction

If multiple steps in reaction mechanism, the slow step always controls the rate