Reaction Energy and Reaction Kinetics
All reactions involve some type of energy.
Thermochemistry - definition
Driving Force of Reactions
- Heat and Temperature
Calorimeter - definition
amount of energy determined by the rise in temperature of the water whose mass is known
Temperature - definition
Joule - definition
more commonly use kilojoule (kJ)
Heat - definition
Moves from area of high temperature to area of low temperature.
Figure 17-1 page
- Heat Capacity and Specific Heat
The amount of energy transferred as heat during a temperature change depends on the
a) nature of the material changing temperature;
b) the mass of the material changing temperature;
c) the size of the temperature change.
Specific heat - definition
units may be joule/gram . degree Celsius or
calories /gram . degree Celsius
Table 17-1 page 513
cp = q / m x delta T or
q = cp x m x delta T
where cp is the specific heat at a given pressure
q is the energy lost or gained
m is the mass of the sample
delta T is the difference between the initial and final temperature
Sample Problem 17-1 page 513
A 4.0 g sample of glass was heated from 274 K to 314 K, a temperature increase of 40K, and was found to have absorbed 32 J of energy as heat.
a) What is the specific heat of this type of glass?
b) How much energy will the same glass sample gain when it is heated from 314 K to 244 K?
Given: m = 4.0 g; delta T = 40 K; q = 32 J; cp = ?
a) cp = q / m x delta T
cp = 32 J / (4.0 g) (40K)
cp = 0.20 J/g . K
b) q = cp x m x delta T
(0.20 J/g.K) x (4.0 g) x (30 K) = 24 J
- Heat of Reaction
Heat of reaction - definition
Similar to the difference between the potential energy of the reactants v products.
A formula equation does not indicate anything about a quantity of heat e.g.
2 H2(g) + O2(g) ---> 2 H2O(g)
Thermochemical Equation - definition
2 H2(g) + O2(g) ---> 2 H2O(g) + 483.6 kJ
The quantity of energy involved depends on the amount of reactants and products e.g.
4 H2(g) + 2 O2(g) ---> 4 H2O(g) + 967.2 kJ
Note: sign convention and where energy term appears in the equation.
Need to include physical states of all reactants and products since this will affect the energy involved.
Cant be measured directly - only change in enthalpy viz. delta
Enthalpy change - definition
can be expressed as:
delta H = H products - H reactants
sign convention when not written as part of a chemical reaction
Figure 17-2 page 516
Figure 17-3 page 517
Things to remember about thermochemical equations:
a) coefficients in the balanced equation represent moles, not molecules
b) physical states of all chemicals must be specified
c) Change in energy is directly proportional to the number of moles of substance undergoing the change
d) the numeric value of delta H is usually not affected by changing temperature
- Heat of Formation
specialized form of enthalpy
molar heat of formation - definition
normally these energy values are given for room temperature and stand pressure
use of zero as a superscript with enthalpy symbol
delta H0 and delta Hf0 i.e. standard enthalpy or standard heat of formation
Appendix A Table A-14
- Stability and Heat of Formation
If a compound is formed during a highly exothermic reaction, heat of formation is negative and large numerically, then the compound is usually stable since it is in a lower energy state than its elements.
Elements in their standard state are considered to have a molar heat of formation of zero.
Table A-14 page --most compounds have a negative heat of formation.
Compounds with positive values of heats of formation or only slightly negative values for heats of formation are relatively unstable and will break down into their elements given the proper conditions.
e.g. HI has a heat of formation = +26.5 kJ/mol. Easily breaks down into iodine vapors which are red/brown in color.
If the heat of formation is a high positive, it may be unstable and dangerous. e.g. ethyne or acetylene, C2H2, has a value of +226.7 kJ/mol. It react violently with oxygen and must be stored in cylinders as a solution in acetone.
Mercury fulminate, HgC2N2O2, has a heat of formation of + 270 kJ/mol and it is used as a detonator for explosives.
- Heat of Combustion
Heat of combustion - definition
N.B. one mole of reactant in definition
Table A-5 page
All are exothermic.
C3H8(g) + 5 O2(g) ---> 3 CO2(g) + 4 H2O(l) delta H0c = -2219.2 kJ/mol
Figure 17-4 Combustion Calorimeter page 519
- Calculating Heats of Reaction
Can rearrange and combine various reactions to calculate the enthalpy change for a reaction not listed in a table of enthalpies. Handle equations just like algebraic equations.
Hesss Law - statement
The route for getting from reactants to products is not critical, just the reactants and products.
want to get the heat of formation for the following reaction
C(s) + 2H2(g) ---> CH4(g)
use the following reactions:
Eq a: C(s) + O2(g) ---> CO2(g) delta H0c = -393.5 kJ/mol
Eq b: H2(g) + 1/2 O2(g) ---> H2O(l) delta H0c = -258.8 kJ/mol
Eq c: CH4(g) + 2 O2(g) ---> CO2(g) + 2H2O(l) delta H0c = -890.8 kJ/mol
Rules to remember:
a) if you reverse the reaction, reverse the sign of the delta H
b) multiply the coefficients of the known equations so that when added together they give the desired thermochemical equation.
Reverse Eq c and reverse the sign of the enthalpy.
Multiply Eq b by two to get two moles of water and multiply the enthalpy by two also.
This gives us three equations which we can manipulate algebraically to get the equation we seek:
Sample Problem 17-2 page 521
- Determining Heat of Formation
We can apply Hesss Law to calculating the heats of formation also.
Figure 17-5 Diagram
for Heat of reaction for carbon dioxide page 523
Sample Problem 17-3 page 523
Calculate the heat of formation of pentane, C5H12, using the information on heats of formation in Appendix Table A-14 and the information on heats of combustion in Appendix Table A-5. Solve by combining the known thermochemical equations.
Two factors allow chemists to predict if a reaction will occur spontaneously and to explain how it occurs:
a) the change in energy of a reaction system;
b) the randomness of the particles in a system.
The Reaction Process
- Enthalpy and Reaction Tendency
Most reactions are exothermic, indicating that the products are in a lower energy state than the reactants. This usually makes the products more stable than the reactants.
Some endothermic reactions occur spontaneously. Thus enthalpy is only one factor determining spontaneity of reactions.
- Entropy and Reaction Tendency
Consider melting of ice.
It is spontaneous, it is endothermic.
The liquid is a more disordered state than the solid state of ice. Both exist at the melting point. Thus endothermic (product in higher energy state than reactant) and also an increase in the disorder of the system.
The increase in disorder can cause an endothermic process to occur spontaneously.
2 NH4NO3(s) ---> 2 N2(g) + 4 H2O(l) + O2(g)
Compare reactants and products regarding disorder.
Tendency to proceed in the direction which produces more disorder.
Compare solids, liquids, and gases and their entropies.
Comparison should be for the same substance in the different physical states.
At absolute zero, random motion ceases, thus, the entropy of a pure crystalline solid is zero at absolute zero.
Standard molar entropy or absolute entropy
units are kj / (mol-K)
Cant measure entropy but can measure entropy change which is the difference
between the entropy of the products and the entropy of the reactants.
If entropy increases, delta S has a positive value.
Forming a solution from a solid and a liquid, for example, results in an increase in entropy.
Figure 17-7 page 528 no link
- Free Energy
We saw the two factors that drive process in nature are low enthalpy and high entropy.
When there is a conflict between these two, free energy, G, determines if the change occurs. Processes tend to occur with a reduction or decrease in free energy.
Cant measure free energy only change in free energy, delta G.
Free energy - definition
Delta G0 = delta H0 - TdeltaS0
units for delta G and delta H are kJ/mol
units for delta S are kJ / (mol.K)
Each term in the equation, except for T, can be either positive or negative.
Table 17-2 page 529 - signs and implications N.B. minus sign operator
Sample Problem 17-4 page 530
For the reaction NH4Cl(s) ---> NH3(g) + HCl(g) at 298.15 K, delta H0 = 176 kj/mol and delta S0 = 0.285 kJ/ (mol.K). Calculate delta G0, and tell whether this reaction can proceed in the forward direction at 298.15 K.
Given: delta H0 = 176 kJ/mol; delta S0 = 0.285 kj/ (mol.K); T = 298.15 K; delta G0 = ?
delta G0 = delta H0 - TdeltaS0
delta G0 = 176 kJ/mol -( 298K [0.285 kJ / (mol.K) ] )
= 176 kJ/mol - 84.9 kJ/mol
= 91 kJ/mol
Positive value indicates reaction does not occur naturally.
Going from reactants to products is not necessarily a direct route. There may be several intermediate steps involved in this transition. The enthalpy, entropy, and free energy is independent of the number of stops involved.
- Reaction Mechanism
We will concentrate on reactions between molecules. Electron clouds of molecules repel each other and so the molecules need a certain minimum kinetic energy to overcome that repulsion. This will allow them to collide and react.
In the reversible reaction:
H2(g) + I2(g) ----> 2 HI(g)
Hydrogen is colorless, iodine is violet, and hydrogen iodide is colorless.
The above reaction is the overall reaction -- it shows the reactants and the products. It does not show any intermediate steps that may have occurred to get from reactants to products.
Reaction mechanism - definition
To determine the reaction mechanism you must design an experiment which will provide that information. Each step in the mechanism shows only those atoms, ions or molecules that participate in that particular step. Algebraically, the steps must equal the overall reaction.
Intermediates - definition
Mechanism page 532 for
two possible reaction mechanisms. Notice how cancellation comes into play to
give the final reactants and products.
Homogeneous reaction - definition.
Homogeneous chemical system - definition
- Collision Theory
No collision between molecules means no reaction.
Collision theory - definition.
For the reaction: AB + AB = A2 + 2B
Figure 17-9 page
Collision must be of sufficient energy and the orientation of the molecules must be favorable since bonds will not form over relatively large distances.
This is what happens in 17-9 c.
Thus a collision may fail to produce a reaction if the molecules collide too softly or the molecules are not properly aligned.
- Activation Energy
For the reaction: 2 H2(g) + O2(g) ---> 2 H2O(g)
the heat of formation is high ( -285.8 kJ/mol at 298.15 K ) and the reaction is exothermic indicating that the product, water is more stable than its reactants.
the free energy change ( delta G ) is also large (-237.1 kJ/mol ).
Yet if you mix hydrogen and oxygen together in a container and do not disturb them, they will not react. YOu need to provide a spark in order to get them to react.
When hydrogen and oxygen molecules collide the electron clouds repel each other and so the molecules bounce off and never actually meet. The molecules need enough kinetic energy to overcome this repulsion and provide a collision of sufficient energy to get a reaction. The spark activates the reaction.
Once such a reaction begins (exothermic) the energy given off is sufficient to sustain the reaction and allow it to run to completion.
Figure 17-10 page
Activation energy - definition
Reverse reaction is endothermic ( products are in a higher energy state than reactants ). Thus the activation energy for the reverse reaction is higher than the activation energy for the forward reaction.
The difference between the activation energy for the forward reaction (Ea) and the activation energy for the reverse reaction (Ea ) represents the heat of reaction or the energy difference between the reactants and products (delta E).
delta E is the same for forward and reverse reaction except the exothermic process has a negative sign, forward reaction in this case, and the endothermic process has a positive sign, the reverse reaction in this case.
- The Activated Complex
Molecules can collide and convert some of their kinetic energy into potential energy within the molecule. If enough energy is converted, molecules with suitable orientation become activated. This allows the formation of new bonds. It is during this time that the bonds break and new bonds form.
Activated complex - definition
Figure 17-11 -
refers to the the reaction of hydrogen and iodine to produce hydrogen
iodide. Shows it as three steps.
Both forward and reverse reactions go through the same activated complex.
The activated complex is at a high energy position. The activated complex
defines the activation energy for the system.
Temperature is a measure of the average kinetic energy. When you raise
the temperature at which a reaction occurs, it causes the molecules the
speed up and therefore collide more often and it also causes molecules
to have a higher energy.
More collisions implies more chance of getting the correct alignment
of molecules and higher energy means more molecules with the activation
The activated complex has bonding characteristic of both reactants and
products. Thus the chance of reforming reactants or forming products
is equally possible.
The activated complex is different from the intermediate products of
a reaction mechanism.
Reaction rate - definition
Concerned with the factors that affect the rate and with the mathematical expressions that reveal the specific dependencies of the rate on concentration.
Chemical Kinetics - definition
- Rate Influencing Factors
Except for decomposition reactions, in other reactions, particles must come into contact in a favorable orientation and with enough energy for activation. i.e. collision frequency and collision efficiency
Five factors influence the rate of a chemical reaction.
- Nature of Reactants
What the reactants are i.e. both reactants or all reactants e.g. hydrogen readily combines with chlorine under certain conditions but under the same conditions it will hardly react with nitrogen.
- Surface Area
Substances in the gaseous state or in solution mix and collide freely and so they can usually react rapidly.
In heterogeneous reactions, the reaction rate depends on the area of contact of the reacting substances.
Heterogeneous reactions - definition
If one of the reactants is a solid, then the larger the surface area of the solid the greater the rate of the heterogeneous reaction.
e.g. Zn(s) + 2 HCl(aq) -----> ZnCl2(aq) + H2(g)
lump of zinc v powdered zinc
Increasing temperature increases average KE and can cause more effective collisions which will increase the reaction rate.
Effective collisions produce energy equal to or greater than the activation energy.
Decreasing temperature has the opposite effect.
A rule of thumb is: starting at room temperature, each ten degree rise in temperature causes a doubling of the reaction rate.
Need to determine actual increase in the lab.
Figure 17-12 page
539 - coal burning in air v pure oxygen.
This is a heterogeneous reaction system. One reactant is a gas and one is a solid. The reaction rate depends not only on the amount of exposed charcoal surface but also on the concentration of the reacting species, O2.
In homogenous reaction systems, reaction rates depend on the concentration of the reactants.
It is difficult to predict a mathematical relationship between concentration and rate in this case since the reaction occurs in steps and only one step determines the reaction rate.
In general, an increase in rate is expected if the concentration of one or more of the reactants is increased.
Figure 17-13 - concentration of reacting species affects the number of collisions and therefore the reaction rate.
The affect of concentration must be determined in the lab.
- Presence of Catalysts
Catalyst - definition
Catalysis - definition
Figure 17-14 catalysis
of the decomposition reaction of hydrogen peroxide by manganese dioxide.
A catalyst provides an alternate pathway to the products using a different activated complex that has a lower activation energy if it is speeding up the reaction.
A catalyst takes part in the reaction - that is why a new activated complex form - but they are regenerated as the products are formed.
Figure 17-15 page 541 different activated complexes for the same reaction using different catalysts.
Homogeneous Catalyst - definition
Heterogeneous Catalyst - definition
Often use metals as heterogeneous catalysts. They work by adsorbing reactants on the metal surfaces, which has the effect of increasing the concentration of the reactants.
- Rate Laws for Reactions
The relationship between the rate of a reaction and the concentration of one reactant is determined experimentally by first keeping he concentrations of other reactants and the temperature of the system constant. Then the reaction rate is measured for various concentrations of reactant in question. Doing this a number of times tells how the concentrations of each reactant affects the reaction rate.
Thermochemistry is the study of the transfers
of energy as heat that accompany chemical reactions and physical changes. back
Calorimeter measures the energy absorbed or
released as heat in a chemical or physical change. back
Temperature is a measure of the average kinetic
energy of theparticles in a sample of matter. back
A Joule is the SI unit of heat as well as all other
forms of energy. back
Heat can be though of as the energy transferred between
samples of matter because of a differenece in their temperatures. back
Specific heat is the amount of energy required
to raise the temperature of one gram of substance by one Celsius degree or
one kelvin. back
Heat of reaction is the quantity of energy
released or absorbed as heat during a chemical reaction. back
Thermochemical Equation is an equation
that includes the quantity of energy released or absorbed as heat during the
An Enthalpy change is the amount of energy absorbed
or lost by a system as heat during a preocess at constant pressure. back
The molar heat of formation is the energy
releaed or absorbed as heat when one mole of a compound is fomred by combination
of its elements. back
Heat of combustion is the energy released
as heat by the complete combustion of one mole of a susbtance. back
Hesss Law states that the overall enthalpy
change in a reaction is equal to the sum of the enthalpy changes for the
individual steps in the process. back
Entropy is a measure of the degree of randomness of
the particles, such as molecules, in a system. back
Free energy of a system is defined as the diffreence
between the change in enthalpy and the product of the Kelvin temperature and
the entropy change at constant temperature and pressure. back
Reaction mechanism is the step by step sequence
of reactions by which the overall chemical change occurs. back
Intermediates are species that appear in some
steps but not in the net equation. back
Homogeneous reaction is a reaction whose
reactants and products exist in a single phase. back
Homogeneous chemical system is a
system in which all reactants and products in all intermediate steps are in
the same phase. back
Collision theory is a set of assumptions regarding
collisions and reactions. back
Activation energy is the minimum energy required
to transform the reactants into an activated complex. back
Activated complex is a transitional structure
that results from an effective collision and persists while old bonds are breaking
and new bonds are forming. back
The reaction rate is the change in concentration
of reactants per unit time as a reaction proceeds. back
Chemical kinetics is the area of chemistry
that is concerned with reaction rates and reaction mechanisms. back
Heterogeneous reactions involve reactants
in two different phases (physical states). back
A catalyst is a substance that chnages the rate of
a chemical reaction without itself being permanently consumed. back
Catalysis is the action of the catalyst. back
A Homogeneous Catalyst is a catalyst that
is in the same phase as all of the reactansts and products in a reaction system. back
A Heterogeneous Catalyst is a catalyst
that is in a different phase from that of the reactants. back