Chemistry Review: Unit 1

Chemistry Review: Unit 1
Collision Theory
Chemical Kinetics study of the rates of chemical processes
Kinetic energy energy of motion
Particles of matter atoms, ions, and molecules
Kinetic theory of matter/ Kinetic molecular theory
o 1. Particles of matter are in a state of constant random motion
o 2. Particles of matter are often colliding
o 3. The energy associated with each collision is conserved
Properties of gasses that distinguish them from solids and liquids
 1) Gasses diffuse to occupy available space
 2) Gasses exert pressure
 3) Gasses have little or no volume compared to the space they occupy
 4) Gasses have little or no attraction for other gas molecules
 5) Gasses move in straight lines and at very high velocities until a collision
causes them to change course.
Collision Theory states that reacting particles must collide for a reaction to occur.
It is important to note that not all collisions result in chemical changes.
Effective collision molecular collision that results in a chemical reaction.
Theory of chemical reactions states that correct orientation and sufficient
intensity are necessary factors to cause old bonds to break and new bonds to form.
 The rate of a chemical reaction is dependent on…
o 1. Kinetic energy of colliding particles
o 2. Orientation of particles on impact
Reaction Rates
 In a chemical reaction, the number of successful collisions is a function of the
kinetic energy and orientation of the colliding particles
 Rate of a chemical reaction a function of the number of successful
collisions between reacting particles per unit of time.
 Particles of motion are much to small to be seen so these collisions are
measured by a change in properties. These properties include…
o 1. Mass
o 2. pH
o 3. Electrical conductivity
o 4. Pressure
o 5. Colour
o 6. Gas volume
 Reaction rate change in the amount of substance over time.
 Average reaction rate the overall rate of reaction.
Chemistry Review: Unit 1
Potential Energy Diagrams
Chemical reactions involve the breaking and forming of chemical bonds
Endothermic a bond breaking process which requires the input of energy
Exothermic a bond forming process which results in a release of energy
Potential Energy
 Energy possessed by something due to its position relative to something else
 Either gained or lost as a result of bond breaking/forming
 Potential energy changes in a chemical reaction are represented using a
potential energy diagram
Temperature and Kinetic Energy
 Kinetic energy energy of motion
 Temperature an indicator of potential energy. It is a measurement of how
fast/slow the particles are moving
 Molecules in the air of a room are in constant random motion
 They are not moving at exactly the same speed
 Therefore, when you take the temperature of that room you are getting the
average kinetic energy of the molecules (that is, the average speed of the all
the molecules)
Percentage of Molecules vs. Kinetic Energy Graph
Relationship Between Kinetic and Potential Energy
 Energy is conserved in chemical reactions  kinetic energy can be converted
to potential energy and vice versa
 Kinetic energy is necessary for particles to collide and form bonds
 At the instant of collision, the kinetic energy is zero
 Threshold energy the amount of energy required for particles to collide
hard enough to cause bonds to stretch to their breaking point.
Chemistry Review: Unit 1
Threshold Energy Graph
Kinetic energy of colliding particles is converted to potential energy
Potential energy diagram shows this conversion of energy as reactants
become products
How to Draw Potential Energy Diagrams
 Activation Energy the minimum gain in potential energy that results in
reactant bonds to stretch to their breaking points. The greater the activation
energy, the slower the reaction.
 Activated complex the unstable group of atoms that are not quite products
or reactants.
 Heat of reaction potential energy difference between reactants and
o 1. The flat region shoes the potential energy of the reacting particles
relative to the products
o 2. Moving particles posses kinetic. The rising part of the graph
represents the increase in potential energy that occurs when
reactants collide. This part of the graph is the activation energy
o 3. The top curve represents the stage where the group is not reactants
or products. This structure of unstable atoms is the activated complex.
The bonds are stretched to their breaking point so potential energy is
at its maximum.
o 4. The falling part of the curve is the energy released when new bonds
are formed. This section where the heat of reaction is labeled from the
end of the activation energy to the products. This section represents
the net energy change of the reaction. When the energy of reactants is
greater than that of the products it is exothermic. When energy of
product is greater, it is endothermic.
o 5. The last flat region is the potential energy of the products
Chemistry Review: Unit 1
Endothermic Reaction
Exothermic Reaction
Reversing a Reaction
 Heat of reaction is negative for exothermic and positive for endothermic
 If the sign of the forward reaction is negative, then the sign of the reverse
reaction is positive
Potential Energy Calculations
 ∆𝐻 = 𝑃𝑜𝑡𝑒𝑛𝑡𝑖𝑎𝑙 𝐸𝑛𝑒𝑟𝑔𝑦𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − 𝑃𝑜𝑡𝑒𝑛𝑡𝑖𝑎𝑙 𝐸𝑛𝑒𝑟𝑔𝑦𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠
 ∆𝐻 = 𝐸𝑎𝑓𝑜𝑟𝑤𝑎𝑟𝑑 − 𝐸𝑎𝑟𝑒𝑣𝑒𝑟𝑠𝑒
 Note: ALL activation energies are positive while heat of reaction may be
positive or negative
Chemistry Review: Unit 1
Rate Laws and Order of Reaction
Rate of reaction and factors affecting the reaction are mathematically related
These relationships must be determined empirically (analyzing data)
Rate Law
 Relationships between rate (r) and the product of the initial concentrations
of the reactants raised to come exponential value
𝒂𝑿 + 𝒃𝑿 → 𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔
𝒓 ∝ [𝑿]𝒎 [𝒀]𝒏
Exponents (m and n) describe relationship between rate and initial
m and n are not to be confused with coefficients of balanced chemical
m and n can be a whole number, zero, or a fraction
Determining the rate of a reaction
 Use k as the “rate constant”
 Rate Law Equation:
𝒓 = 𝒌 [𝑿]𝒎 [𝒀]𝒏
The values for k, like r, must also be determined
Orders of Reaction
 Overall order of reaction the sum of all the individual orders (exponents)
 Ex. 𝑟 = 𝑘 [𝑋]1 [𝑌]2 [𝑍]0
The order of reaction for [X] is 1, [Y] is 2, [Z] is 0
The overall order of reaction is 3
What Does the Order of Reaction Mean?
For 1st order – If concentration is doubled, rate is doubled (21)
For 2nd order – If concentration is doubled, rate is quadrupled (22)
For 3rd order – If concentration is doubled, rate is increased by a factor of 8
For 0 order – If concentration is doubled, there is no change in the rate (20)
*Note: If a reaction is zero order for one of the reactants, that reactant is left
out of the equation
Chemistry Review: Unit 1
Sample Problem
Given the following equations and experimental data, write the correct
a. Rate Law Expression
b. Reaction Order
c. Determine k, the Specific Rate Constant (including units)
1. A2 + B2  2 AB
Exp #
(mole L-1 s-1)
For A2 compare 4 & 5: [A2] x2 rate x4=2 , 2 order for A2
For B2 compare 1 & 2: [B2] x2 rate x2=21 , 1st order for B2
r  k [ A ] [B ]
[ A2 ] [ B2 ]
(0.0010) 2 (0.0010)
k  1.0 x10 7
r  1.0 x10 7 [ A ] [ B ]
reaction order = 3
Chemistry Review: Unit 1
Thermochemistry the study of energy changes that accompany physical
and chemical changes
When a chemical system undergoes a change heat, q, is transferred between
Ex. An ice pack
o Energy is being absorbed by the forming of new bonds once the water
pack is broken
o Therefore we feel the cold (heat is absorbed)
Measuring Energy Changes
 Calorimeter a device that measures energy in a reaction. It takes place in
an isolated compartment and the energy change can then be measured.
 Heat gain/release can be calculated using the following formula:
𝒒 = 𝒎𝒄∆𝑻
Where… q = heat (J or kJ)
m = mass (g)
c = specific heat capacity (J/g or kJ/kg)
∆T = change in temperature (C or K)
Heat of Reaction
 Enthalpy amount of heat produced or used during a chemical reaction
 Represented using the symbol ∆H and the unit of kJ/mol
 ∆H is also equal to the difference in the enthalpies of products and reactants
 ∆H is not affected by a catalyst
Energy Transfer
 Energy cannot be created or destroyed but it can be transferred
𝒒𝒈𝒂𝒊𝒏𝒆𝒅 = −𝒒𝒍𝒐𝒔𝒕
Note: when you react a “cold” object with a “hot” object, energy will be
transferred and the final temperatures will be the same unless otherwise
Chemistry Review: Unit 1
Phase Diagrams
A substance may change state
This state change can be represented using a phase diagram
The amount of energy necessary for that phase change to occur can be
Latent heat heat required for a phase change
When calculating q for a state change from solid to liquid, you use the latent
heat of fusion constant (varies for each element) multiplied by mass
𝒒 = 𝒎𝒍𝒇
When calculating q for a state change from liquid to gas, you use the latent
heat of vaporization constant (varies for each element) multiplied by mass
𝒒 = 𝒎𝒍𝒗
To solve a potential energy diagram, it is important to first sketch a graph of
the phase changes (be sure to note when the element changes state for
accurate calculations)
The total value of q is found by calculating the sum of all the line segments of
the sketch
Calculating ∆𝐇 (kJ/mol)
There are multiple ways to calculate ∆H
o 1. Using the value of q and the number of moles
o 2. Using Hess’ Law
o 3. Using Standard Heats of Formation
o 4. Using bond energies
o 5. Potential energy diagrams
Standard Heats Formation
 Standard enthalpy of formation (H°f) energy associated with making a
substance from its elements
 Elements have a H°f of 0 and the values for compounds can be found on the
table provided
 To calculate ∆H…
o 1. Balance the chemical equation and write the states of matter
o 2. Calculate ∆H using the following formula
∆𝑯 = ∑𝒏𝑯°𝒇𝒑𝒓𝒐𝒅𝒖𝒄𝒕𝒔 − ∑𝒏𝑯°𝒇𝒓𝒆𝒂𝒄𝒕𝒂𝒏𝒕𝒔
Chemistry Review: Unit 1
Sample Problem
C3H8 (g) + 5O2 (g)  3CO2 (g) + 4H2O (g)
∆𝐻 = ∑𝑛𝐻°𝑓𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − ∑𝑛𝐻°𝑓𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠
∆𝐻 = 3CO2 + 4H2 O − C3 H8 + 5O2
∆𝐻 = [3 (−393.5
∆𝐻 = −2043
𝑚𝑜𝑙 ) + 4 (−241.8 ⁄𝑚𝑜𝑙 ] − [(−104.7 ⁄𝑚𝑜𝑙 ) + 5 (0 ⁄𝑚𝑜𝑙 )]
Calculating ∆𝑯 Using Bond Energies
Enthalpy of a reaction can be calculated using bond energies
∆𝑯 = ∑𝒃𝒐𝒏𝒅𝒔 𝒃𝒓𝒐𝒌𝒆𝒏 − ∑𝒃𝒐𝒏𝒅𝒔 𝒇𝒐𝒓𝒎𝒆𝒅
Before summing, the bond energies must be multiplied by the number of that
bond present and the coefficients from the balanced chemical equation
Structures of each substance must be drawn to determine bonds present
Sample Problem
C2H6 + 7/2 O2  2CO2 + 3H2O
Bonds Broken
1 C-C (347 kJ/mol)
6 C-H (414 kJ/mol)
3.5 O- - O (498.7 kJ/mol)
=4576.4 kJ.mol
Bonds Formed
4 C - - O (804 kJ/mol)
6 H – O (460 kJ/mol)
=5980 kJ/mol
∆𝐻 = ∑𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛 − ∑𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑
∆𝐻 = 4576.4 − 5980
∆𝐻 = −1404 kJ/mol