Overview

Key Skills

• calculate reaction rates from experimental data (bleach decomposition, zinc in hydrochloric acid, iodine clock reaction), consider the factors affecting the rates
  • nature of reactants
  • concentration
  • temperature
  • surface area
• analyzing reaction mechanism (roles and application of catalysts for analysis)
• graphically representing energy changes in reactions.

Reference

• periodic table

Specific Knowledge

• Know that reactions occur at different rates
• Experimentally determine reaction rate
• Know collision theory (with respect to reaction rates)
• Describe the energies associated with reactants becoming products
• Apply collision theory to explain how reaction rates can be changed
• Analyze the reaction mechanism for a reacting system
• Represent graphically the energy changes associated with catalyzed and uncatalyzed reactions.
• Describe the uses of specific catalysts in a variety of situations

Sig fig matters

• Significant Figures and Scientific Notations

Review Chemical Reaction

• Chemical reaction - a process where a substance undergoes a rearrangement of molecules to produce a new substance, which can be described in a written format, Chemical equation
• Reactants - the beginning materials consumed in a reaction
• Products - the new substances produced in a reaction
• We call the change in energy the Enthalpy change ($\mathrm{\Delta }H$),
  • positive in endothermic reactions (where energy is absorbed), and
  • negative in exothermic reactions (where energy is dissipated, like the combination of hydrogen and oxygen, 283 kJ of heat is dissipated for every mole of water).
• Conditions surrounding a reaction determine its rate
• Workbook P1
  • Reaction kinetics studies the rates of reactions and the factors which affect the rates.

Collision Theory

• states that molecules act as small, hard spheres which collide with and bounce off each other and transfer energy among themselves during their collisions.
• 5 rules
  • reactant molecules must collide to react
  • the rate of a reaction is proportional to the frequency of collision of reactant molecules
  • not all collisions lead to the formation of products
  • colliding reactant molecules must have sufficient kinetic energy (to overcome the repulsion between the electron clouds of molecules) and a favorable orientation in order for the collision to be effective.
  • energy changes in collision as chemical bones are broken and formed
• understand what causes reaction rates to change when the conditions of the reactions are altered
  • the effect of concentration (narrowed hallway, more students)
  • the effect of temperature (less relax time) - as the temperature increases, then the kinetic energy of the molecules increases and their moving speed increases.
  • surface area (break the pack)
  • match maker (to lower activation energy)
• extension of Kinetic Molecular Theory (states that the molecules in matter move at a rate dependent upon the heat present in the system).
• the minimum amount of kinetic energy required to overcome the repulsion (to break chemical bonds) is activation energy ($E_a$).
  • when a bond breaks, energy is taken in -> potential energy (PE) increases; when a bond forms, energy is given out -> PE decreases.
  • PE increases when two molecules approaches each other
• Meanwhile the reacting molecules must have the correct collision geometry for the collision to be successful

Reaction Rates

• Reaction rate, the rate of appearance of a product or rate of disappearance of a reactant, units are moles/second, or grams/second, proportional to the coefficients in a balance equation
  • = (amount of product forms) / (time interval) = (amount of reactant used) / (time interval).
  • (disregard the sign, positive or negative) is equal to the slope of graph of change on color intensity (Mg ions are colorless), temperature, pressure in a sealed vessel of reactions producing gases, or mass if only solid present versus time
  • the rate can change on a graph, not necessarily a straight line
  • precise rate at a particular time is the slope of the tangent passing the point on the graph
• Remember to check the status of the reactants and products
  • $\mathrm{H}{\phantom{A}}_2(\mathrm{g})+\mathrm{Cl}{\phantom{A}}_2(\mathrm{g})⟶2\mathrm{HCl}(\mathrm{g})$
• complete -> net Ionic equations (don't break down the gas) + spectator ions

Factors affecting reaction rates

• temperature
• concentration
• pressure
• nature of the reactants
• surface area (only applicable when reactants are solids, powered platinum to remove the pollutants).
• catalyst (chemicals increasing the reaction rate when added to a reaction) and inhibitors (chemicals reducing the reaction rate when added to a reaction)

Homogenous reactions (all the reactants are in the same phase or state, reactions between two gases, between two solutions, liquids)

Heterogeneous reactions (the reactants are in different phases/states, like a reaction between a solid piece of zinc and a hydrochloric acid solution, between steel wool and oxygen in the air)

• the factors include, temperature, concentration, and surface area.

Influence of Reaction Form on Rate

• A reaction of all reactants in aqueous ionic form is very fast.
  • bonds are pre-broken by the dissociation process when the solution was mixed.
  • ions of opposite charges attract each other and are inclined to collide much more frequently
  • the combining of oppositely charged ions has a very low $E_a$.
  • e.g. $\mathrm{Sr}{\phantom{A}}^{+2}(\mathrm{aq})+\mathrm{CO}{\phantom{A}}_3{\phantom{A}}^{-2}(\mathrm{aq})⟶\mathrm{SrCO}{\phantom{A}}_3(\mathrm{s})$ (fast)
• A reaction that has all reactants in aqueous molecular or liquid form is usually slow (subject to $E_a$)
  • bonds need to break first
  • collisions are random
  • e.g. $\mathrm{C}{\phantom{A}}_{12}\mathrm{H}{\phantom{A}}_{16}(\mathrm{l})+\mathrm{H}{\phantom{A}}_2\mathrm{O}(\mathrm{l})⟶\mathrm{C}{\phantom{A}}_{12}\mathrm{H}{\phantom{A}}_{17}\mathrm{OH}(\mathrm{l})$ (slow)
• A reaction with few reactant particles is typically faster than one with many reactant particles
  • e.g. $4\mathrm{NH}{\phantom{A}}_3(\mathrm{g})+5\mathrm{O}{\phantom{A}}_2(\mathrm{g})⟶4\mathrm{NO}(\mathrm{g})+6\mathrm{H}{\phantom{A}}_2\mathrm{O}(\mathrm{g})$ (slow)
• Homogeneous reactions are typically faster than heterogeneous reactions, because the latter are limited by surface area.

Enthalpy changes in Chemical Reactions

• changes in total energy (enthalpy) during chemical reactions in relation to collision theory.
  • enthalpy $H$ is the total kinetic and potential energy in an open system at constant atmospheric pressure.
  • $\mathrm{\Delta }H=H_{prod}-H_{react}$ is the change in enthalpy during the course of a reaction
  • if the reaction is endothermic, $\mathrm{\Delta }H>0$ ; if the reaction is exothermic, $\mathrm{\Delta }H<0$
• 2 ways to show the reaction and the enthalpy changes
  • $\mathrm{\Delta }\mathrm{H}\mathrm{notation}:2\mathrm{N}{\phantom{A}}_2+\mathrm{O}{\phantom{A}}_2⟶2\mathrm{N}{\phantom{A}}_2\mathrm{O},\mathrm{\Delta }\mathrm{H}=+164\mathrm{kJ}$.
  • thermochemical equation, or standard notation (energy, positive on the either side of the equation) $2\mathrm{N}{\phantom{A}}_2+\mathrm{O}{\phantom{A}}_2+164\mathrm{kJ}⟶2\mathrm{N}{\phantom{A}}_2\mathrm{O}$
  • $E{a}_{(forward)}=25kJ,E{a}_{(reverse)}=40kJ\to \mathrm{\Delta }{H}_{(forward)}=-15kJ$
• curve of potential energy (PE, different from kinetic energy (KE)) vs per the x-axis is labelled the progress of reaction (time cannot be reversed, only in successful reaction)
  • PE is the energy that results from an object's position in space, including the sum of all attractive and repulsive forces; KE is the energy that results from the movement in or of the system.
  • Bond energy, is the energy required to break a bond between two atoms
  • due to the electrons on the outer edges between atoms (Hydrogen and Chlorine), they slow down, the PE rises (like a pressed spring), as they reach the peak, the moment of decision for the reaction occurs, forming the activated complex (temporary unstable species)
  • if the collision is successful, old ones break, new bonds form, and the product separates, and the PE falls down to the right side of the curve
  • if the collision is unsuccessful, the reactant molecules bounce off each other and return down the PE curve from where they came.
  • both reactants and products have low PE and high kinetic energy (KE)
• $Ea$, considered as barrier for the reaction to be successful, like to go pay rice
  • fixed by the nature of the reactants (numbers and strengths of bonds in reactants)
  • not affected by temperature, concentration or the surface area of the solid
• Temperature determines how many (or what fraction of the) molecules will have $KE\ge Ea$, hence to make it over to the barrier).
  • kinetic energy distribution graph
  • It is faster at high temperature because a larger fraction of the molecules had sufficient KE to form the activated complex.

Kinetic energy distributions

• Kinetic energy is the energy of moving particles (may move at random speed), as $KEor{E}_k$
• $E_k=\frac{1}{2}m{v}^2$
  • the greater the speed is, the higher the kinetic energy
• the total number of particles with kinetic energies within a certain range is indicated by the area under the curve
  • activation energy is the minimum KE (sometimes as ME) particles must have to have a successful collision.
  • as the temperature increases, there will be more fractions of the particles have kinetic energies above the activation energy.
  • average kinetic energy is proportional to temperature
  • under higher temperature the curve moves to the right and spreads out, shorter and wider, with the areas below the curve adds to the same
  • 0 particles at kinetic energy, because all particles are moving
• When the activation energy is near the right side of the curve, increasing the temperature by $10\mathrm{°}C$
  • about a 3% increase in the number of collisions due to increased molecular speed
  • doubles (100% increase) the fractions of the molecules that can overcome the activation energy barrier for successful collisions (NOT the increased number of collisions),
  • for the most slow reactions at or near room temperature, the reaction rate is doubled
  • if the reaction is already fast, you don't see this doubling effect.
• Temp (in Kelvin) = Temp (in degrees Celsius) + 273.15
  • increased with 10 Celsius degree from 25 to 35, the Kelvin reflects more accurately the percentage increase in kinetic energy (from 298.15 to 308.15).

Activation Energy

• activated complex is the arrangement of atoms that occurs when the reactants are in the process of rearranging to form products; of high energy and unstable.
  • activation energy is the minimum potential energy required (the difference) to change the reactants into the activated complex.
• catalysts and enzymes take the reactants through an alternative pathways, a different activated complex forms.
  • still higher than the product energy, but lower than the activation energy without catalysts.
• when temperature was increased, the potential energy of reactants, products, and activated complex remain unchanged.

Reaction mechanism

• reaction mechanism - the actual sequence or series of steps that make up an overall reaction.
  • deduced through much study and research
• e.g. $4\mathrm{HBr}+\mathrm{O}{\phantom{A}}_2⟶2\mathrm{H}{\phantom{A}}_2\mathrm{O}+2\mathrm{Br}{\phantom{A}}_2$, has 3 steps, elementary processes.
  • step 1, an HBr and an $\mathrm{O}{\phantom{A}}_2$ molecule collide; $\mathrm{HBr}+\mathrm{O}{\phantom{A}}_2⟶[\mathrm{HOOBr}]\mathrm{activated}\mathrm{complex}⟶\mathrm{HOOBr}$ (slow), limits the rate of the overall reaction, the Rate-Determining Step (RDS).
  • step 2, $\mathrm{HBr}+\mathrm{HOOBr}⟶2\mathrm{HOBr}$ (fast)
  • step 3, $\mathrm{HBr}+\mathrm{HOBr}⟶[\mathrm{H}{\phantom{A}}_2\mathrm{OBr}{\phantom{A}}_2]⟶\mathrm{H}{\phantom{A}}_2\mathrm{O}+\mathrm{Br}{\phantom{A}}_2$ (very fast)
• reactants - substances which first appear on the left hind side of the mechanism and also on the left side of the net reaction
  • product first appear on the right hind side of the mechanism and also are on the right side of the net reaction
  • intermediates first appear on the right hind side of the mechanism and then appear on the left hind side in a following step of the mechanism, and later cancelled out and does not appear in the net reaction, also called temporary products.
  • catalysts first appear on the left hind side of the mechanism and then appear on the right hind side in a following step of the mechanism, only need very little as they will get recycled throughout the reaction.
• enthalpy changes in catalyzed chemical reactions
  • 3 bumps on the curve correspond to 3 steps in the mechanism
  • the highest bump corresponds to the RDS
  • the $Ea$ of the overall action always is the Ea of the highest bump.
  • $\mathrm{\Delta }H$ remains the same for uncatalyzed and catalyzed reactions.
• Concentration vs Time - the absolute value of the slope (can be negative, rise/run) of the graph is the reaction rate (always the positive).

Catalysts

• catalysts provide alternative mechanisms for a reaction with a lower overall activation energy.
  • By lowering activation energy, this increases the probability that a reaction will occur because a greater number of molecules will have the kinetic energy needed to effectively collide and react.
• catalysts speed up the reaction.
• enzymes are proteins that act as catalysts, the effectiveness is affected by the temperature, pH
• yeast (catalase), laundry detergent, amylase in the saliva, enzymes in liver.

Glossary

• STP (Standard Temperature and Pressure).
• aq (aqueous state)
• Molar mass - in the periodic table
• whole number (suggesting a decimal)