Chapter 6 - Rate Of Reactions

Rate Of A Reaction

Change in concentration of a product or reactant in a given time period

Can be change gas, pH, colour

given in mol L^-1 s^-1 (moles per litre per second - how much is being depleted)

Easier way fro rate of reaction

How much of products is being used over a period of time

How much of reactants is being used over time

Using Volume Of Gas

Explanation of Example

A steeper gradient indicates a higher rate of reaction

Eventually, it evens out - less reaction occurring

Becomes zero - reaction has concluded

Rate of reaction can be calculated through gradient of graph

Using Mass Loss

By measuring weight of apparatus after during regular intervlas (CO₂) gas is lost

Conditions For A Reaction To Take Place (Collision Theory)

Must collide with each other

Must collide with sufficient force to break bonds within reactants

Must collide in the correct orientation to break bonds among reactants and allow the formation to create new products

If these are not satisfied = no reaction

Meeting E_a threshold

Minimum amount of energy for a reaction to occur (energy developed through collisions)

Particles are always constantly colliding (typically collisions have less energy than E_a) - the kinetic energy turns into potential energy

External stimulus (light/heat/sound) - can increase particle motion, leading to a greater collision - surpassing E_a for a reaction

Exothermic Vs. Endothermic (Unclear)

Endothermic: Products are higher energy than reactants, more external stimulus required for an appropriate collision

Exothermic: Products have a lower energy than reactants, so only a small amount of external stimulus is required

Endothermic Vs. Exothermic

Transition State/Activated Complex

The site of maximum potential energy within a reaction

Bond-breaking (reactants), bond-forming (products) is being done

Is highly unstable

Like a “hill” - can go back to being products, or go back to being reactants (determined by relative energies of the products and reactants and random molecular motion) - the activated complex is most likely to move towards the side with less energy

e.g. in exothermic - products have less energy - favors the forward reaction

in endothermic - products have more energy - more likely to fall back to reactants, unless sufficient energy is provided

In reversible reactions, and equilibrium may be formed

Equilibrium

Forward reactions = Backward reactions

Forward Reactions - Transition state to products

Backward Reactions - Transition state to reactants

On average, no. of forward reactions = no. of backward reactions

Orientation Of Colliding Molecules

The correct orientation of collision must occur - if not, no matter how great the activation energy, there will be no reaction

Must react in a specific orientation allowing the bonds in the reactants are broken and bonds in the products are formed (bonds must break entirely, not partially)

Finding Activation Energy Of Reverse Equation (and Enthalpy)

E_a of reverse = E_a of foward - Enthalpy

-(Enthalpy) = new Enthalpy of reverse equation

If a reaction is endothermic, its reverse is always exothermic (and vice versa)

How Rates Of Reactions Can Be Altered

Increasing surface area of solid reactant

Concentration of reactants in a solution (more collisions)

Pressure of gaseous reactants

Temperature of the reaction

Prescence of a catalyst

In order to increase reaction rate

Increase the frequency of successful collisions

Increasing the energy of all collisions so that the proportion of collisions that have energy higher than E_a is higher

Increasing concentration/Pressure

When concentration increases (more particles in a given area) then subsequently the frequency of collisions for which E (energy of collision) > E_a also increases

When adding more gas into a fixed volume or a smaller volume, it increases the concentration of gas molecules = more successful collisions

Increasing Surface Area

By expanding surface area (for a solid), more successful reactions can occur, as they only happen at the surface

More reactant particles are exposed, which results in more successful collisions, triggering a reaction

e.g. burning paper than a log - has more surface area so it catches fire, and burns easier

Higher oxygen concentration in shallow water causes more frequent successful collisions, increasing the rate of corrosion.

Common sense- oxygen drops when the water gets deeper

In shallow waters, water has a higher concentration to oxygen atoms and comes in contact with temperatures high than deep water, effectively increasing the rate of the reaction

Though more pressure is found in deep water, it is not enough to mitigate the higher concentration of oxygen atoms in shallow waters

Increasing Energy Of Particles

Only occurs through increase of temperature

Relation between kinetic energy and temperature

KE = 1/2mv²

When the temperature increases, v increases (v is the temp, m is the mass)

Since v is squared, a small increase can result in much more kinetic energy, resulting in more reactions

Greater proportion of collisions will have E> E_{a -} will have energy that is greater than of = to the required activation energy

Mxwell-Boltzmann Distribution Curve

Mean - the average of energies across all particles (is boosted by th particles in the high end)

Most probable energy point - The energy level that most particles will have

Increasing Temperature/ Adding Catalyst - Effect On Boltzmann Graph

When temperature increases, does activation energy decrease?

No - only proportion of particles that can overcome the Ea barrier increases - not the physical barrier itself

What happens if a catalyst itself has larger surface area

It provides more reaction sites between the particles and the catalyst - leads to accelerated reaction

Role Of Catalysts

Reduces required activation energy, and introduces an alternative pathway for the reaction to take place (less activation energy required), and lowers the energy level of the transition state

Normally reactants must get over a high activated complex, but through catalysts they can go through another pathway that reduces the height of “the hill” required

How Catalysts Lower Activation Energy

Bringing reactants together correctly

• Increases the frequency of effective collisions

• Ensures correct orientation for bond breaking/forming

🔹 Weakening existing bonds

• The catalyst may temporarily bond with reactants

• This weakens bonds, so less energy is needed to break them

🔹 Stabilising the transition state

• The transition state becomes lower in energy

• A lower-energy transition state = lower activation energy

Key Points To Catalystr Interaction

Activation energy decreases

Enthalpy remains the same (smaller bump in activation energy only)

Catalyst is not consumed (is still chemically unchanged, and is present in the mixture at the end)

Catalyst Effect

Types Of Catalysts

Homogenous Catalysts, Heterogenous Catalysts

Homogenous Catalysts

Are in the same form as the reactants (solid, liquid, gas)

Mixes uniformly with reactants

Creates temporary intermediates

Has good contact with reactants

Catalyst is regenerated in the end

Temporary Intermediates

When the catalyst makes bonds with the reactants for a very short time, and allow the reaction to occur in multiple smaller steps

Each step has lower activation energy - thus activation energy is lower

Are consumed before the reaction finishes (not in overall reactions)

Another explanation: A short-lived species formed during a reaction that exists between reactants and products, but is not the transition state. -Has lower energy than the transition state

Heterogenous Catalysts

Catalysts that have a different form to the reactants (Usually a solid when reactants are gases and liquids)

The reactants adsorb onto the catalyst and held onto the right position (correct orientation for a reaction to occur) - increase the surface area for possible collisions

No collisions occur, as reactants are held in place and their bonds are stretched and vulnerable to breakage

Then they desorb and go away as products

Higher surface area = more sites of adsorption

Catalysis

The process of employing a catlyst

Example of Homogenous Catlysis

Chlorine atoms acting as catalysts, breaking done ozone gas into oxygen gas - depleting the ozone layer

Why Heterogenous Catalysts Are Used For Industrial Purposes

More easily seperated from products of reaction - can be reused and recycled

Much easier to use

Able to be used at high temperatures

• Many homogeneous catalysts are molecular compounds (e.g., transition metal complexes).

• High temperatures can:

  • Break down the catalyst (decompose it chemically)

  • Change its oxidation state or structure, making it inactive

• They are dissolved in solution, so thermal stability is often low.

Result: The catalyst cannot survive long at high temperatures → reaction may slow or stop.

2. Heterogeneous catalysts

Definition: Catalysts in a different phase than reactants (usually solid with gas or liquid reactants).

Why they can withstand high temperatures:

• Typically metals or metal oxides → very thermally stable

• Structure doesn’t decompose easily

• Can maintain activity even at hundreds of °C

Summary Of How To Alter Reaction Rates

Why use a gas syringe

Used for accurately measuring collecting and dispensing volumes of mass

Why is cotton wool used on top of the container

It acts as a loose plug to allow gases to escape the vessel, whilst keeping the liquids within the vessel (preventing spraying or spilling)- especially during violent reactions

Surface Area Relationship

Larger the particle - surface area is smaller - smaller sites of reaction

Smaller the particle - Surface area is larger (e.g. powder)

SA = 1/particle size (inversely proportional - the larger the individual particle, the smaller it's surface area)