Chapter 7 - Extent Of Chemical Reactions

Closed System

Matter and energy can be exchanged with surroundings

e.g. A steam radiator only exhales heat (energy), whilst water is held within (matter)

Open System

Energy and matter is exchanged with surroundings

e.g. Boiling pot with water - heat (energy) and water (matter) is being exchanged

Non-reversible/Irreversible reactions

Products cannot be converted back into reactants (e.g. combustion reactions)

How reversible arrows are depicted

With double arrow

The collisions between particles in products can reform reactants

Examples of reversible batteries

Recchargeable batteries

Reversible Reactions Reaching Equilibrium

The rate at which the reactants keep turning into products = rate at which products keep turning into reactants

Rate of forward reaction = rate of reverse reaction

Concentrations of all products and reactants remain constant

Reversibility

If endothermic reaction (larger activation energy), reverse is exothermic (smaller activation energy) (and vice versa)

Dynamic State Of Equilibrium

Forward and reverse reactions continue to occur

Reaction is incomplete (reactants and products are all in the mixture)

Bonds are constantly being broken and new bonds are being formed (reactants are turning into products and vice versa)

Occurs only in a closed system

Example

No further change between 8 seconds - 3hrs (dynamic equilibrium)

Extent Of The Reaction

How far the reversible reaction has been converted into products/how much product is formed when the system reaches equilibrium

Is expressed as a fraction of the reactants that have been converted into products

Is from 0-1 (no products - full conversion to producst)

What Does Extent Of Reaction Not Indicate

Speed of reaction

Can range from being fast - slow (not directly related to extent of reaction)

Example Of Extent Of Reaction (Unclear)

Example Of Concentration Graph

Equilibrium Law

It states if more products are being produced, more reactants are being produced, or if rate of production is equal

Equation for Equilibrium Law

Explaining Law

Products^coefficients /Reactants^coefficients

If there are more than one product or reactants, you must multiply them together

K is the equilibrium constant (K_c or K_p)

When multiplying products or reactants

Are in variable form - e.g. “x” or “M”

This is for their units - deciding what unit output is at the end, not the actual concentration itself

E.g

Reaction Quotient

Same as the equilibrium law, but is compared with “Q”

If K>Q, then more products are formed (shifted to the right)

If K<Q, then more reactants are formed (Shifted to the left)

If K=Q, then products and reactants are equal (equilibrium)

Q (Where reaction is at now), K (Where reaction should be) - theorised value

K vs Q Relation

Explanation Of K vs Q

If

K>Q (right now)

that means:

• The equilibrium ratio requires more products

• The system currently has too few products (or too many reactants)

So the reaction shifts forward to increase the numerator (products) and decrease the denominator (reactants).

Other way: too few reactants - goes backward to reach equilibrium

Homogenous System

Where all reactants and products are in the same phase

Pure Solids and Pure Liquids In Reaction Quotient

Have constant concentrations - adding them to the equation wouldn’t result in a change - hence they are left out for simplicity

Equilibrium Yield

Amount of products present at equilibrium

Meaning Of Concentrations (and their value)

If K_c is very small

The reactants (denominator) is very large compared to the products

If K_c is very large

The products (numerator) is very large compared to the reactants

How does Temperature only affect K_c

It changes how many particles have enough energy to react (more kinetic energy)

It affects the forward and reverse reactions differently

It changes the relative rates at which forward and reverse reactions occur

When coefficients are halved

K is square rooted (from original)

When coefficients are doubled

K is squared (from original)

When doing changes to constant “K”, does that affect the units “M”?

No - they must be reworked with the equilibrium law

Example Of Using ‘Ice Box’ Method when concentrations are unknown

Le Chatlier’s Principle

If the system if offset from equilibrium by some force, it itself will adjust accordingly to attain equilibrium

Equilibrium will shift to relieve stress that it is placed under

Equilibrium Position

The actual amounts of products and reactants (K is simply a ration - don’t get them mixed up!)

Changes To An Equilibrium System

Adding/Removing a reactant/product

Changing pressure by changing the volume (for gases)

Dilution (for aqeuous solutions)

Changing Temperature

Adding a reactant/product

When more reactant is added, the system is temporarily not in equilibrium

Collisions between reactants occur more frequently - creates more product

But the products also collide (more particles - more collisions) to produce reactants

Equilibrium is slowly achieved (backward = forward)

Even if one is altered, it affects all concentrations

Important Note

After changes in the conditions, the system does not return to the initial equilibrium position

K value is not changed (unless temperature is changed)

All concentrations don’t go up - if reactant is added, it favors the forward reaction more (more reactants are consumed), despite the products being formed back into reactants

e.g. If one reactant is added, it favors the forward reaction, but the other reactants go down in concentration (as they were never added) - net concentration goes down

Forward Reaction Vs Backward Reaction - When product vs reactant is added

Predicting effect using equilibrium law

When more reactants are added - denominator increases - and Q is momentarily < K

This triggers a production in products (position shifts to the right)

When more products are added - numerator increases - Q is temporarily > K

More reactants are produced (Position shifts to left)

Changing Pressure by changing Volume

If pressure is higher - it favors the reaction which creates less particles, thus reducing the overall pressure

If pressure is lower - it favours the reaction which creates more particles, thus increasing the pressure

e.g. 2SO_2(g) + O_{2 (g)} - 2SO_3(g)

When there is more pressure, it favours the forward reaction and as three gas particles turn into 2 - reduces pressure (equilibrium position shifts)

Applying equilibrium law to explain change of pressure

By halving the volume - all concentrations are doubled

Multiplying “2” to the power of their coefficient = to a factor of K

Pressure Changes For Liquids & Solids

They are already too tightly packed - an increase in pressure has a negligible effect on volume

When number of particles in reactants and products are equal

No matter in which direction the system shifts, the output is the same

Hence it does not oppose a volume change

When volume is doubled

The concentration of the system is halved

The system shifts to wherever more particles are being produced

Increasing pressure with an inert gas

By adding an inert gas, the overall pressure increases, but the concentration of the reactants/products is unaffected (are not squashed)

If the container is not rigid, the volume expands - lessening pressure