Module 5 - CHEM Notes PDF

Preview

Summary

Download Module 5 - CHEM Notes PDF

Description

MODULE 5: EQUILIBRIUM AND ACID REACTIONS

 Qi saq u a nt i t yt ha tc h a n g e sa sar e a c t i ons y s t e ma p pr o a c he se q ui l i br i um.  Ki st hen ume r i c a lv a l u eo fQa tt h e" e n d "o ft her e a c t i on ,wh e ne q u i l i br i u m i sr e a c h e d REVERISIBITY OF REACTIONS Irreversible Reactions: Cannot be reversed and goes to completion. - Neutralisation, combustion, acid and metal reactions e.g. 2Mg(s) + O2(g)  2MgO(s) We cannot convert magnesium oxide back to Mg and O 2 Reversible Reactions: Reaction that can proceed in either direction depending on conditions. e.g. Iron (III) nitrate and potassium thiocyanate  blood-red complex iron(III) thiocyanate Fe3+(aq) + SCN-(aq) ⇌ [FeSCN]2+ Both the forward and reactions are occurring at that same time. The reactants do not all get used up and at any instant all species (reactants and products) are present in the reaction vessel. Adding extra iron ions or thiocyanate ions separately makes the colour a deeper red. Fe3+(aq) + SCN-(aq)  [FeSCN]2+ Adding extra iron cyanate to the mixture produces a lighter colour as reaction is pushed in reverse. [FeSCN]2+  Fe3+(aq) + SCN-(aq) REVERSIBLE AND IRREVERSIBLE REACTIONS All physical changes and some chemical reactions are reversible under suitable conditions. - The products can recombine and re-form reactants - Forward and reverse reactions will continue to occur at the same time OPEN AND CLOSED SYSTEM Open System: Both energy and matter are constantly moving between system and the environment. Closed System: Energy may still be able to flow in and out, between the system and the environment, but matter cannot enter or leave the system.  When a reaction reaches equilibrium, no energy enters ore leaves the system. Physical changes and Chemical changes can occur in both open and closed systems Some examples of reversible reactions:  Dinitrogen tetroxide: N2O4(g) ⇌ 2NO2(g)  Ionisation of water: H2O(aq) ⇌ H+(aq) + OH-(aq)



Living organisms: reactions involved in the transport of oxygen and carbon dioxide around the body in the blood: haemoglobin + O 2(g) ⇌ oxyhaemoglobin

STATIC VS DYNAMIC EQUILIBRIUM Equilibrium: Refers to a situation where no changes can be seen happening in a chemical system, which macroscopic properties are constant. STATIC: Where nothing is happening, the forward and reverse reaction rates are both zero DYNAMIC: There may seem to be no observable change taking place, but the forward and reverse reactions are occurring at the same rate and there is constant change at an atomic level. STEADY STATE: Physical system which matter and energy are entering and leaving the system at a constant rate so that conditions have becomes stable within the system. HOMOGENOUS EQUILIBRIUM: all involved species (products&reactants) are in the same state. HETEROGENOUS EQUILIBRIUM: Substances can be in different states CHARACTERISTICS OF EQUILIBRIUM SYSTEMS 1. It’s a closed system --- if a reversible reaction is to reach and maintain equilibrium no matter or energy can enter or leave the system. 2. Rate of forward reactions equals the rate of the reverse reaction. 3. Macroscopic properties stay constant --- no change in pH, state, colour, temperature or pressure 4. Concentrations of all reactants and products stay constant 5. The equilibrium can be approached from either direction ENTHALPY AND ENTROPY IN EQUILIBRIUM Entropy and enthalpy are both drivers of reactions such as photosynthesis and combustion. When ΔG = 0, the system is at equilibrium Therefore, when a system is in equilibrium, enthalpy and entropy is no longer considered as driving forces. EQUILIBRIUM AND COLLISION THEORY In a reversible reaction, the reaction in one direction is exothermic, the other is endothermic. Reactions go to completion when: 1. Activation energy is low, products have less energy than reactants 2. Activation energy is too high, reactants are more stable than products. A reversible reaction which is exothermic in the forward direction will become endothermic in the reverse reaction, vice versa. EQUILIBRIUM POSITION Refers to where the equilibrium lies with reference to the equation.

SO2(g) + O2(g) ⇌ 2SO3(g) If the equilibrium position is described as being to the left, that means there are higher concentrations of sulfur dioxide and oxygen than sulfur trioxide. If it lies to the right, then more sulfur trioxide is present at equilibrium. Equilibrium position can be affected by temperature, pressure and concentration. Catalyst has no effect but will just help the reaction reach equilibrium faster. INVESTIGATING CHANGES IN EQUILIBRIUM Cobalt Chloride Equilibrium [Co(H2O)6]2+(aq)(pink) + 4Cl-(aq) ⇌ [CoCl4]2-(aq)(blue) + 6H2O(l) Some   

changes you can make to this equilibrium reaction include: Chloride ion concentration can be increased Ethanol can be added to remove water from equilibrium The reaction mixture can be heated or cooled. o Higher temperatures, the mixture becomes a deeper blue o Low temperatures, mixture because more pink.

NO2 and N2O4 Equilibrium 2NO2(g)(brown) ⇌ N2O4(g)(colourless) ΔH = -57.2 kJ mol-1  Nitrogen dioxide can be heated or cooled and will change equilibrium position. The Chromate-dichromate Mixture 2CrO42+(aq) (yellow) + 2H+(aq) ⇌ Cr2O72-(aq) (orange) + H2O The Ferrocyanate Equilibrium Fe3+(aq) (yellow) + SCN-(aq) (colourless) ⇌

[FeSCN]2+(Blood-red)

Some changes you can make to this equilibrium reaction include:  Warming or cooling reaction changes equilibrium position  Adding dilute solution of Iron(III) Chloride CHANGING EQUILIBRIUM Equilibrium position can be disturbed by: Changing Concentration More reactants: drives forward reaction More products: drives reverse reaction Having more reactant particles will increase the frequency of collisions and some of these colliding particles will have energy levels equal to or higher than the activation energy allowing them to react.  therefore, equilibrium position changes Changing temperature Exothermic reaction: Increase temperature: drives reverse reaction Decrease temperature: drives forward reaction Endothermic reaction: Increase temperature: drives forward reaction Decrease temperature: drives reverse reaction Increase in temperature  favours the higher activation energy (endothermic)

Decrease in temperature  favours the lower activation energy (exothermic) Changing Pressure/Volume (on Qs) More moles of a gas can also increase pressure as particles are closer together. Increase in pressure: favour forward reaction  more collisions, increase rate of reaction Reduce in pressure: favour reverse reaction Pressure won’t have an effect when:  The number of moles of gas on either side of equation are equal  There are no gases present LA CHATELIER’S PRINCIPLE If a system is in equilibrium and it is disturbed or changed in any way, then the system will adjust itself to minimise the amount of change Changes in temperature: 2NO2(g) ⇌ N2O4(g) + heat (∆H is negative) If it is cooled, it will react to increase the temperature – to raise it to its original value, thus minimising the change. Increasing the temperature favours the endothermic direction of an equilibrium and decreasing the temperature favours the exothermic direction.

Changes in concentration: Cl2(g) (pale green) + ICl(l) (brown) ⇌

ICl3(s) (yellow)

If Cl2 concentration is increased the forward reaction will be favoured in order to use up the added Cl2. This decreases its concentration again – back to its original value. Changes in pressure for reactions involving gas CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g) If pressure was increased, the system will adjust by reducing the pressure to minimise change. There are fewer moles on the left side so we can predict that the reverse reaction will be favoured (the presence of fewer gas molecules means lower pressure). EQUILIBRIUM GRAPHS CO(g) + Cl2(g)

⇌

COCl2(g) (∆H is negative)

Changes in temperature Temperature has decreased at 7.5s. System is moving in a net forward direction to raise temperature to original value.

Reaches new equilibrium at 15s.

Changes in concentration Addition in CO concentration at 20s System responds by moving in a net forward direction to partially remove the added CO

Changes in pressure A sharp rise at 20s suggests a volume decrease (which causes a concentration increase in the CO) Moves in a net forward direction.

ADDING A CATALYST Adding a catalyst to a system at equilibrium will have no effect on the system. - Catalyst provides alternative pathways reducing activation energy, so both forward and reverse reactions will occur at faster rate. THE EQUILIBRIUM EXPRESSION AND CONSTANT aA + bB ⇌ cC +dD a, b, c, d are the number of moles of substances A, B, C and D that are present at equilibrium. Equilibrium constant (Keq): Defined as the product of the concentrations of products present at equilibrium.

K eq =

[ C]c [ D ]d [ A ]a [B ]b

1. Concentration values to calculate Keq are concentrations of substances present at equilibrium. 2. Concentrations of solids and pure liquids are not changed by adding to or removing them from an equilibrium system, so they are considered to have a concentration value of 1. 3. Keq stays the same for a particular equilibrium regardless of the concentrations you start with.

4. Keq not affected by: - Changes in concentration of pressure (only change equilibrium position) - The addition of a catalyst (only speeds up forward & reverse reactions) 5. Keq is only affected by temperature. When temperature increases: - Decreases for exothermic reaction (favours formation of reactants) - Increases for endothermic reaction (favours formation of products) RELEVANCE AND USES OF Keq VALUES: If Keq is large (e.g. >103), there is more products present in equilibrium than reactants. If Keq is small (e.g. Ksp

Supersaturated

THE COMMON ION EFFECT

Precipitation No Precipitation Solution is at equilibrium Precipitation Precipitation

Mixing two solutions with a common ion reduces the solubility of the ionic precipitate. Example Adding sodium chloride to a silver chloride precipitate at equilibrium (both have Cl- ion).  effect can be predicted by Le Chatelier’s principle Ag+ (aq) + Cl-(aq)

⇌ AgCl(s)

The more chloride ions added from sodium chloride, the equilibrium will shift to the right. This produces more AgCl salt and the silver chloride has become less soluble....