# 5.1 Rates Equilibria and pH

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• Created by: ryan
• Created on: 18-12-12 17:49

## 1.1Rate graphs and orders

Rate of reaction:The change in the amount of reactants or products per unit time                               (per second).

order:How the reactants concentration affects the rate.

•       shows the number of molecules of that reactant which are                                   involved in the rate determining step

rate constant:The larger it is the faster the reaction.

half-life:The time for half the reactant to disappear.                                                                  a first order reaction's half life is independent of the concentration

rate-determining step: The slowest step in a multi step reaction.

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## Rate equations; rate constants

K=[A]^m[B]^n

m and n are orders 0,1 or 2

Temperature changes:

reactions occur when particles collide with enough energy to break existing bonds. By increasing the temperature the reactant particles speed up, so they collide more often, with more of a chance of reacting as they have more energy.

So, Increasing temperature increases the reaction rate.

According to the rate equation, reaction rate is affected only by the rate constant and the reactant concentrations.  therefore, it must change the rate constant.

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## Rate determining step

Using a rate determining step it is possible to create a rate equation. for example by looking at the mechanism between chlorine free radicals and ozone.

As both chlorine free radicals and ozone are in the rate determining step they must both be in the rate equation.

As there is only one Chlorine free radical and one Ozone molecule they must both be to the order 1.

Making the equation:

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## 1.2 Equilibrium

Dynamic equilibrium can only happen in a closed system at a constant temperature.

H2+I2 <=>2HI

If 1.0 mole of Hydrogen gas is mixed with 1.0 mole of Iodine gas at a constant temperature 640K. At equilibrium, if there is 1.6 moles of Hydrogen Iodide there will be 0.2 moles of the starting gases, no matter how long you leave it.the equilibrium amounts never change.

Kc

by using the molar concentrations of each substance at equilibrium it is possible to work out the equilibrium constant, Kc, at a particular temperature.

aA+bB<=>dD+eE,    Kc=[D]^d[E]^e/[A]^a[B]b

products go on the top line [] mean concentration in moldm^-3

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## more on Kc

Temperature changes alter Kc

increasing the temperature in an exothermic reaction (-dH) will shift the reaction in the endothermic direction (+dH) to absorb the extra heat.

Decreasing the reaction temperature will shift the equilibrium in the exothermic direction to replace the lost heat.

If the change means more product is formed Kc will increase. if it means less is formed, Kc will decrease.

Concentration, pressure changes will cause the equilibrium to counteract the change meaning Kc will remain the same.

A catalyst will not affect Kc, it will just ensure equilibrium is reached faster.

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## 1.3 Bronsted-Lowry acids and bases

Acid - A species that can donate a proton                                                                     Base - A species that can accept a proton

Metal atoms will become oxidised by donating electrons to H+ ions, reducing them, Hydrogen gas is given off.

Carbonates react with an acid forming water and carbon dioxide.

Acids and bases will react and neutralise each other.                                                    Most insoluble bases are metal oxides and are neutralised too.

when an acid is added to water  the acid HA donates a proton to the water, in the reverse reaction however, the A- acts as a base and accepts. Water in this case is a base as it accepts a proton.

When a base is added to water, the water donates a proton so can also be considered an acid.

The conjugate pairs for both instances will be the acids with their corresponding bases

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## strong and weak acids

A strong acid/base is one which will dissociate completely.

A weak acid/base is one which will dissociate negligibly.

To find the pH of a weak acid you use Ka (the acid dissociation constant).

Ka=[H+][A-]/[HA]

as all H+ ions come from the acid you can make [H+]=[A-] making Ka=[H+]^2/[HA]

pKa=-log(Ka)    Ka=10^-pKa

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## pH and [H+(aq)]

pH=-log[H+]=10^-pH

for strong monobasic acids [H+]=[Acid]

For strong bases [OH-]=[Base]

Kw=[H+][OH-] (mol^2dm^-6)

using the Kw expression the pH of a strong base can be worked out, by taking the concentration of the base as the concentration of OH- ions you can calculate [H+] by rearranging the Kw equation.

[H+]=Kw/[OH-] then taking the -log of [H+]

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## Buffers

• A buffer is a system that minimises pH changes on the addition of a small amount of acid or base.
• if an acid is added the H+ concentration increases. most of the extra H+ ions combine with the B- ions to form HB. Shifting the equilibrium to the left.