Acids and bases

Acids

Acids

Proton donors (H+)

Bases

Proton acceptors

Alkalis

A type of base that dissolves in water to form hydroxide ions

Acid + alkali →

Salt + Water

Acid + carbonate →

Salt + water + carbon dioxide

Acid + metal →

Salt + hydrogen gas

Monoproctic acids

Each molecule can release 1 ptoron

Strong acids dissociate…

Fully in water

General acid dissociation equation

HA + H2O → H3O+ + A-

Weak acids

Partially dissociate in water

What type of reaction are weak acid dissociations

Reversible

General equation for weak acid dissociation

HA + H2O → ← H3O+ + A-

Where does equilibrium lie in weka acid dissociations

Well over to the left

Ka=

Acid dissociation constant

What are the Ka values for weak acids

Less than 1

Ka equation

Ka= ([H+(aq)] [A-(aq)]) / [HA]

General acid dissociation equation

HA→ H+ + A-

What does a larger Ka indicate about PoE

Favours the right hand side- is a stronger acid

What does a lower pKa value indicate

A stronger acid

pKa=

-log10(Ka)

Ka using pKa =

Ka= 10^-(pKa)

[H+ (aq) ] =

10^-(pH)

pH=

-log10([H+(aq)])

What type of scale is the pH scale and what does this mean about each value

Is logarithmic- each value differs by a factor of 10

What does a high [H+] indicate about pH

A low pH

What would a pH change of 1 mean for the [H+]

It would change x10

What equations calculate the pH of strong acids

[H+(aq)] = [HA(aq)] then pH= -log10^[H+(aq)]

How t roughly check pH of strong acid calculations

Should be less than 4

Equations too work out pH of weak acids

[H+(aq)] = sq rt ( Ka x [HA(aq)] ) then pH= -log10([H+(aq)]

What assumptions do we make when calculating pH of weak acids

1) equilibrium constant of HA= conc. of undissociated HA

2) conc. of H+ ions are = A- ions

Amphoteric

Substance can act as both an acid and a base

Example of water acting as a base

H2O + HCl → H3O + + Cl-

Water acting as an acid

H2O + NH3 → NH4+ + OH-

Kc equation for water

Kc= [OH-][H+] / [H2O] → will be less than 1

Extent of ionisation of water

Very small- approx 1 water molecule. In every 500,000 dissociates

Dissociation reaction of water

H2O (l) → H+(aq) + OH-(aq)

Where does PoE lie on water dissociation reaction

Well to the left

Kw definition

The ionic product of water

What is Kw made up from

Kc and [H2O]

Kw=

Kc x [H2O(l)] = [H+(aq)][OH-(aq)]

At 25 deg c, pH of H2O is 7 what is [H+(aq)]

10^-7

At 25 deg c, pH of H2O is 7, what does [OH-(aq)] =

10^7 as [H+]=[OH-}

At 25 deg c, pH of H2O is 7, what is Kw

[H+]x[OH-]= 10^7 × 10^7 = 10^14, Kw= 1 × 10^-14 mol² dm^-6

In water and neutral solutions [H+(aq)] =

[OH-(aq)]

In acidic solution s [H+(aq)]

>[OH-(aq)]

In alkaline solution [H+(aq)]

< [OH-(aq)]

At 25 deg c, what must Kw=

1×10^14 mol² dm^-6

Dissociation of NaOH

NaOH→ Na+ + OH-

[NaOH]=

[OH-]

How to find the pH of NaOH

Kw= [H+][OH] therefore [H+]= Kw/[OH-] , pH= -log10([H+])

What is the pH of KOH with a conc. of 0.05 mol dm^-3

[KOH]=[OH-] = 0.05, Kw= 1×10^14, [H+]= Kw/[OH-} = 1×10^-14/ 0.05= 2×10^-13, pH= -log10(2×10^-13)= 12.7

Buffer solution

A mixture that minimises pH changes on addition of small amounts of acid or base

What are buffers made from

Weak acid (HA) and its conjugate base (A-)

When an acid (H+) is added to a buffer solution, what happens

[H+] increases, conjugate base A- reacts with the excess H+ ions, PoE shifts to the left removing most of the excess H+ ions

When an alkali (OH-) is added to a buffer solution what happens

[OH-} increases, the small conc. of H+ ions react with the OH- ions (H+ + OH- → H2O), HA dissociates shifting the equilibrium right to restore most of the H+ ions that have reacted

What is pH dependent on

Acid dissociation constan, Ka, of buffer solution, conc. ration of weak acid: conjugate base

Find pH at 25deg c of a buffer containing 0.06 mol dm^-³ CH3COOH and 0.1 mole dm^-³ CH3COO-Na+, fro CH3COOH, Ka= 1.7 × 10^- 5

[H+] = (1.7 × 10^-5) x (0.05/ 0.1) = 8.5 × 10^6, pH= -log10(8.5 × 10^6) = 5.08

Step by step method to work out pH of Buffers

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

[H+] = Ka x ([HA]/[A-])

pH=-log([H+])

Example o

Carbonic acid, hydrogen carbonate buffer in blood

How does the Carbonic acid, hydrogen carbonate buffer work

Blood pH has to be between 7.35-7.45, carbonic acid is the wea acid, hydrogen carbonate is the conjugate base, most materials released into blood are acidic so HC3O- ions are removed by being converted into H2CO3 which then gets converted to dissolved CO2 and removed by the lungs

Equation for Carbonic acid, hydrogen carbonate buffer

H2CO3 → ← HCO3- + H+

What happens when

an acidic/ alkaline substance is added

Equivalence point

When the solution of an acid- base titration has been neutralised

End point

When the colour changes

Features of a strong acid strong base pH curve

Low pH of starting acid, sharp rise in oh around equivalence point, plateaus at alkaline pH

Features of a strong acid weak base pH curve

Starts at pH 1, very little pH change in initial 20cm3, sharp change n pH over addition of less than ½ drop of alkali, curve levels of at pH 10 due to excess 0.1m of weak alkali

Features of weak acid strong base pH curve

Starts at pH4, sharp change in pH over addition of less than ½ drop of strong base, curve levels of at pH13

Feature of weak acid weak base pH curve

Starts a pH4, no sharp change in pH, curve leves of due to excess 0.1m weak alkali

How to choose a pH indicator

If thee ph range of the indicator falls within the rapid pH change of the titration

Indicators that can be used for strong acid strong base

Phenolphthalein, litmus, methyl orange

Indicator for weak acid strong base

Phenolphthalein

Indicator for strong acid weak base

Methyl orange

Indicator for weak acid weak base

No suitable indicator

pH range of colour change for methyl orange

3.5-5.5

pH range of colour change for Litmus

6.5-8.5

pH range of colour change for Phenolphthalein

8.5-10.5

What must acid base indicators have

An easily observable colour change, must change immediately in required pH range over the addition of half a drop of reagent

Lattice enthalpy definition

Measure of strength of ionic bonding in a giant ionic lattice

What is lattice enthalpy

The enthalpy change that accompanies the formation of one mole of an ionice compound form its gaseous ions under standard conditions

Features of lattice enthalpy

Involves ionice bond formation from espérate gases or ions, is exotérmica so value for enthalpy change will always by negative

Standard enthalpy change of formation

The change that takes place when one mole of a compound is formed form its elements under standard conditions with all reactants and products in their standard states

Standard enthalpy change of Atomisation

Enthalpy change that takes place for the formation of one mole of gaseous atoms form the elements in its standard state under standard conditions

Is AaeH endothermic or exothermic

Endothermic- bonds are broken to form gaseous atoms- is a positive value

First ionisation energy

Enthalpy change required to remove one electron from eat atom in one mole of gaseous atoms to form one mole of gasoues 1+ ions

Why are FIEs endothermic

Energy is required to overcome the attraction between a negative electrons and the positive nuclesu

First electron affinity

Enthalpy change that takes place when one electrons is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions

Why are FEAs exothemic

The electron being added is attracted towards the nucleus

When are Successive electron affinities required

Wen an anion has a charge greater than -1

How to write SEAs

Same as successive ionisation energies

Why are second electron affinities endothermic

A second electron is gained ya. Negative ion which repels the electron await so energy must be it in to force the negatively charged electrons onto the negative ion