Moduel 5 Chapter 20
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What does the Arrehenius model define acids and bases as
The ions released when in an aqueous solution .Acids - H+ .Alkalis OH-
What does the Bronstead-Lowry model define acids and bases as
Their role in proton transfers, Acid are proton donors Bases are proton acceptors
What can protons only dissociate in and what do they form
They can only dissociate in solution, so they form hydroxonium ions (H3O +) with water
What is the hydroxonium ion often simplified to in equations
H+
What does a conjugate acid base pair contain
two species that can be interconverted by proton transfer
In a conjugate acid-base pair what categorises something being a conjugate base and a conjugate acid
When a species releases a proton it forms a conjugate base When a species accepts a proton it forms a conjugate acid
What are the three types of acid
Monobasic - donates one proton Dibasic - donates two protons Triassic - donates three protons
What can the reactions of acids be represented by
Ionic equations using H+ as the active species
Who produced the PH scale that we use today
Sorensen used the negative logarithm of H+ ion concentration to produce a more manageable score that we use today
What are the different H+ ion concentrations that correlate to the different PH values
H+ ion concentration of 10^-1 - 10^-14 converts to PH values of 1-14
What is the temperature at which the PH scale works at
25*C
What is the formula for PH
PH = -log[H+ (conc)]
How can the H+ ion concentrations be calculated from the formula
[H+ (conc)] = 1/10^ph
What is a feature of strong acids and how can this be used for working out PH
They fully dissacociate, so the concentration of the acid [HA] is equal to the concentration of [H+]
What are the equations for the disassociation of a strong and weak acid
HA —> H+ + A- HA <-> H+ + A-
Why can the constant Ka be used to represent the dissociation of a weak acid
As weak acids form an equilibrium their dissociation can be represented by the acid dissociation constant Ka
What is the formula for Ka
Ka = [H+ (aq)] [A- (aq)] / [HA (aq)]
What does the value of Ka change with
It changes with temperature and recorded Ka values are normally standardised at 25*C
What occurs as the value of Ka gets larger
The more the equilibrium lies to the right and so the lower the pH.
What is the formula for pKa
pKa = -log[Ka]
Why is pKa used
As the values for Ka are very small it is easier to give their negative logarithm pKa instead of
What are happens to the values of Ka and pKa as the acid gets weaker
Smaller Ka Larger pKa
What does the [H+] of a weak acid depend on
The value of Ka and [HA]
What are the two approximations that can be used to simplify the Ka equation for a monoprotic acid that is weak
HA dissociates to form equal amounts of [H+] and [A-] The change [HA] is negligible [HA (start)] = [HA (eqm)]