acids, bases and buffers
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Acids, Bases and Buffers. History and Definitions. Your task is to research the history of acids and bases. Doing this you will need to find out about Arrhenius, Bronsted & Lowry and Lewis. - PowerPoint PPT PresentationTRANSCRIPT
Acids, Bases and Buffers
History and Definitions
• Your task is to research the history of acids and bases.
• Doing this you will need to find out about Arrhenius, Bronsted & Lowry and Lewis.
• By the end of the lesson you will also need to definitions of acids and bases- with examples as equations.
Bronsted-Lowry
• Bronsted-Lowry acids and bases– A Bronsted-Lowry acid is any substance from
which a proton can be removed– A Bronsted-Lowry base is any substance that can
remove a proton from an acid
• A single proton doesn’t really exist in a solution. Acids only release protons if a base can accept it.
Conjugate Acid-Base Pairs
• Instead of a floating proton in solution, water molecules accept protons to form hydronium ions, H3O+
(aq)
• This is sometimes called an oxonium ion.
This is an acid
This is a thing that can accept a proton, it’s the acids conjugate baseWhich one is the base and conjugate
acid between these two?
An acid-base pair is a set of two species that transform into each other by gain or loss of a proton
Calculations
• Practice questions on page 139.
pH
pH
• Don’t ask what it means. Noone knows.
• pH is all about the concentration of hydrogen ions in solution.
• It is a logarithmic scale of concentration of hydrogen ions.
pH calculations
pH = -log[H+(aq)]
[H+] = 10–pH
Your Calculator
Calculations
• Attempt calculations on page 141 of text book.
Strong and Weak Acids
Strong Acids
• Strong acids completely dissociate in aqueous solution.
• Only a few exist, the rest are weak.– HCl -HI– HNO3 -HClO4
– H2SO4
– HBr
Weak Acids
• Weak acids only partially dissociate in aqueous solution, the equilibrium lies well to the left.
Ka The Acid Dissociation Constant
• A weak acid has the following equilbrium:HA H+ + A-
• The expression for the acid dissociation constant is:
Ka =
• Units are always:
Ka Context
• A strong acid has a high Ka value.• A weak acid has a small Ka value.
• Can also convert these into logs, which makes the numbers more manageable.
pKa = -log10Ka
Ka = 10-pKa
• Taking logs inverts the values. High pKa is a weak acid and vice versa.
pH of Strong Acids
• For a strong acid:HA(aq) H+
(aq) + A-(aq)
–HA totally dissociates: [HA] = [H+]–Use pH = -log[H+]
• A bottle of HCl has a concentration of 1.22 x10-3 mol dm-3. What is the pH?
pH of Weak acids
• For a weak acid:HA(aq) H+
(aq) + A-(aq)
–HA only partially dissociates.–H+ and A- are formed equally. [H+]=[A-] –In our equation for Ka: [H+][A-] = [H+]2
–Due to the small partial dissociation we can assume that the equilibrium concentration of HA is the same as the start concentration. This gives us the equation:
pH of Weak Acids
Ka = [H+]2
[HA]
Or
[H+] = Ka x [HA]
Weak Acid Practice
• A sample of nitric acid, HNO2, has the concentration 0.055 mol dm-3. Ka = 4.70 x10-4
mol dm-3 at 25oC. Calculate the pH.
• Ka =• [H]+ =• pH =
Kw
Ionisation of Water
H2O H+ + OH-
equilibrium
Kc = [H+][OH-]
[H2O]
Kw
• Rearranging this gives usKc x [H2O] = [H+] [OH-]
• [H2O] is always 55.6 mol dm-3
• Kc is a constant too. We can create a new constant Kw
• Kw = [H+] [OH-]
Significance of Kw
• At 25oC Kw = 1.00 x 10-14 mol2 dm-6
– This number is found on your data sheet.• This is due to the pH of water being 7.• So [H+] = [OH-] = 10-7
• This value of Kw lets us calculate [OH-] if we know [H+].
pH of Bases
pH of Bases
• Using the Kw equation we can calculate the pH of a strong base.
• We wont have to calculate pH for a weak base.
• The strength of a base is a measure of its dissociation in solution to produce OH- ions.– Exactly like acids.
pH Calculations
• To work out the pH of a base we need to know two things
• [H+]• Kw
• For a strong base it totally dissociates. [OH-] = [base]
Example
• A solution of KOH has a concentration of 0.050 mol dm-3. What is it’s pH?
• Kw = • [H+] = • pH =