THEME 2: THE CHEMISTRY OF ACIDS AND BASES
BASICS OF ACIDS AND BASES TO REMEMBER
‐ Acids:
o Corrosive
o Sour
o React with carbonates
o React with some metals to form H2
o Change colour of indicators
o Changes blue litmus red
‐ Bases:
o Corrosive
o Bitter
o Soapy
o Change colour of indicators
o Changes red litmus blue
16.1 THE BRONSTED-LOWRY CONCEPT OF ACIDS AND BASES
‐ Arrhenius acid: substance that, when dissolved in water, increases the concentration
of hydrogen ions, H+
‐ Arrhenius base: substance that increases the concentration of hydroxide ions, OH -,
when dissolved in water
‐ Bronsted-Lowry acid: is a proton (H+) donor
‐ Bronsted-Lowry base: is a proton acceptor
‐ Wide variety of Bronsted-Lowry acids such as molecular compounds such as nitric
acid, cations such as NH4+, anions such as HSO4- and hydrated metal cations
‐ Bronsted-Lowry bases include molecular compounds and anions
‐ Acids such as HF, HCl, HNO3 and CH3CO2H (acetic acid) are capable of donating one
proton and are called monoprotic acids
‐ Other acids calls polyprotic acids → able to donate 2 or more protons
Polyprotic Acids and Bases
Acid form Amphiprotic form Base form
H2S HS- S2-
H3PO4 H2PO4- PO43-
HPO42-
H2CO3 HCO3- CO32-
H2C2O4 HC2O4- C2O42-
‐ Such things as polyprotic acids (donate more than one proton) and polyprotic bases
(accept more than one proton)
‐ Some molecules and ions can act as a Bronsted acid or base → Amphiprotic – e.g.
water
Images with examples come from Mr de Beers class notes
,CONJUGATE ACID-BASE PAIRS
‐ Reaction of a Bronsted acid and base produces a new
acid and base
‐ Refer to pg 711 for a diagram of conjugate acid and base
pairs
‐ A conjugate acid-base pair consists of two species that
differ from each other by the presence of a hydrogen ion
‐ Every reaction between a Bronsted acid and base involves 2 conjugate acid-base
pairs
Images with examples come from Mr de Beers class notes
, 16.2 WATER AND THE pH SCALE
WATER AUTOIONISATION AND THE WATER IONISATION CONSTANT, Kw
‐ Water ionisation is reactant-favoured at equilibrium
‐ Kw = [H3O+][OH-] = 1.0 x 10 -14 at 250C
‐ Important aspects to this equation:
o Equilibrium constant is given a special symbol (Kw) and is known as the
autoionization constant of water
o Electrical conductivity of pure water shows that [H3O+] = [OH-] = 1.0 x 10-7 M
o Do not include the concentration of water in the expression for Kw
‐ If acid or base added to pure water, equilibrium is disturbed
o Adding an acid raises concentration of H3O+ ions – to oppose the increase, Le
Chatelier’s principle predicts that a small fraction of these ions will react with
OH- ions from water autoionization to from water. This lowers [OH-] until the
product of [H3O+] and [OH-] is equal to 1.0 x 10-14
o Adding a base to pure water gives a basic solution because OH- ion
concentration has increased. Le Chateliers principle predicts some of the
added OH- ions will react with H3O+ ions present in the solution from water
autoionization, lowering [H3O+] until the value of the product of [H3O+] and
[OH-] equals 1.0 x 10-14
‐ In a neutral solution, [H3O+] = [OH-] → both equal 1.0 x 10-14 M
‐ In an acidic solution, [H3O+] > [OH-] → [H3O+] > 1.0 x10-14 M and [OH-] < 1.0 x 10-14
‐ In a basic solution, [H3O+] < [OH-] → [H3O+] < 1.0 x 10-14 M and [OH-] > 1.0 x 10-14
THE pH SCALE
‐ pH scale compresses range of concentrations to values from roughly 15 to -1
‐ pH of solution defined as the negative of the base-10 logarithm (log) of the
hydronium ion concentration
‐ pH = -log [H3O+]
‐ can define pOH as the negative of the base-10 logarithm of the hydroxide ion
concentration
‐ pOH = -log [OH-]
‐ pure water → hydronium and hydroxide ion concentrations both 1.0 x 10-7 M –
therefore at 250C pH = -log (1.0x10-7) = 7.00
‐ pKw = 14.00 = pH + pOH
‐ don’t forget !! When working with log – the number of significant figures of the
number is the number of decimal places in the pH answer
‐ figure 16.1 pg 714 Is very important
Images with examples come from Mr de Beers class notes
BASICS OF ACIDS AND BASES TO REMEMBER
‐ Acids:
o Corrosive
o Sour
o React with carbonates
o React with some metals to form H2
o Change colour of indicators
o Changes blue litmus red
‐ Bases:
o Corrosive
o Bitter
o Soapy
o Change colour of indicators
o Changes red litmus blue
16.1 THE BRONSTED-LOWRY CONCEPT OF ACIDS AND BASES
‐ Arrhenius acid: substance that, when dissolved in water, increases the concentration
of hydrogen ions, H+
‐ Arrhenius base: substance that increases the concentration of hydroxide ions, OH -,
when dissolved in water
‐ Bronsted-Lowry acid: is a proton (H+) donor
‐ Bronsted-Lowry base: is a proton acceptor
‐ Wide variety of Bronsted-Lowry acids such as molecular compounds such as nitric
acid, cations such as NH4+, anions such as HSO4- and hydrated metal cations
‐ Bronsted-Lowry bases include molecular compounds and anions
‐ Acids such as HF, HCl, HNO3 and CH3CO2H (acetic acid) are capable of donating one
proton and are called monoprotic acids
‐ Other acids calls polyprotic acids → able to donate 2 or more protons
Polyprotic Acids and Bases
Acid form Amphiprotic form Base form
H2S HS- S2-
H3PO4 H2PO4- PO43-
HPO42-
H2CO3 HCO3- CO32-
H2C2O4 HC2O4- C2O42-
‐ Such things as polyprotic acids (donate more than one proton) and polyprotic bases
(accept more than one proton)
‐ Some molecules and ions can act as a Bronsted acid or base → Amphiprotic – e.g.
water
Images with examples come from Mr de Beers class notes
,CONJUGATE ACID-BASE PAIRS
‐ Reaction of a Bronsted acid and base produces a new
acid and base
‐ Refer to pg 711 for a diagram of conjugate acid and base
pairs
‐ A conjugate acid-base pair consists of two species that
differ from each other by the presence of a hydrogen ion
‐ Every reaction between a Bronsted acid and base involves 2 conjugate acid-base
pairs
Images with examples come from Mr de Beers class notes
, 16.2 WATER AND THE pH SCALE
WATER AUTOIONISATION AND THE WATER IONISATION CONSTANT, Kw
‐ Water ionisation is reactant-favoured at equilibrium
‐ Kw = [H3O+][OH-] = 1.0 x 10 -14 at 250C
‐ Important aspects to this equation:
o Equilibrium constant is given a special symbol (Kw) and is known as the
autoionization constant of water
o Electrical conductivity of pure water shows that [H3O+] = [OH-] = 1.0 x 10-7 M
o Do not include the concentration of water in the expression for Kw
‐ If acid or base added to pure water, equilibrium is disturbed
o Adding an acid raises concentration of H3O+ ions – to oppose the increase, Le
Chatelier’s principle predicts that a small fraction of these ions will react with
OH- ions from water autoionization to from water. This lowers [OH-] until the
product of [H3O+] and [OH-] is equal to 1.0 x 10-14
o Adding a base to pure water gives a basic solution because OH- ion
concentration has increased. Le Chateliers principle predicts some of the
added OH- ions will react with H3O+ ions present in the solution from water
autoionization, lowering [H3O+] until the value of the product of [H3O+] and
[OH-] equals 1.0 x 10-14
‐ In a neutral solution, [H3O+] = [OH-] → both equal 1.0 x 10-14 M
‐ In an acidic solution, [H3O+] > [OH-] → [H3O+] > 1.0 x10-14 M and [OH-] < 1.0 x 10-14
‐ In a basic solution, [H3O+] < [OH-] → [H3O+] < 1.0 x 10-14 M and [OH-] > 1.0 x 10-14
THE pH SCALE
‐ pH scale compresses range of concentrations to values from roughly 15 to -1
‐ pH of solution defined as the negative of the base-10 logarithm (log) of the
hydronium ion concentration
‐ pH = -log [H3O+]
‐ can define pOH as the negative of the base-10 logarithm of the hydroxide ion
concentration
‐ pOH = -log [OH-]
‐ pure water → hydronium and hydroxide ion concentrations both 1.0 x 10-7 M –
therefore at 250C pH = -log (1.0x10-7) = 7.00
‐ pKw = 14.00 = pH + pOH
‐ don’t forget !! When working with log – the number of significant figures of the
number is the number of decimal places in the pH answer
‐ figure 16.1 pg 714 Is very important
Images with examples come from Mr de Beers class notes
