Chapter 14

1
Chapter 14 Acids and Bases
2
History of Acids Bases

• Vinegar was probably the only known acid in
  ancient times.
• Strong acids such as sulfuric, nitric and
  hydrochloric acids were not discovered until
  after the 12th century.
• Over the years, there have been many attempts to
  define acids and bases.

3
Old Definitions of Acids and Bases

• At first, acids and bases were defined in terms
  of their observed properties such as taste,
  effects on indicators and reactions with other
  substances.
• In the 17th century, Boyle described the
  properties of acids in terms of taste, their
  action as solvents and how they changed colour of
  certain vegetable materials.
• He also noticed that alkalis (soluble bases)
  could reverse the effects of acids.
• Lavoisier, in the 18th century, thought that
  acidic properties were due to the presence of
  oxygen.
• In 1810, Davy suggested that the acid properties
  of substances were associated with hydrogen and
  not oxygen.
• In 1887, Arrhenius defined acids as substances
  that produced hydrogen ions (H) in water while
  bases produced hydroxide ions (OH-) in water.
• According to his theory, when acids and bases
  react together, the H and OH- form water
  according to the equation
• H OH- ? H2O Arrhenius called this a
  neutralisation reaction.

4
Definitions cont

• There were, however, limitations to these
  theories.
• Arrhenius definition for example was restricted
  to acids and bases in water.
• One of the more useful definitions used today was
  first proposed by the Bronsted and Lowry
• Bronsted and Lowry described the reactions of
  acids as involving the donation of a hydrogen ion
  (H).
• A hydrogen ion is a hydrogen that has lost its
  only electron.
• In most cases, a hydrogen ion is a proton.

5
Bronsted-Lowry Acids and Bases

• According to the Bronsted-Lowry theory, a
  substance behaves as an acid when it donates a
  proton, ie H to a base.
• A substance behaves as a base when it accepts a
  proton from an acid. Hence
• Acids are proton donors and
• Bases are protons acceptors.

6
Bronsted-Lowry Acids and Bases

• As protons are exchanged from an acid to a base,
  this definition explains why acids and bases
  react together.
• In an aqueous solution of hydrogen chloride,
  nearly all the hydrogen chloride is present as
  ions virtually no molecules of hydrogen
  chloride remain.
• This solution is known as hydrochloric acid.
• In this reaction, each hydrogen chloride molecule
  has donated a proton to a water molecule.
• According to the Bronsted-Lowry theory, the
  hydrogen chloride has acted as an acid.
• The water molecule has accepted a proton from the
  hydrogen from the hydrogen chloride, so has acted
  as a base.
• HCl(g) H2O(l) ? H3O(aq) Cl-(aq)

7
Acid-base Conjugate Pairs

• Because HCl and Cl- can be formed from each other
  by the loss or gain of a single proton, they are
  called a conjugate acid/base pair.
• Similarly, H3O and H2O are also a conjugate
  pair.
• A conjugate pair is two species which differ by a
  proton.
• For the reaction between HCl and H2O, the
  conjugate pairs are shown as
• HCl(g) H2O(l) ? H3O(aq) Cl-(aq)
• Blue bases
• Red acids

8
The H ion in Water

• A hydrogen ion (or proton) in solution is
  represented as H3O(aq) or more simply H(aq) and
  is called the hydronium ion.
• The hydronium ion itself attracts more water
  molecules and is further hydrated.
• However, these water molecules are not as
  strongly attracted and their number is not
  constant.

9
Some Common Acids Bases
10
Amphiprotic Substances

• Some substances can behave as either acids or
  bases, depending on what they are reacting with.
• These substances are given the name amphiprotic
  substances.
• In equation 1 below, water readily accepts a
  proton from sulfuric acid and acts as a base.
• In equation 2, water donates a proton to the
  oxide ion and acts as an acid.
• Eqn 1H2SO4(aq) H2O(l) ? HSO4-(aq) H3O(aq)
• Eqn 2 O2-(aq) H2O(l) ? OH-(aq) OH-(aq)

11
Amphiprotic Substances cont

• If the solute is a stronger acid than water, then
  water will act as a base.
• If the solute is a stronger base than water, then
  the water will act as an acid.

12
Amphiprotic Substances cont

• When an amphiprotic substance is placed in water,
  it reacts as both an acid and a base.
• For example, the hydrogen carbonate (HCO3-) ion
  reacts according to the equations
• HCO3-(aq) H2O(l) ? H2CO3(aq) OH-(aq)
• HCO3-(aq) H2O(l) ? CO32-(aq) H3O(aq)
• Since HCO3- can act as both acid and base, it is
  amphiprotic.
• Although both reactions are possible for all
  amphiprotic substances in water, generally one of
  these reactions occurs to a greater extent.
• The dominant reaction can be identified by
  measuring the pH of the solution.

13
Acid Base Strength

• Experiments show that different acid solutions of
  the same concentration do not have the same pH.
• Some acids donate a proton more readily than
  others.
• The strength of an acid is based on its ability
  to donate hydrogen ions.
• The strength of a base is based on its ability to
  accept hydrogen ions.
• Since aqueous solutions of acids and bases are
  most commonly used, it is convenient to use an
  acids tendency to donate a proton to water, or a
  bases tendency to accept a proton, as a measure
  of its strength.

14
Strong Acids

• Acids that ionise completely in solution are
  called strong acids.
• Strong acids donate protons easily.
• Solutions of strong acids would contain ions and
  virtually no unreacted acid molecules.
• The most common strong acids are hydrochloric
  acid, sulfuric acid and nitric acid.

15
Weak Acids

• An acid that does not fully ionise is called a
  weak acid.
• An example of a weak acid is ethanoic acid.
• Only a small proportion of ethanoic acid
  molecules are ionised.
• A weak acid can be shown be the presence of
  reversible arrows.
• CH3COOH(l) H2O(l) CH3COO-(aq) H3O(aq)

16
Strong Bases

• The ionic compound sodium oxide (Na2O)
  dissociates in water, releasing sodium ions (Na)
  and oxide ions (O2-).
• The oxide ions react completely with the water,
  accepting a proton to form hydroxide ions (OH-).
• The oxide ion is an example of a strong base.
• Strong bases accept protons easily.

17
Weak Bases

• Ammonia is a covalent molecular compound that
  ionises in water by accepting a proton.
• This ionisation process can be represented by the
  equation
• NH3(aq) H2O(l) NH4(aq) OH-(aq)
• Only a small proportion of ammonia molecules
  ionise.
• This is shown in the equation by the presence of
  reversible arrows.
• Ammonia is a weak base in water.

18
Polyprotic Acids

• Some acids are capable of donating more than one
  proton from each molecule and are said to be
  polyprotic.
• The number of hydrogen ions an acid can donate
  depends on the structure of the acid.
• Monoprotic acids can donate only one proton and
  include HCl, HF, HNO3, CH3COOH.
• Diprotic acids can donate two protons and
  include H2SO4, H2CO3,
• Triprotic acids can donate three protons and
  include H3PO4, H3BO3.

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Polyprotic Acids cont

• Polyprotic acids do not donate all protons at
  once, but do so in steps when reacting with a
  base.
• Sulfuric acid (H2SO4) is diprotic, meaning it has
  two protons that it can donate to a base.
• A diprotic acid ionises in two stages, for
  example
• STAGE 1 H2SO4(l) H2O(l) ? HSO4-(aq) H3O(aq)
• STAGE 2 HSO4-(aq) H2O(l) ? SO42-(aq) H3O(aq)

20
Polyprotic Acids cont

• When added to a base stronger than water, a weak
  acid will ionise to a greater extent.
• For example, a strong base such as OH- will
  accept a second proton from H2SO4 and the second
  and third proton from H3PO4.
• Similarly a weak base will ions to a greater
  extent if added to a strong acid.
• Sometimes there are more hydrogens in a molecule
  than can actually be donated.
• For example CH3COOH contains four hydrogen and
  yet will only donate one.
• Only the hydrogen involved in the polar OH- bond
  is donated.
• In general each hydrogen ion that is donated by
  an acid molecule is involved in a polar bond.

21
Relative Strengths of Acid Base Pairs
22
Strength vs. Concentration

• It is important that the terms strong and weak
  are not confused with the terms concentrated and
  dilute.
• Concentrated and dilute describe the amount of
  acid or base dissolved in a given volume of
  solution.
• The terms strong and weak describe how readily an
  acids donates, or base accepts a proton.

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Strength vs. Concentration cont
24
Qualitative vs. Quantitative

• Terms such as concentrated and dilute, or weak
  and strong are qualitative, or descriptive terms.
• Solutions can be more accurately described by
  stating concentration in mol/L or g/L.
• This is a quantitative description.

25
Acidic, Basic and Neutral Solutions

• The acidity of a solution is a measure of the
  concentration of hydrogen ions present.
• The higher the concentration of hydrogen ions,
  the more acidic the solution.
• Water has the ability to act as either an acid or
  a base.
• Pure water undergoes self ionisation to a small
  extent with allows it to conduct electricity
  slightly.
• This can be represented by the equation
• H2O(l) H2O(l) H3O(aq) OH-(aq)

26
Acidic, Basic and Neutral Solutions cont

• Acidic solutions contain a greater concentration
  of H3O than OH-.
• Neutral solutions contain equal concentrations of
  H3O and OH-.
• Basic solutions contain a lower concentration of
  H3O than OH-.

27
Measuring Acidity

• H3O x OH- 10-14M2
• Pure water is neutral so H3O OH-
• If either the H3O or OH- in an aqueous
  solution is increased, the other must decrease
  proportionally.
• At 25C, a solution is
• Acidic if H3Ogt10-7M and OH-lt10-7M
• Neutral if H3O 10-7M OH-
• Basic if H3Olt10-7M and OH-gt10-7M

28
Acidity Example

• In a 5.6x10-6M HNO3, solution at 25C, calculate
  the concentration of
• H3O ions
• HNO3 is a strong acid and ionises completely to
  produce 5.6x10-6M of H ions.
• b. OH- ions
• H3O x OH- 10-14
• 5.6x10-6 x OH- 10
• OH- 10-14/5.6x10-6
• OH- 1.79 x 10-9M

29
The pH Scale

• This scale is a useful way of indicating the
  acidity of a solution.
• pH -log10H3O
• The pH of a solution decreases as the
  concentration of hydrogen ions increases.
• Acidic solutions have a pHlt7
• Basic solutions have a pHgt7
• Neutral solutions have a pH7

30
Calculating pH Example 1

• What is the pH of a solution in which H
  0.0135M
• pH -logH
• pH -log(0.0135)
• pH -(-1.87)
• pH 1.87

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Calculating pH Example 2

• What is the pH of a 0.0050M of Ba(OH)2?
• Step 1 Find concentration of H
• Ba(OH)2(aq) ? Ba2(aq) 2OH-(aq)
• Ba(OH)2 is completely dissociated in water and
  each mole of Ba(OH)2 dissociates to release 2
  moles of OH- ions
• So, OH- 2 x Ba(OH)2
• 2 x 0.0050
• 0.010M
• Since H x OH- 10-14
• H x 0.010 10-14
• H 10-14 / 0.010
• H 10-12
• Step 2 Calculate the pH
• pH -logH
• -log(10-12)
• 12

32
Calculating the Concentration of H in a solution
of a given pH

• H 10-pH
• If the pH is 5.00, what is the H?
• H 10-5
• 0.0001M