Acids and Bases

1
Acids and Bases
2
Brønsted-Lowry Theory

• Brønsted-Lowry describes 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.
• Chemists often use the terms hydrogen ion and
  protons interchangeably.

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Brønsted-Lowry Theory

• According to Brønsted-Lowry theory, a substance
  behaves as an acid when it donates a proton to a
  base.
• In other words it donates a H to a base.
• A substance behaves as a base when it accepts a
  proton from an acid.

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HENCE

• Acids are proton donors
• Bases are protons acceptors

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Brønsted-Lowry Theory

• As protons are exchanged from an acid to a base,
  this definition explains why acids and bases
  react together.
• For example
• Hydrochloric acid is very soluble in water. The
  molecules ionise in water.
• HCl(g) H2O(l) ? H3O(aq) Cl-(aq)
• In an aqueous solution of hydrogen chloride,
  nearly all the hydrogen chloride is present as
  ions, virtually no molecules of HCl remain. This
  ionised solution is the hydrochloric acid we use
  regularly.

6
Hydrogen Chloride

• In the reaction each hydrogen chloride molecule
  donated a proton to the water molecule.
• According to the Brønsted-Lowry theory is it an
  acid or a base?
• The water molecule has accepted a proton.
• So is water an acid or a base?

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Conjugate Pairs

• HCl and Cl- can be formed from each other by the
  loss or gain of a single H (proton).
• These are called conjugate acid/base pairs.
• Similarly H3O and H2O are also a conjugate pair.
  
• A conjugate pair is two species which differ by a
  proton.

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Conjugate Pairs

• What are the conjugate pairs in this reaction.
• NH3(aq) H2O(l) ? NH4(aq) OH-(aq)

Base
Base
Acid
Acid
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Some Common acids and bases
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Amphiprotic Substances

• Some substances can be acids or bases depending
  on what they react with.
• They can donate or accept protons.
• These substances are said to be amphiprotic.
• Can you name any amphiprotic substances?

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Water

• Water is an amphiprotic substance.
• It can be an acid and a base.
• 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
  water will act as an acid.

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Amphiprotic Substances
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Acid and Base Strength

• 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 or a base is its ability
  to donate or accept an proton.
• We generally use an acids tendency to donate a
  proton to water or a bases tendency to accept a
  proton from water, as a measure of its strength.

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Strong Acids

• Acids that ionise completely in solution are
  called strong acids.
• Solutions of strong acids would contain ions,
  with virtually no unreacted acid molecules
  remaining.
• HCl(g) H2O(l) ? H3O(aq) Cl-(aq)
• H2SO4(aq) H2O(l) ? H3O(aq) HSO4-(aq)
• HNO3(aq) H2O(l) ? H3O(aq) NO3-(aq)

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Weak Acids

• In water it ionises to produce ethanoate ions and
  hydronium ions.
• CH3COOH(l) H2O(l) CH3COO-(aq)
  H3O(aq)
• However only a small proportion (less than 1) of
  the ethanoic acid molecules actually ionise.
• So in water more is present as CH3COOH than
  CH3COO-
• We use a reversible arrow to represent a weak
  acid

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Strong and weak bases

• A strong base dissociates completely in water,
  all of the compound is now in the form of ions.
  Hydroxide ions are a strong base
• A weak base does not dissociate completely in
  water. Ammonia is a weak base.
• We also represent weak bases by the reversible
  arrows.

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Polyprotic acids

• Acids that are capable of donating more than one
  proton are polyprotic.
• Monoprotic acids can donate only one proton
• These include HCl, HF, HNO3, CH3COOH
• Diprotic acids can donate two protons
• Sulfuric acid H2SO4 and carbonic acid H2CO3 are
  diprotic acids
• Triprotic acids can donate three protons.
• Phosphoric acid H3PO4 is a triprotic acid

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Polyprotic acids

• Polyprotic acids do not donate all their protons
  at once, but do so in steps when reacting with a
  base.
• It also depends on the strength of the acid.
• Sulfuric acid (H2SO4) is diprotic.
• A diprotic acid ionises in two stages.

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Stage 1

• H2SO4(aq) H2O(l) HSO4-(aq)
  H3O(aq)
• Sulfuric acid is a strong acid in water so this
  stage occurs to completion.
• That is all the sulfuric acid molecules have
  ionised into hydrogen sulfate and hydronium ions.

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Stage 2

• HSO4-(aq) H2O(aq) SO42-(aq)
  H3O(aq)
• Hydrogen sulfate is only a weak acid so only a
  proportion ionise.
• A solution of sulfuric acid therefore contains
  hydrogen ions, hydrogen sulfate ions and sulfate
  ions.

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Strength versus Concentration

• Strong and weak refer to acids.
• They are not the same as concentrated and dilute.
• Concentrated and dilute describe the amount of
  acid or base dissolved in a given volume of
  solution.

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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.
• Quite often we use the H3O instead of the
  hydrogen ion.

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Water

• Water is both an acid and a base
• Pure water undergoes self ionisation to a very
  small extent.
• H2O(l) H2O(l) H3O(aq) OH-(aq)
• Water behaves as a very weak acid and a very weak
  base, producing one hydrogen ion (H3O) for every
  hydroxide ion (OH-).

Acid
Base
Acid
Base
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Acidic Solutions

• Pure water is neutral because the concentration
  of H3O ions is equal to the concentration of OH-
  ions present.
• If an acid is added to water, more H3O ions are
  produced. The concentration of H3O ions becomes
  greater than that of OH- ions.
• This results in an acidic solution.

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Basic Solutions

• A basic solution is the opposite, if a base is
  added to water more OH- ions are produced and the
  concentration of OH- ions becomes greater than
  that of H3O ions.

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Acid, basic and neutral solutions

• Therefore
• Acidic solutions contain a greater concentration
  of H3O ions than OH- ions.
• A neutral solution contains equal concentrations
  of H3O and OH-.
• Basic solutions contain a lower concentration of
  H3O ions than OH- ions.

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Measuring Acidity

• Experimental measurements show that all aqueous
  solutions contain both H3O ions and OH- ions and
  that the product of their molar concentrations is
  always 10-14 at 25C.
• This relationship is called the ionic product and
  can be represented by
• H3O x OH- 10-14 M2 at 25C
• The square brackets mean the concentration of the
  ions.

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H3O x OH- 10-14 M2 at 25C

• Pure water is neutral, so H3OOH-
• Since 10-7 x 10-7 10-14 M2
• H3O 10-7 M and OH- 10-7 M at 25C
• What happens to the OH- as we increase H3O?

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At 25C

• A Solutions is
• Acidic if H3O gt 10-7 M and OH- lt 10-7 M
• Neutral is H3O 10-7 M OH-
• Basic if H3O lt 10-7 M and OH- gt 10-7 M

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Worked Example

• 0.1 mol of hydrogen chloride (HCl) gas was
  bubbled into sufficient water to produce 1L of
  solution. Calculate the solution concentration
  of
• H3O ions
• OH- ions

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Worked Example

• In a 5.6x10-6 M HNO3 solution at 25C, calculate
  the concentration of
• H3O ions
• OH- ions

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The pH Scale

• The pH scale is a useful way of indicating the
  acidity of a solution.
• pH is defined as
• pH -log10H3O
• Where H3O is measured in mol L-1.
• The pH of a solution decreases as the
  concentration of hydrogen ions increases

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pH

• Since pH is a logarithmic scale, increasing the
  concentration of H by a factor of 10 results in
  a decrease of one pH unit.
• For example H 0.001 M at 25C
• Then the pH -log H
• -log 0.001
• -log 10-3
• -(-3)
• 3

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pH

• If H 0.01 M at 25C
• What would the pH be?
• If H 10-7 M at 25C
• What would the pH be?

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pH
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Calculating the pH of aqueous solutions

• In the following examples H is used represent
  H3O, since the terms can be used
  interchangeably.
• In order to calculate the pH of an aqueous
  solution, you must first calculate the
  concentration of the H ions and apply the
  formula
• pH -log10H3O

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Calculating the pH of aqueous solutions

• If the OH- ion concentration is given then the
  equation H3O x OH- 10-14 M2
• Must be used first to determine the hydrogen ion
  concentration in the solution at 25C

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Worked Example

• What is the pH of a solution in which
• H 0.0135M?

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Worked examples

• What is the pH of a 0.0050 M solution of Ba(OH)2?
• What is the pH of a solution, at 25C, that
  contains 1.0g NaOH in 100mL solution?
• 30.0mL of 0.100M HNO3 is added to 50.0mL water.
  What is the pH of the diluted solution.

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Calculating the concentration of H in a solution
of a given pH

• If a pH of a solution is known, it can be used to
  determine the concentration of hydronium ions.
• The pH relationship can be used in the form
• H 10-pH
• If pH 5.00, H 10-5.
• 0.0000100 M

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Worked Example

• What is H in a solution of pH 3.47?
• What is the concentration of OH- ions in a
  solution of pH 10.4?

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Strong acid/Strong base
The steepest slant of the curve demonstrates the
equivalence point on the titration curve
Small volume of strong acid added produces a
large changes in pH. This is demonstrating a
sharp end point.
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pH Curves
pH curves showing change of pH during a titration
of a a strong base with a strong acid, and b a
weak base with a strong acid. Phenolphthalein,
which changes colour in the pH range 8.210,
givesa sharp end point in a but a broad end
point in b. Methyl orange, which changes colour
between pH 3.1 and 4.5, would be a more suitable
indicator for the second titration.
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Which indicator would be best to identify the
equivalence point?
Chapter 4 Q6. The graphs in Figure 4.7 show the
pH curves for titrations involving combinations
of acids and bases of various strengths. You have
a choice of phenolphthalein and methyl orange
indicator. Phenolphthalein changes colour over a
pH range 8.2 to 10.0. Methyl orange changes
colour between pH 3.2 and 4.4. Decide which
indicators would be suitable to identify the
equivalence point for each reaction. Provide
reasons for your selections.
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Which indicator would be best to identify the
equivalence point?
Change in pH during a titrations of a a
strong acid with a strong base b a strong acid
with a weak base c weak acid with a strong base
d weak acid with a weak base.
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Answer
Chapter 4 A6. a The equivalence point occurs in
the range pH 3 to pH 11. Both indicators will
change colour over this pH range. Both indicators
will provide a sharp end point, i.e. they will
change colour at the equivalence point with the
addition a small volume, 1 drop, of acid. b The
equivalence point occurs in the pH range 3 to 7.
Methyl orange provides a sharper end point over
this pH range. c The equivalence point occurs in
the pH range 7 to 11. Phenolphthalein provides
the sharper end point. d Both indicators will
provide a broad end point and neither would be
suitable.