Acid and Base Equilibria

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Acid and Base Equilibria

• Chapter 8

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Acid and base definitions based on experiments.

• Acids
• Tastes sour
• Conducts electricity
• Changes the colour of litmus paper from blue to
  red
• Turns neutral (green) Bromthymol blue to yellow
• Reacts with active metals such as zinc and
  magnesium, liberating hydrogen gas (H2).
• Reacts with carbonates, releasing carbon dioxide
  gas (CO2)

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• Bases
• Tastes bitter
• Conducts electricity
• Changes the colour of litmus paper from red to
  blue
• Turns neutral (green) Bromthymol blue to blue
• Turns colourless phenolphthalein to pink
• Reacts with an acid to destroy its properties
• Feels slippery

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Arrhenius Theory Review

• Acids are solutes that produce hydrogen
  ions/protons in aqueous solutions (increases the
  concentration of H) or hydronium ions (H3O).
  Ex
• HCl H(aq) Cl-(aq) 
• OR H2O(l) HCl(g) H3O (aq) Cl-(aq)
• Bases produce hydroxide ions when dissolved in
  water.
• NaOH(s) ? Na(aq)  OH-(aq)
• However, this model does not account for basic
  properties of compounds that do not contain
  hydroxide ions, such as ammonia NH3(aq)
• NH3(g) H2O(l) NH4 (aq) OH-(aq)

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Bronsted-Lowry Theory

• According to this theory, an acid is a proton
  donor, and a base is a proton acceptor. A
  substance can only be classified as one or the
  other for a particular reaction (as it can change
  from one reaction to another). ? focus on proton
  transfer
• Ex. of acid HCl when hydrogen chloride reacts
  with water, a proton is transferred from HCl to
  H2O
• H2O HCl(g) H3O (aq) Cl-(aq)
• base acid conj. acid conj. base.
• Ex. of base when ammonia reacts with water,
  water now acts as an acid because it donates a
  proton to ammonia, which is the Bronsted-Lowry
  base
• NH3(g) H2O(l) NH4 (aq) OH-(aq)
• base acid conj. acid conj. base.

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Amphoteric

• As you can see in the above reactions, water can
  act as either a base or as an acid. A substance
  that can act as either a Bronsted-Lowry acid or
  B-L base given the reaction is called amphoteric
  (amphiprotic) it can donate or accept a proton.
• amphoteric may act as an acid or base
  amphiprotic may accept or donate protons for
  Bronsted-Lowry acids and bases, amphiprotic is
  always amphoteric, but not in more general
  definitions equilibrium conditrion like water
  or bicarbonate ion in baking soda

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Amphoteric Water

• Another example of water as amphoteric
• HCO3-(aq) H2O(l) H2CO3(aq) OH-(l)
• base acid
• HCO3-(aq) H2O(l) CO32-(aq) H3O(aq)
• acid base
• (autoionzation)

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Neutralization

• benefit of B-L rather than Arrhenius is that it
  defines acids and bases in terms of chemical
  reactions, so that neutralizations do not have to
  produce water and salt, as according to Arrhenius
• Arrhenius neutralization acid-base
  neutralization produces water and salt as in
• NaOH(aq) HCl(aq) ? H2O(l) NaCl(aq)
• NH4OH(aq) HCl(aq) ? H20(l) NH4Cl
• Acid-base neutralization without hydronium ions,
  hydroxide ions, or water
• NH3(g) HCl(g) ? NH4Cl
• (proton transferred from Cl atom to N atom)

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Reversible Acid-Base Reactions
PromptWhich is the acid? How do we know
that? Where is the proton transfer?

• (Equilibrium implied in Bronsted-Lowry reaction)
• In each proton transfer reaction at equilibrium,
  both forward and reverse reactions involve
  Bronsted-Lowry acids and bases. On each side of
  the reaction are acids and bases, which are
  called conjugate acid-base pairs.
• For ex.
• HC2H3O2(aq) H2O(l) ? C2H3O2-(aq) H3O(aq)
• acid base conj base conj acid
• In any acid-base equilibrium, there will always
  be two acids and two bases.
• The base on the right (product) is formed by the
  removal of the proton from the acid on the left.
• The acid on the right (product) is formed by the
  addition of a proton to the base on the left.
• A conjugate acid-base pair is a pair of
  substances whose molecular formula differ by a
  single H ion (proton). ( ?the acid has one more
  proton than the base)
• The acid in the forward reaction is a proton
  donor and the base is the proton acceptor.
• In the reverse reaction, the conjugate acid is
  the proton donor and the conj. base is the
  acceptor.

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Competition for protons (p.530)

• View acid-base reactions as a competition for
  protons between two bases.
• Strong acids in reactions go to almost 100
  ionization or percent reaction. (the proton
  transfer is almost a complete forward reaction
  and almost no reverse proton transfer occurs)
  therefore, an equilibrium is not established in
  this case
• Weak acids have a lower ionization, so their
  equilibrium position favours the reactants rather
  than the products.
• The stronger an acid, the weaker its conjugate
  base the acid has a weaker affinity for the
  proton, and easily loses/transfers the proton to
  its conjugate base.
• The weaker an acid, the stronger its conjugate
  base. The stronger the base, the stronger the
  attraction for protons.

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Strong and weak acids

• Bronsted-Lowry explanation of strong and weak
  acids
• HA(aq) is used as general symbol for acid, and
  A-(aq) as its conjugate base.
• Ionization reaction
• HA(aq) H2O(l) A-(aq) H30(aq)
• The strength of the acid HA is determined by the
  extent of the proton transfer.
• The ionization equation of an acid in water is
  often abbreviated
• HA(aq) H2O(l) ? A-(aq) H30(aq)
• to
• HA(aq) ? A-(aq) H(aq)

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continued

• For ex, if we go back to our previous example,
• HC2H3O2(aq) H2O(l) C2H302-(aq)
  H30(aq)
• acid base conj base conj acid
• then it reduces to
• HC2H3O2(aq) ? C2H302-(aq) H(aq)
• Although this reduced equation shows the change
  that takes place, it does not demonstrate the
  important role that water plays in causing the
  acid to ionize, or that the proton most likely
  exists as a hydronium ion (H3O).

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Strong Acids (p.534 )

• an acid that is assumed to ionize quantitatively
  (completely) in aqueous solution (100
  ionization is gt .99 - however, we assume its
  100 in calculations)
• hydrochloric acid HCl(aq), hydrobromic acid
  HBr(aq), sulfuric acid H2SO4(aq), nitric acid
  HNO(aq), phosphoric acid H3PO4(aq)
• Monoprotic acid an acid that possesses only one
  ionizable (acidic) proton ex. HCl
• Diprotic acid an acid that possesses two
  ionizable (acidic) protons ex. H2SO4
• Triprotic acid possesses three ionizable
  hydrogen atoms ex. H3PO4
• Note Dissociation Chemistry
• a. The process by which the action of a solvent
  or a change in physical condition, as in pressure
  or temperature, causes a molecule to split into
  simpler groups of atoms, single atoms, or ions.
• b. The separation of an electrolyte into ions of
  opposite charge.

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Strong Bases (p.537)

• an ionic substance that (according to Arrhenius)
  dissociates completely in water to release
  hydroxide ions
• dissolve in water (dissociate completely)
• LiOH(s), NaOH(s), KOH(s), RbOH(s), CsOH(s)
• Mg(OH)2(s), Ca(OH)2(s), Ba(OH)2(s), Sr(OH)2(s)
• For every mole of metal hydroxide that is
  dissolved, one mole of hydroxide ion is produced.
• NaOH(s) ? Na(aq) OH-(aq)
• Group 2 elements dissolve in water and form 2
  hydroxide ions
• Ba(OH)2(s) ? Ba2(aq) 2OH-(aq)
• Hydroxide ions react with hydrogen ions shifts
  equilibrium to right, causing undissolved salts
  to dissolve and produce higher hydroxide ion
  concentrations

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Hydrogen Ion Concentration and pH (p. 540)

• There is a wide range of concentration of
  hydrogen ions and hydroxide ions in acidic and
  basic solutions. The pH scale was developed to
  measure hydrogen ion concentration.
• pH logH(aq)
• Example Calculate the pH of a solution with a
  hydrogen ion concentration of 5.3 x 10-9.
• pH log H(aq)
• pH log (5.3 x 10-9) (two sig digits)
• pH 8.28
• The solution has a pH of 8.28.

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• (p. 541)
• pH of pure (neutral) water and any neutral
  solution at SATP is 7.00 as the hydrogen and
  hydroxide ions are equal.