Acids, bases & buffers

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Created by
Divya Lambotharan

Cards (50)

    • A Bronsted Lowry acid is a proton donor
    • a Bronsted Lowry base is a proton acceptor
  • Bronsted Lowry acids & bases:
    • acids dissociate and release H+ ions in aqueous solutions
    • alkalis dissociate and release OH- ions in aqueous solutions
    • H+ ions are neutralised by OH- ions to form water
    • H+ (aq) + OH- (aq)—> H2O (l)
    • An alkali is a soluble base
    • HCl —> H+ (aq) + Cl (aq)
    • HCl (aq) and Cl- (aq) are called a conjugate acid base pair
    • In the forward direction, HCl releases a proton to form its conjugate base Cl-
    • In the reverse direction Cl- accepts a proton to form its conjugate acid HCl
    • HCl + OH- -> H2O + Cl-
    • HCl is an acid as it donates H+
    • OH- is a base as it accepts H+
  • Monobasic acids —>An acid that contains only one hydrogen atom that can be replaced in an acid base reaction
  • Dibasic acid —> An acid that has two acidic hydrogen atoms in its molecule
  • Tribasic acid —> A substance that produces three hydrogen ions per molecule of the acid, when it gets completely dissociated in water
    • acid + metal —> salt + hydrogen
    • acid + carbonate —> salt + water + carbon dioxide
    • acid + base —> salt + water
    • acid + alkali –> salt + water
  • the pH scale:
    • pH less than 7 shows increasing acidity
    • pH greater than 7 shows increasing alkalinity
    • pH 7 is neutral
  • pH - a logarithmic scale:
    • A low value of [H+ (aq) ] matches a high value of pH
    • A high value of [H+(aq)] matches a low value of pH
    • pH = -log[H+ (aq)]
    • [H+ (aq)] = 10^-pH
  • Calculating the pH of strong acids:
    • In aqueous solutions, a strong monobasic acid, HA completely dissociates
    • HA (aq) —> H+ (aq) + A- (aq)
    • For a strong acid, [H+ (aq)] is equal to the concentration of the acid, [HA (aq)]
    • The pH of a strong acid can be calculated directly from the concentration of the acid
  • The ionic product of water:
    • Kw = [H+][OH-]
    • As temperature increases, the equation moves right to oppose the increase in temperature therefore [H+] and [OH-] increases
    • Kw increases and therefore pH decreases
    • However, the water is still neutral as [H+] = [OH-]
    • Kw = [H+]^2
  • Neutral --> [H+] = [OH-]
  • Dilution of strong base:
    • concentration x old volume / new volume
    • [H+] = Kw / [OH-]
    • pH = -log( Kw/ [OH-] )
  • Calculating pH - weak acids:
    • Can't be calculated by just knowing the concentration
    • Need to know the original concentration & the extent of the ionisation
    • The extent of the ionisation is given by Ka. The dissociation constant for a weak acid
    • CH3CO2H --> CH3CO2- + H+
    • HA --> A- + H+
  • Ka = [H+][A-] / [HA]
    • each acid has its own Ka
    • The bigger the Ka (more dissociated) - the less weak the acid
  • pKa:
    • pKa = -log(Ka)
    • Ka = 10^-pKa
    • For pKa, the smaller the number the stronger the acid
  • Assumption 1: [HA] at equilibrium is presumed to be the same as the original concentration. This is more of a problem for weak acids because [HA] will differ more
    Assumption 2: [A-] = [H+]
    • A strong acid HA completely dissociates : [H+] = [HA]
    • A weak acid HA partially dissociates: [H+] doesn't equal [HA]
  • [H+] depends on:
    • the concentration of the acid
    • the acid dissociation constant Ka
  • Determination of Ka:
    Experimentally the Ka for a weak acid can be determined by
    • preparing a standard solution of weak acid of known concentration
    • measuring the pH of the standard solution using a pH meter
  • Approximation 1: The dissociation of water is negligible
    Approximation 2: The concentration of acids is much greater than the H+ concentration at equilibrium
    • Water ionises very slightly acting as both an acid and as a base
    • H2O --> H+(aq) + OH-(aq)
    • Ka = [H+(aq)][OH- (aq)] / [H2O(l)]
    • Kw is called the ionic product of water - the ions in water ( H+ and OH-) multiplied together
    • Kw = [H+ (aq)][OH-(aq)]
    • Kw varies with temperature
    • Kw = 1.00 x 10^-14 mol^2dm^-6
    • On dissociation, water is neutral - it produces the same number of H+ and OH- ions
    • [H+(aq)] = [OH-(aq)]
    • Kw is essentially an equilibrium constant that controls the concentration of H+ and OH- in aqueous solutions
    • A solution is acidic when [H+(aq)] > [OH-(aq)]
    • A solution is neutral when [H+(aq)] = [OH-(aq)]
    • A solution is alkaline when [H+(aq)] < [OH-(aq)]
    • A strong base is an alkali that completely dissociates in solution
    • NaOH (aq) --> Na+ (aq) + OH-(aq)
    • NaOH is a monobasic base as each mole of NaOH releases one mole of OH- ions
  • The pH of a strong base can be calculated from:
    • the concentration of the base
    • the ionic product of water Kw
  • Buffer solution --> solutions which resist changes in pH when small quantities of acid or alkali are added.
  • Uses of buffer solutions:
    • Standardising pH meters
    • Buffering biological systems
    • Maintaining the pH of shampoos
    • The weak acid, HA, removes added alkali
    • The conjugate base, A-, removes added acid
  • Making a buffer:
    • Add salt directly
    • Add approximately half of the mole of strong acid - leaving the acid in excess
  • Action of the buffer solution:
    • The conjugate acid-base pair, HA / A- in an acid buffer solution controls the pH
    • HA --> H+ + A-
    • The control of pH can be explained in terms of shifts in the equilibrium position using Le Chateliers principle
  • Conjugate base removes added acids:
    On addition of an acid H+
    • [H+] increases
    • H+ ions react with the conjugate base A-
    • The equilibrium position shifts to the left, removing most of the H+ ions
  • Weak acid removes added alkali:
    on addition of an alkali, OH-
    • [OH-] increases
    • The small concentration of H+ ions reacts with the OH- ions
    • HA dissociates, shifting the equilibrium position to the right to restore most of H+ ions
  • A buffer is most effective at removing either added acid or alkali when there are equal concentrations of the weak acid and its conjugate base
  • When [HA] = [A-]:
    • the pH of the buffer solution is the same as the pKa value of HA
    • the operating pH is typically over about 2 pH units, centered at the pH of the pKa value
    • the ratio of the concentrations of the weak acid and its conjugate base can then be adjusted to fine-tune the pH of the buffer solution
  • Buffers in blood:
    • Human blood contains a buffer of carbonic acid (H2CO3) & bicarbonate anion (HCO3-) in order to maintain blood pH between 7.35 & 7.45
    • CO2 is an important component of the blood buffer in regulation in the body
    • The effect of this can be important when the human body is subjected to strenuous conditions
    • H3O+ + HCO3- --> H2CO3 + H2O