acids and bases

  • Created by: Nataliagx
  • Created on: 29-11-18 09:59
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  • Acids & Bases
    • Arrhenius Theory
      • Acids: substances that produce hydrogen ions in solution.
      • Bases: substances that produce hydroxide ions in solution.
      • Neutralisation is from H+ ions and OH- ions that react to produce H2O
        • However, when HCl reacts with NH3, no OH-ons are needed
          • NH3 + H2O is a reversible reaction- that is how OH-ions are generated
    • Bronsted-Lowry Theory
      • Acid: Proton (H+) donor
      • Base: Proton (H+) acceptor
      • ammonia = base; accepts proton (H+) that is attached to a lone pair on the nitrogen of NH3
    • Strong acids:
      • HCl
        • H2O + HCl > H3O+ + Cl-
          • FORWARD: HA (HCl) is an acid. H2O is a base
          • REVERSE: H3O+ is an acid and A- (Cl-) is a base
          • The reversible reaction contains 2 acids and 2 bases: CONJUGATE PAIRS
            • conjugate pairs differ  by the presence/ absence of the H+.
          • H2O acts as acid & base: AMPHOTERIC
            • H3O+ is a very strong acid,
              • Kw= [H3O+] x [OH-]
                • OH- is a very strong base
                • little of the water is ionised and so, the conc remains unchanged
                • varies with temperature
                  • 1.00 x 10-14 mol2 dm-6 at room temperature
                    • INCREASE temp = FORWARD reaction  favoured,  more H+  and OH-  formed. This increases f Kw with temp.
                  • pH of pure water FALLS as the temp INCREASES
                    • although the pH changes, it still remains NEUTRAL
                • pKw = -log(Kw)
                  • pKw at room temp = 14
                • 1. find [H+] or [OH-] (they're the same thing)
                  • at room temp, Kw is:[H+] [OH-] = 1.00 x 10-14
                    • but as [OH-] = [H+] it can be written as [H+]2 = 1.00 x 10-14
                      • then square root: [H+] = 1.00 x 10-7 mol dm-3
                        • pH = -log [H+]
            • OH- is a very strong base
      • HNO3
      • HBr
      • H2SO4
    • Weak aids
      • HSO4-
      • HNO2
      • C6H5COOH
      • H3PO4
    • pH
      • concentration of hydrogen ions in a solution.
        • pH = -log [H+]
          • STRONG ACIDS: fully dissociate
            • if the conc of acid is 0.1 mol dm-3, then the conc of H+ is also 0.1 mol dm-3.
          • WEAK ACIDS: partially dissociates
            • The further to the left Eq is, the weaker the acid is.
            • Ka = [H+] [A-] / [HA]
              • Weak acid= small Ka
              • pKa= -log(Ka)
                • low pKa = Strong acid
                • Ka= 10 (pka)
                  • Ka = [H+] [A-] / [HA]
                    • Weak acid= small Ka
                    • pKa= -log(Ka)
                      • low pKa = Strong acid
                      • Ka= 10 (pka)
            • pH curve & Titration
              • indicator changes colour= END POINT
                • Weak acid: LITMUS PAPER
                  • OH- ions
                  • H+ ions
                  • OH- = H+
              • equation proportions = EQUIVALENCE POINTS
      • Buffer Solutions
        • resists changes in pH when small amounts of  acid or alkali is added
          • Acid Buffer= pH lower than 7; weak acid + salt
            • Ethanoic acid (weak acid); equilibrium shift left.
              • ethanoate ions from the sodium ethanoate
                • un-ionised ethanoic acid
                  • H+ combine w/ the ethanoate ions = ethanoic acid.
                    • ethanoate ions from the sodium ethanoate
                      • un-ionised ethanoic acid
                        • H+ combine w/ the ethanoate ions = ethanoic acid.
                • enough H+  to make the solution acidic
                  • Alkali removal w/ ethanoic
                    • Acid reacts w/ OH- to make ethanoate ions + H2O
                  • Alkali removal w/ H+
                    • OH- combine with H+ from ethanoic acid = H2O
                      • Eq shifts RIGHT to replace lost H+ ions.
                • 0.10 mol dm-3 of ethanoic acid and 0.20 mol dm-3 of sodium ethanoate.  pH?
                  • Ka = [H+] [A-] / [HA]
                    • assume that the ethanoate ion conc = conc of sodium ethanoate; 0.20 mol dm-3.
                    • [H+] = Ka x [acid]/[salt]
                  • Ka for ethanoic acid is 1.74 x 10-5 mol dm-3
                    • to calculate PROPORTIONS- reverse it
                      • 1. In-log the pH (10) = [H+]
                        • 2. Ka / [H+]
                          • conc of ethanoate ions (from the sodium ethanoate) in the solution has to be 0.5 times that of the conc of the acid.
              • Alkali Buffer= pH higher than 7; weak base + salt
                • ammonia +  ammonium chloride: weak base = Eq shift LEFT
                  • unreacted ammonia
                  • ammonium ions from the ammonium chloride
                  • enough OH- to make the solution alkaline
                  • Acidic Removal w/ ammonia
                    • NH3 + H+ = NH4 +
                  • Acidic removal w/ OH-
                    • H+  combine w/ OH-  to make H2O
                      • Eq shifts RIGHT to restore OH-
                  • pH w/ 0.100 mol dm-3 of ammonia and 0.0500 mol dm-3 of ammonium chloride
                    • Ka  ammonium ion is 5.62 x 10-10 mol dm-3
                      • [H+]= [alkali]/[salt]

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