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CHEMISTRY U1/2 > CHAPTER 11: ACIDS AND BASES > Flashcards

CHAPTER 11: ACIDS AND BASES Flashcards

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1
Q

bronsted lowry theory of acids and bases

A
  • acids are proton donors → donate H+ ions
  • bases are proton acceptors → accept H+ ions
  • an acid-base reaction involves an exchange of protons from an acid to a base
    • eg. HCl (g) + H2O (l) → H3O+ (aq) +Cl- (aq)
    • H3O+ - hydronium ion
    • eg. HCl (aq) + NH3 (aq) → NH4+ (aq) + Cl- (aq) → NH4Cl (s)
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2
Q

conjugate acid base pairs

A

a conjugate acid-base pair is two species that differ by a proton ie H+
eg.
- HCl (g) + H2O (l) → H3O+(aq) + Cl-(aq)
- conjugate acid-base pair: HCl/Cl- and H3O+/H2O
- NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)
- conjugate acid-base pair: NH4+/NH3 and H2O/OH-

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3
Q

amphiprotic substances

A

amphiprotic substances are substances that can donate or accept protons depending on what they are reacting with (act as either an acid or a base)

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4
Q

monoprotic acids

A
  • monoprotic acids can donate only one proton
  • in CH3COOH2 only the hydrogen that is part of the highly polar O-H bond is donated
  • this hydrogen atom is called the acidic proton
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5
Q

polyprotic acids

A

polyprotic acids are acids that can donate ore than one proton

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6
Q

diprotic acids

A
  • diprotic acids can donate two protons
  • stage 1
    • H2SO4(l) + H2O (l) → HSO4-(aq) + H30+(aq)
      • strong acid
      • occurs almost to completion
  • stage 2
    • HSO4-(aq) + H2O(l) ⇆ SO4 2- (aq) + H3O+(aq)
      • weak acid
      • only partially ionised
      • a double reversible arrow indicates that an incomplete reaction occurs
        (use the product of first stage in second stage)
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7
Q

triprotic acids

A
  • can donate 3 protons
  • stage 1
    • H3PO4(aq) + H2O(l) ⇆ H2PO4-(aq) + H3O+(aq)
      • Ka= 6.8 x 10^-3
    • acid dissociation → Ka
  • stage 2
    • H2PO4-(aq) + H2O(l) ⇆ HPO4 2-(aq) + H3O+(aq)
    • Ka2 = 6.2 x 10^-8
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8
Q

strength of acid and bases

A
  • the strength of an acid is its ability to donate hydrogen ions to a base
  • the strength of a base is a measure of its ability to accept hydrogen ions from an acid
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9
Q

strong acids

A
  • hydrochloric acid HCl
  • sulfuric acid H2SO4
  • nitric acid HNO3
  • complete ionisation
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10
Q

weak acids

A
  • ethanoic acid CH3COOH
  • carbonic acid H2CO3
  • phosphoric acid H3PO4
  • incomplete ionisation
  • double arrows
    • some reactants and some products
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11
Q

strong bases

A
  • sodium oxide NaO
  • sodium hydroxide NaOH
  • potassium hydroxide KOH
  • calcium hydroxide Ca(OH)2
  • complete reaction
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12
Q

weak bases

A
  • ammonia NH3
  • incomplete reaction
  • double arrows
    • some reactant and some products are present
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13
Q

relative strength of conjugate acid-base pair

A
  • the stronger an acid is, the weaker its conjugate base
  • the stronger a base is, the weaker its conjugate acid
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14
Q

strength vs concentration

A
  • strength: full completion/ionisation of a reaction
  • concentration: the amount of a given substance (the number of molecules) in a given volume of water
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15
Q

neutralisation reaction

A
  • acid + metal hydroxide → salt + water
  • eg. H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2H2O (l)
  • ionic equation: hydrochloric acid and sodium hydroxide
    • H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) → Na+ (aq) + Cl- (aq) + H2O (l)
    • spectator ions: Na+ (aq) and Cl- (aq)
    • ionic equation: H+ (aq) + OH- (aq) → H2O (l)
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16
Q

acid and metal carbonate

A
  • acid + metal carbonate → salt + water + carbon dioxide
    • eg. 2HCl (aq) + Na2CO3 (aq) → 2NaCl (aq) + CO2 (g) + H2O (l)
  • acid + metal hydrogen carbonate → salt + water + carbon dioxide
    • eg. HCl (aq) + NaHCO3 (s) → NaCl (aq) + H2O (l) + CO2 (g)
17
Q

acid and reactive metals

A
  • acid + reactive metal → salt + hydrogen
  • eg.
    • full chemical equation: 2HCl (aq) + Zn(s) → ZnCl2 (aq) + H2 (g)
    • full ionic equation: 2H+ (aq) + 2Cl- (aq) + Zn (s) → Zn 2+ (aq) + 2HCl- + H2 (g)
    • ionic equation: 2H+ (aq) + Zn (s) → Zn2+ (aq) + H2 (g)
18
Q

antacid

A
  • weak bases used to treat the symptoms of excess HCl in the stomach
  • they neutralise the acid
19
Q

ionic product of water

A

H2O (l) + H2O (l) ⇆ H3O+ (aq) + OH- (aq)

  • in pure water at 25 degrees Celsius, the H3O+ and OH- concentrations are each 10^-7 M
    • M = mol/L
20
Q

ionisation constant of water

A
  • Kw = [H3O+][OH-]= 1.00 X 10^-14m^2 at 25 degrees celsius
  • if the [H3O+] increases, the [OH-] decrease
  • if the [OH-] increases, the [H3O+] decreases
  • according to LeCharelier’s principle, an increase in either the [H3O+] or [OH-] will drive the equilibrium between water and its ions to the right, reducing the number of [H3O+] and [OH-] ions
21
Q

pH formula

A

pH = -log10[H3O+]
[H3O+] = 10^-pH

22
Q

what are indicators

A
  • is a weak acid with a colour that is different to the colour of its weak conjugate base
  • natural Indicators
    • Litmus
    • In the presence of acids, litmus turns red and it turns blue in basic solutions.
23
Q

types of indicators

A
  • red cabbage indicator
  • bromothymol blue
  • methyl orange
  • phenolphthalein
  • universal indicator (mixture of diff indicators)
24
Q

methyl orange

A
  • the indicator changes colour between pH 3.1 and pH 4.4.
  • between these pH values the indicator has an orange colour
25
Q

phenolphthalein

A

it changes colour over the pH range 8.3 to 10.0.

26
Q

accuracy vs precision

A
  • accurate measurements are close to the true value
  • precise measurements are close to each other and are reproducible.
27
Q

red cabbage juice

accuracy and precision

A
  • accuracy: sample value can be determined to within 2 pH units of the true value
  • precision: generally reproducible to within 2 pH units
28
Q

universal indicator

accuracy and precsion

A
  • accuracy: sample value can be detemined to within 1 pH of the true value
  • precision: generally reproducible to within 1 pH unit
29
Q

pH meter

accuracy and precision

A
  • accuracy: sample value can be determined to within 0.01 pH units of the true value
  • precision: generally reproducible to within 0.01 pH of the true value
30
Q

metal ocides vs non-metal oxides

A
  • metal oxides produce basic solutions
  • non-metal oxides produce acidic solutions

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