CHAPTER 11: ACIDS AND BASES Flashcards
bronsted lowry theory of acids and bases
- 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)
conjugate acid base pairs
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-
amphiprotic substances
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)
monoprotic acids
- 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
polyprotic acids
polyprotic acids are acids that can donate ore than one proton
diprotic acids
- diprotic acids can donate two protons
- stage 1
- H2SO4(l) + H2O (l) → HSO4-(aq) + H30+(aq)
- strong acid
- occurs almost to completion
- H2SO4(l) + H2O (l) → HSO4-(aq) + H30+(aq)
- 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)
- HSO4-(aq) + H2O(l) ⇆ SO4 2- (aq) + H3O+(aq)
triprotic acids
- can donate 3 protons
- stage 1
- H3PO4(aq) + H2O(l) ⇆ H2PO4-(aq) + H3O+(aq)
- Ka= 6.8 x 10^-3
- acid dissociation → Ka
- H3PO4(aq) + H2O(l) ⇆ H2PO4-(aq) + H3O+(aq)
- stage 2
- H2PO4-(aq) + H2O(l) ⇆ HPO4 2-(aq) + H3O+(aq)
- Ka2 = 6.2 x 10^-8
strength of acid and bases
- 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
strong acids
- hydrochloric acid HCl
- sulfuric acid H2SO4
- nitric acid HNO3
- complete ionisation
weak acids
- ethanoic acid CH3COOH
- carbonic acid H2CO3
- phosphoric acid H3PO4
- incomplete ionisation
- double arrows
- some reactants and some products
strong bases
- sodium oxide NaO
- sodium hydroxide NaOH
- potassium hydroxide KOH
- calcium hydroxide Ca(OH)2
- complete reaction
weak bases
- ammonia NH3
- incomplete reaction
- double arrows
- some reactant and some products are present
relative strength of conjugate acid-base pair
- the stronger an acid is, the weaker its conjugate base
- the stronger a base is, the weaker its conjugate acid
strength vs concentration
- strength: full completion/ionisation of a reaction
- concentration: the amount of a given substance (the number of molecules) in a given volume of water
neutralisation reaction
- 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)
acid and metal carbonate
-
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)
acid and reactive metals
- 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)
antacid
- weak bases used to treat the symptoms of excess HCl in the stomach
- they neutralise the acid
ionic product of water
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
ionisation constant of water
- 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
pH formula
pH = -log10[H3O+]
[H3O+] = 10^-pH
what are indicators
- 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.
types of indicators
- red cabbage indicator
- bromothymol blue
- methyl orange
- phenolphthalein
- universal indicator (mixture of diff indicators)
methyl orange
- the indicator changes colour between pH 3.1 and pH 4.4.
- between these pH values the indicator has an orange colour
phenolphthalein
it changes colour over the pH range 8.3 to 10.0.
accuracy vs precision
- accurate measurements are close to the true value
- precise measurements are close to each other and are reproducible.
red cabbage juice
accuracy and precision
- accuracy: sample value can be determined to within 2 pH units of the true value
- precision: generally reproducible to within 2 pH units
universal indicator
accuracy and precsion
- accuracy: sample value can be detemined to within 1 pH of the true value
- precision: generally reproducible to within 1 pH unit
pH meter
accuracy and precision
- 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
metal ocides vs non-metal oxides
- metal oxides produce basic solutions
- non-metal oxides produce acidic solutions
