Understanding Points |
Reactivity 3.1.1—Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton
acceptor.
Reactivity 3.1.2—A pair of species differing by a single proton is called a conjugate acid–base pair.
Reactivity 3.1.3—Some species can act as both Brønsted–Lowry acids and bases.
Reactivity 3.1.4—The pH scale can be used to describe the [H+] of a solution:
pH = –log10[H+]; [H+] = 10–pH
Reactivity 3.1.5—The ion product constant of water, Kw, shows an inverse relationship between [H+] and [OH–]. Kw = [H+] [OH–]
Reactivity 3.1.6—Strong and weak acids and bases differ in the extent of ionization.
Reactivity 3.1.7—Acids react with bases in neutralization reactions.
Reactivity 3.1.8—pH curves for neutralization reactions involving strong acids and bases have
characteristic shapes and features. |
Brønsted–Lowry Theory
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Brønsted–Lowry acid: proton (H+) donor
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Brønsted–Lowry base: proton (H+) acceptor
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e.g. HCl + NH3 → NH4+ + Cl-
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HCl has donated a proton to NH3, so HCl is an acid
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NH3 has accepted a proton from HCl, so NH3 is a base
Common Acids and Bases
Common Acids | Common Bases |
HCl, H2SO4, HNO3
H2CO3, H3PO4, CH3COOH, C2H5COOH, HCOOH | NaOH, Mg(OH)2, CuO - metal (hydro)oxides
NH3, Organic amines- CH3NH2, C6H5NH2, etc. |
Conjugate pairs
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Differ by just one proton (H+)
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Conjugate base: an acid which has lost a hydrogen proton
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Conjugate acid: a base which has gained a hydrogen proton.
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HA (acid) + B (base) ⇌ A (conj. base) + BH+ (conj. acid)
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As the base B gained a H+ to form BH+, BH+ is the conjugate acid of B
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As the acid HA lost a H+ to form A-, A- is the conjugate acid of HA
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HA and A- → form a conjugate acid-base pair
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B and BH+ → conjugate base-acid pair
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TIP: to find the conjugate acid of a base, add an H+ to the base
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To find the conjugate base of an acid, remove an H+ and adjust the charge as necessary
Amphiprotic species
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Amphoteric species: can either donate or accept a proton
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Amphiprotic species: can act as both Brønsted-Lowry acids and bases
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All amphiprotic compounds are amphoteric but not all amphoteric compounds are amphiprotic, e.g. aluminium oxide (refer to “Period 3 Oxides”)
Water as an amphiprotic species:
Water | |
As a Base | As an Acid |
CH3COOH + H2O ⇌ CH3COO– + H3O+
H2O accepts a proton from CH3COOH | NH3 +H2O ⇌ NH4+ + OH–
H2O donates a proton to NH3 |
Other examples = conjugate bases of diprotic weak acids, e.g. H2CO3, H2SO4, H3PO4
•
H2CO2, H2SO4, H3PO4 can act as both acids and bases in HCO3- and HSO4-, H2PO4- states
(Base)
HSO4- + H2O ⇌ H2SO4 + OH-
Base Acid C. Acid C. Base | (Acid)
HSO4- + H2O ⇌ SO42- + H3O+
Acid Base C. Base C. Acid |
pH scale
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pH: ‘potential of hydrogen’
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Indicates the acidity or basicity of a given solution.
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pH = –log10 [H+ (aq)] → rearranging the equation gives [H+] = 10–pH
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e.g. A solution that has [H+] = 0.1 mol dm–3 ⇒ [H+] = 10–1 mol dm–3 ⇒ pH = 1
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e.g. A solution that has [H+] = 0.01 mol dm–3 ⇒ [H+] = 10–2 mol dm–3 ⇒ pH = 2
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General Rules:
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pH numbers are usually positive and have no units
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The pH number is inversely related to [H+]
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A change of one pH unit represents a 10-fold change in [H+]
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e.g. pH=3 → [H+]=10-3 and pH=4 → [H+]=10-4 differ by a factor of 10
Autoionization of water
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2H2O (l) ⇌ H3O+ (aq) + OH– (aq)
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(H2O (l) ⇌ H+ (aq) + OH– (aq))
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2 H2O molecules act as a base / acid to form hydronium (H3O+) and hydroxide (OH-)
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Kc = [H+][OH-][H2O]2 but since H2O is the solvent we can disregard the H2O concentration
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Kw = [H+][OH–] = Ionic product constant of water (ionisation constant)
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At 298K, Kw = 1.00 × 10–14
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From the reaction we know that [H3O+]=[OH-] → Kw=[H3O+]2 or Kw=[H+]2
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∴ At 298K, [H+] = √Kw = 1.00 × 10–7/pH = 7 for pure H2O
Solution | [H+] = Kw/[OH-] | pH + pOH = 14 |
Acidic | [H+] > [OH-] | pH < 7 |
Neutral | [H+] = [OH-] | pH = 7 |
Basic | [H+] < [OH-] | pH > 7 |
Strong and weak acids/bases
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Dissociation: the separation of ionic compounds into its respective cation and anion in solution
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Strong Acid/Base: fully dissociates into its ions in aq. solution (i.e. more ionization)→ more H+ or OH- ions
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Reactive H+ can easily be donated to another species
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Weak Acid/Base: partially dissociates into its ions in aq. Solution (i.e. less ionization) → less H+ or OH- ions, lower electrical conductivity
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Reactive OH- can easily abstract a proton from another species
Strong (indicated with single arrow) | Weak (indicated with equilibrium arrow) | |
Acid | HCl → H+ + Cl- | CH3COOH ⇄ CH3COO- + H+ |
Base | NaOH → Na+ + OH- | NH3 + H2O ⇄ NH4+ +OH- |
1.0 mol dm-3 solution of HCl | 1.0 mol dm-3 of CH3COOH |
Fully dissociated | Partially dissociated |
All of HCl is ionized to form H+ | Only some of CH3COOH is ionized to form H+ |
[H+]=[Cl-] = 1.0 mol dm-3 | [H+]=[CH3COO-] < 1.0mol dm-3 |
More ions to carry charge → higher conductivity | Less ions to carry charge → lower conductivity |
Distinguishing through gas formation:
Gas formed | Strong Acid | Weak Acid | |
Acid + metal | H2 gas | Vigorous bubbling | Light bubbling |
Acid + metal oxide | n/a | n/a | n/a |
Acid + metal hydroxide | n/a | n/a | n/a |
Acid + metal hydrogen carbonate | CO2 gas | Vigorous bubbling | Light bubbling |
Acid + metal carbonate | CO2 gas | Vigorous bubbling | Light bubbling |
Distinguishing strong/weak with other methods
Strong Acid/Base | Weak Acid/Base | |
Electrical conductivity*
• Dependent on mobile ions | Higer | Lower |
Rate**
• Dependent on [H+] or [OH-] | Faster | Slower |
pH
• Dependent on [H+] or [OH-] | Close to 1 | Close to 14 |
Reactions of acids with metal/base/carbonates
1.
Acid + reactive metal* → salt + hydrogen gas
HCl (aq) + Na (s) → NaCl (aq) + H2 (g)
2H3PO4 (aq) + 6Li (s) → 2Li3PO4 (aq) + 3H2 (g)
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Examples of reactive metals: group 1 and 2 metals
1.
Acid + Metal Oxide → salt + water
2HCl (aq) + MgO (s) → MgCl2 (aq)+ 2H2O (l)
H2SO4 (aq) + Li2O (aq) → Li2SO4 (aq) + H2O (l)
1.
Acid + Metal Hydroxides → salt + water**
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
H2CO3 (aq) + Mg(OH)2 (aq) → MgCO3 (aq) + 2H2O (l)
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*More generally, an acid and base can react in a neutralization reaction
Acid + Base → salt + water
HNO3 (aq) + NH4OH (aq) → NH4NO3 (aq) + H2O (l)
H2SO4 (aq) + NH4OH (aq) → (NH4)2SO4 (aq) + H2O (l)
1.
Acid + Metal Carbonate → Salt + CO2 gas+ water
2HCl (aq) + K2CO3 (aq) → 2KCl (aq) + CO2 (g)+ H2O (l)
2CH3COOH (aq) + CaCO3 (aq) → (CH3COO)2Ca (aq) + CO2 (g) + H2O (l)
1.
Acid + metal hydrogen carbonate → salt + water + carbon dioxide gas
2HCl (aq) + Na2CO3 (s) → 2NaCl (aq) + H2O (l) + CO2 (g)
HCl (aq) + NaHCO3 (s) → NaCl (aq) + H2O (l) + CO2 (g)
H2SO4 (aq) + 2NaHCO3 (s) → Na2SO4 (aq) + 2H2O (l) + 2CO2 (g)
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For carboxylic acid, organic weak acid, metal cation (M+) is written at rear end of chemical formula to emphasize where the electrostatic attraction between the negative and positive ions occurs
2CH3COOH (aq) + 2Na (s) → 2
CH
3
COONa
(aq) + H2 (g)
2CH3COOH (aq) + Mg (s) → (CH3COO)2Mg (aq) + H2 (g)
3CH3COOH (aq) + Al (s) → (CH3COO)3Al (aq) + 1½H2 (g)
CH3COOH (aq) + NaOH (aq) → CH3COONa (aq) + H2O (l)
Tips on balancing acid reactions:
1.
Write out non-salt products first (H2, H2O, CO2 + H2O)
2.
Deduce salt chemical formula with correct cation and anion charge balance (net charge=0)
3.
Balance molar ratio between products and reactants
a.
Balance the anion in the acid with the salt anion
b.
Balance the salt cation with the base cation
c.
Add coefficients as necessary to remaining byproducts (H2, H2O, CO2, H2O)
Neutralization reaction: a reaction between an acid and a base to form salt and water
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HA (aq) + M(OH) (aq) → MA (aq) + H2O (l)
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H+ + A- + M+ + OH → M+ + A- + H2O
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In solution, A- and M+ are spectator ions.
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The enthalpy of the neutralization reaction is determined by the reaction between H+ and OH-: H+ (aq) + OH- (aq) → H2O (l) ΔH<0 (exothermic, ∵ bond making)
Salt: product of acid-base neutralization rxn = anion of the acid and cation of the base (with balanced charge)
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Mg(NO3)2 ← HNO3 + Mg(OH)2
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NaCO3 ← H2CO3 + NaOH
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K2SO4 ← H2SO4 + KOH
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CH3COONa ← CH3COOH + NaOH
Non Hydroxide Bases
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Alkali: A soluble base (i.e. can dissociate to ions in water)
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Not all bases are hydroxides; Metal Oxides, Metal carbonates, Ammonia are also bases
Alkalis | Insoluble bases |
Metal Oxide:
K2O ⇌ 2K+ + O2-
O2- + H2O → 2OH-
Overall: K2O (s) + H2O (l) → 2K+ (aq) + 2OH– (aq)
Ammonia:
NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH– (aq)
Metal Carbonate:
Na2CO3 (aq) → Na+ (aq)+ CO32- (aq)
CO32–(aq) + H2O(l) ⇌ HCO3–(aq) + OH–(aq)
Overall: Na2CO3 (aq) + H2O(l) → Na+(aq) + HCO3– (aq) + OH– (aq)
Metal Hydrogen Carbonate:
KHCO3 → K+ + HCO3-
HCO3–(aq) ⇌ CO2(g) + OH–(aq)
Overall: KHCO3 → K+ + CO2 + OH- | Metal Oxides
• except Na2O, K2O, SrO, BaO, CaO, which are slightly soluble
Metal Hydroxides
• Except NaOH, KOH, Sr(OH)2, Ba(OH)2, Ca(OH)2, which are slightly soluble
• They become more soluble as you go down the column (lower lattice enthalpy) |
Acid-Base Titration: Experiment to measure unknown concentration of acid/base by neutralisation reaction
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Titrant : solution with known concentration, add from burette
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Analyte is the solution with unknown concentration, fixed volume
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Equivalence point vs end point
Equivalence point | End point |
Theoretical
*can be determined with pH probe/titration curve | Experimental through titration |
Complete neutralisation point | |
acid and base mixed according to stoichiometric molar ratio & pure salt + water (no acid/base)
*not neutral point/pH 7 b/c resulting salt might be acid or basic depending on parent acid and base | Point of colour change, almost identical to equivalence point
*indicator needs to be picked such that the indicator’s pH range matches the equivalence point of the reaction |
Acid rain
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The atmosphere contains CO2 which can dissolve in water to form carbonic acid:
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CO2 + H2O ⇄ H2CO3
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Which partially dissociates to form H+ and HCO3-
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H2CO3 ⇄ H+ + HCO3- ~pH=5.6
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Causes of acid deposition (Sulphur & nitrogen oxides)
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Acid deposition is when the pH of rain drops below 5.6 due to anthropogenic sources of pollution, namely SO2 and NOx gases.