acids bases salts

note: for IB, use reversible arrows as long as WA/WB is involved. only use -> when it is a salt that is dissociating fully.

acids properties

• pH 1 to 7 @ 25C

• releases H+ when dissolved in water

• sour

• reacts with bases to form salt + water

• reacts with carbonates to form salt + water + CO2

• reacts with metals to form salt + H2

arrhenius theory

• acid: has H in formula and produces H+ ions when dissolved in water

• base: has OH in its formula and produces OH- ions when dissolved in water

• limitations

  • restricted to reactions that occur in aqueous solutions (excludes gases eg HCl + NH3 → NH4Cl)

  • many substances that release OH- in water do not have OH in the formula (eg ammonia, amines)

bronsted-lowry theory

• acid: proton donor (ie: must have H)

• base: proton acceptor (need not have H)

• involves transfer of proton from acid to the base (need not be in water)

  • if acid reacts with alkali, it acts as bronsted-lowry acid

  • if acid reacts with stronger acid, it acts as bronsted-lowry base

• conjugate acid-base pairs differ by a proton (H+)

  • acid donates proton to become conj base

  • base accepts proton to become conj acid

  • there is an acid and base on both sides of the reaction. system reaches state of equilibrium based on relative strengths of acid, base and their conjugates.

amphiprotic

species that can act as both bronsted-lowry acid and bronsted-lowry base

polyprotic

acid that has many protons (H+)

lewis theory

• acid: electron-pair acceptor, to form dative covalent bond

• base: electron-pair donor, to form dative covalent bond

• product is called adduct (a species that contains a new dative covalent bond)

  • all atoms in the adduct have octet electron configuration

• example: complex ion formation

  • lewis base: ligand (electron-rich, electron pair donor)

  • lewis acid: central metal cation (electron-deficient, electron pair acceptor)

• widest scope: doesn’t need to have H or OH

alkalis

• bases that are soluble in water

• pH>7 @ 25C

• releases OH- ions when dissolved in water

• bitter

• reacts with acids to form salt + water

• reacts with ammonium salts to give salt + water + NH3 (g)

  • note: this is an acidic salt, since NH4+ + conjugate acid → NH3 (conjugate base)

acid deposition

• process by which acid forming pollutants (NOx, SO2) are deposited on earth’s surface

• effects

  • deforestation

  • leaching of minerals from soils → increasing acidity of lakes and rivers + uptake of toxic materials from soil by plants/shellfish

  • corrosion of limestone buildings

pH of rainwater

• naturally around 5.6

• CO2 (g) + H2O (l) <=> H2CO3 (aq) ← carbonic acid (weak acid)

• H2CO3 (aq) <=> H+ (aq) + HCO3- (aq)

• HCO3- (aq) <=> H+ (aq) + CO3 2- (aq)

oxides of nitrogen

• formation: high temperature in internal combustion engines, nitrogen and oxygen in air react to form nitrogen monoxide → which further reacts with oxygen from atmosphere to form nitrogen dioxide

  • N2 (g) + O2 (g) → 2NO (g)

  • 2NO (g) + O2 (g) → 2NO2 (g)

• NO2 effects

  • brown smog

  • acid rain: 2NO2 (g) + H2O (l) → HNO3 (aq) + HNO2 (aq)

    then oxidised by oxygen in atm: 2HNO2 (aq) + O2 (g) → HNO3 (aq)

sulfur dioxide

• from volcanic eruptions and combustion of fossil fuels with sulfur impurities

• formation of sulfurus acid: SO₂ (g) + H₂O (l) <=> H₂SO₃ (aq)

  oxidised by oxygen in atm to form sulfur trioxide: SO₂ (g) + O₂ (g) <=> 2SO₃ (g)

  then dissolves in rainwater to form sulfuric acid: 2SO₃ (g) + O₂ (g) → SO₄ ^2-

ocean acidification

• 50% of CO2 produced by fossil fuels is dissolved by oceans → ocean becomes more acidic, inhibits shell growth + causes reproductive disorders in fish

• CO2 dissolves in rainwater to form carbonic acid: CO2 (g) + H2O (l) → H2CO3 (aq)

• carbonic acid forms hydrogencarbonates/carbonates

  • H2CO3 (aq) <=> HCO₃^- (aq) + H+ (aq)

  • H2CO3 (aq) <=> CO₃^2- (aq) + H+ (aq)

self-dissociation of water

• partial dissociation: most remains as H2O molecules → since partial, can use reversible arrow + write eq constant (Kc)

• H2O (l) + H2O (l) <=> H3O⁺ (aq) + OH^- (aq)

  • note: this is an acid-base reaction where one H2O is acid and the other is base

• endothermic reaction

  • when temp increases, POE shifts right (favour endothermic) to oppose the temp increase → Kw increases as temperature increases

Kw

Kw = [H+][OH-]

Kw = Ka X Kb

pKw

pKw = pH + pOH

[H+] in strong acids

[H+] = [SA] + [H+]_H2O

(if significant ie [SA] <10^-7 @ 25C)

strong acids

• undergo complete dissociation in water → high concentration of H3O+ ions

• conjugate base is neutral since Kb << Kw

• examples: HNO3, H2SO4, hydrogen halides (except HF), HClO4

  • acid strength of hydrogen halides increase down group 17: halogen atom radius increase, H-X bond length increase so weakened → less energy needed to break to form H+

neutral solution

[H+] = [OH-]

strong bases

• undergo complete dissociation in water → high concentration of OH- ions

• conjugate acid is neutral since Ka << Kw

• examples: group 1/2 metal hydroxides

weak acids

• undergo partial dissociation in water since low tendency to donate proton (use reversible arrow) → low concentration of H3O+ ions

  • definition is independent of pH (concentration), it depends on pKa (extent of dissociation)

• conjugate base of a WA is a WB

  • the weaker the WA, the stronger its CB (but they’re all still weak) → POE lies towards the acid

weak bases

• undergo partial dissociation in water since low tendency to accept proton (use reversible arrow) → low concentration of OH- ions

  • definition is independent of pH (concentration), it depends on pKb (extent of dissociation)

• conjugate base of a WB is a WA

  • the weaker the WB, the stronger its CA (but they’re all still weak) → POE lies towards the base

experimental determination of acid/base strengths

• enthalpy change of neutralisation is less exothermic for WA-SB or SA-WB reactions

  • WA/WB only partially dissociates → less energy needed to break bonds

• electrical conductivity lower for WA/WB

  • WA/WB only partially dissociates → lower concentration of mobile ions

• rate of reaction lower for WA/WB

  • WA/WB only partially dissociates → concentration of H+/OH- lower → rate = k[H+][OH-]

• pH: at same temperature and same concentration, pH of WA higher than SA , pH of WB lower than SB (related to concentration of H+/OH- ions)

Ka

Ka = [H+]²/[HA]

• acid dissociation constant: extent to which acid dissociates in water → measure of strength of acid

  • if ask for which eqn gives expression, chose the one involving water

• the larger the Ka, the stronger the acid

• the larger the pKa value, the weaker the acid

Assumes [HA] @ equilibrium = [HA] initial

Kb

Kb = [OH-]²/[B]

• measure of strength of base

• the larger the Kb, the stronger the base

• the larger the pKb, the weaker the base

salt hydrolysis

• when salt dissolved in water, undergoes further reaction with water to produce H3O+ or OH- ions, resulting in an acidic/alkaline solution.

• is a reversible reaction

occurs for salts that contain

• an anion that is a CB of a WA

  • WB + SA → acidic salt (WA) + neutral

  • WA + H2O <=> WB + H3O+

  • Ka = [WB][H+]/[acidic salt] <1

• a cation that is a CA of a WB

  • WA + SB → basic salt (WB) + neutral

  • WB + H2O <=> WA + OH-

  • Kb = [WA][OH-]/[basic salt] <1

• a cation of high charge density (Al³⁺, Cr³⁺, Fe³⁺)

  • high charge density → form coordinate bond with ligands to form complex ion

  • high charge density of metal cation polarises H2O → H2O pulls electrons closer to itself, O-H bond breaks

  • H+ released, complex ion acts as a weak acid when hydrolysed → therefore salt with a transition metal is an acidic salt.

  • [Al(H₂O)₆]³⁺ + H₂O <=> [Al(H₂O)₅(OH)]²⁺ + H₃O⁺

explain whether 0.1M NaNO3 solution is acidic, alkaline or neutral

NaNO3 → Na+ + NO3-

NaNO3 formed from strong acid HNO3 and strong base NaOH. neutral since Na+ and NO3- do not undergo hydrolysis.

explain whether 0.1M Na2CO3 solution is acidic, alkaline or neutral

Na2CO3 → 2Na+ + CO3 2-

Na2CO3 is formed from strong base NaOH and weak acid H2CO3

Na+ ions do not undergo hydrolysis because metal cations do not react with water

CO3 2- + H2O <=> HCO3- + OH- (WB + H2O <=> WA + OH-)

the conjugate base (CO3 2-) of the weak acid undergoes hydrolysis to form basic solution due to OH- formed

(recall if salt has anion that is CB of WA, then it undergoes hydrolysis)

buffer solutions

• solutions that maintain an approximately constant pH when small amonts of either acid or base are added/when solution is diluted

• has acidic component and basic component

• formed by direct 1:1 mixing of WA/WB + conjugate

  OR 2:1 excess weak + limited strong

how does acidic buffer work?

• salt fully dissociates (→) to give large amount of conj base anions (WB) + cation from SA

• WB reacts with H+ from forward reaction to produce conj acid (the WA we started with)

• this backward reaction suppresses the forward dissociation reaction (of the WA)

when small amount of H+ added, CB + H+ <=> CA. added H+ is removed as CA, [H+] remains constant.

when small amount of OH- added, CA + OH- <=> CB. added OH- is removed as CB, [OH-] remains constant.

pH = pKa + lg [salt]/[acid]

pH points to consider for pH titration curves

1. initial pH → SA/SB/WA/WB

2. equivalence pH @ Veq → gradient vertical

   pH = pOH

   • pH = 7: SA + SB

   • pH < 7: SA + WB

   • pH > 7: WA + SB

3. max buffer pH @ ½ Veq if before end point / 2 Veq if after end point → gradient horizontal

   pH = pKa (if acidic buffer) / pOH = pKb

   • pH < 7: WA buffer (excess WB + limited SA)

   • pH > 7: WB buffer (excess WA + limited SB)

factors influencing max buffer

• dilution: affects buffer capacity

  • reduces concentration of components → lowers buffer capacity

  • doesn’t change ratio → does not change pH of buffer solution

• temperature: affects Ka/Kb → changes pH of buffer solution

pH curve for SA SB

1. low initial pH due to SA

2. pH changes gradually until equivalence point

3. very sharp jump in pH at equivalence point → pH @ equivalence point = 7

4. after equivalence point, curve flattens out at high pH due to SB

pH curve for SA WB

1. low initial pH due to SA

2. pH changes until equivalence point reached

3. very sharp jump in pH at equivalence point → pH @ equivalence point < 7

4. after equivalence point pH increases gradually due to basic buffer (max buffer @ 2Veq)

pH curve for WA SB

1. high initial pH due to WA

2. pH changes gradually until equivalence point due to acidic buffer

3. very sharp jump in pH at equivalence point → pH @ equivalence point > 7

4. after equivalence point the curve flattens out at high pH due to SB

pH curve for WA WB

1. high initial pH due to WA

2. addition of base causes pH to rise steadily

3. change in pH at equivalence point is not sharp → unable determine

4. after equivalence point the curve flattens out at low pH due to WB

maximum buffer capacity

amount of base/acid that ca be absorbed without significant changes to the pH

indicators: definition, how it works, explain useful range

• acids/bases that undergo neutralisation to change colour

• usually weak organic acids with complex structures

  • HInd is one colour

  • Ind- (conjugate base) is another colour

• HInd <=> H+ + Ind-

  • if [H+] increases, POE shifts left, forms more HInd

  • if [H+] decreases, POE shifts right, forms more Ind-

  • when orange, red = yellow → [HInd] = [Ind] → max buffer

• useful pH range for colour change of indicator is pKa (indicator) +- 1

  • when [Ind-]/[HInd] = 1/10, will start to see colour change

  • to measure useful range: add alkali/acid 1.0cm³ at a time to solution containing indicator + buffer, and continuously measure pH using pH meter → pH at start of colour change + end of colour change is the range (need buffer so that pH does not change too rapidly when add 1.0cm³)

choosing indicator

useful range should cover sharp pH change at eq point

1. determine whether reacting SA/SB/WB/WA → pH of equivalence point

2. if pH = 7, any indicator; if pH < 7, choose pKa <7; if pH > 7, choose pKa > 7

explaining choice: “active pH range of indicator coincides with the sharp change in pH at the equivalence point”

note: indicator range falls within actual range, not the other way around

how to determine acid requiring greatest volume of base for complete reaction?

determine total [H+], which depends on [acid] X basicity

ie 0.1M H3PO4 > 0.1M H2SO4 regardless of given pH (measure of H+ at a point)