Chemistry: Module 6 - Acid Base Reactions

what is an inorganic Acids and give examples

simple and have a hydrogen and an ion: HCl, HNO₃, H₃PO₄

define Organic Acids

compounds that are organic but have acidic properties (e.g. alkanoic acids: have -COOH)

what is the nature of Alkalis

Bases soluble in water → include soluble hydroxides and ammonia (all Group 1 metals)

Monoprotic

Not all hydrogens in an acid molecule are ionizable: monoprotic = can donate one H⁺

define Diprotic

Can donate two H⁺

hydronium ion

Is the more stable version of an aqueous H⁺_(aq): H⁺_(aq)+ H₂O_(l) → H₃O⁺_(aq)

H⁺_(aq) = H₃O⁺_(aq)

properties of acids (6)

Produce H+ ions in water

Have a pH < 7

Sour taste

Conducts electricity in solution

Turns blue litmus red

Corrosive

acid dissociation equation

HA → H⁺ + A^-

base properties (6)

Produce hydroxide ions (OH-) in solution

pH > 7

Soapy feel, bitter taste

Conducts electricity in solution (if soluble)

Turns red litmus blue

Corrosive (caustic)

types of bases and an example (4)

Metal Hydroxides	Metal Oxides (non-metal oxides like CO2 and SO2 are acidic)	Metal Carbonates/Hydrogen Carbonates	Organic Bases (include amines: -NH2)
NaOH, Ba(OH)2	Na2O, CuO	K2CO3, NaHCO3	caffeine, nicotine, morphine

what is the nature of ammonia?

a weak base

metal oxides are either:

bases or amphoteric

what is the acetate ion and its nature?

a weak base, CH₃COO^-

what is amphoteric

• can act as either an acid or a base (Al₂O₃) → will be the opposite of whatever it is reacting with

base: Al₂O_{3 (s)} + 6HCl_(aq) → 2AlCl_{3 (aq)} + 3H₂O_(l)

acid: Al₂O_{3 (s)} + 2NaOH_(aq) → 2AlCl_{3 (aq)} + 2Na(Al(OH)₄)_(aq)

what is amphiprotic:

• Amphiprotic: can donate or accept a proton 

H₂O_(l) + HCl → H₃O⁺_(aq) + Cl^-_(aq) (water accepted the H⁺, HCl donated)

H₂O_(l) + NH_{3 (aq)} → NH₄⁺_(aq) + OH^-_(aq) (water donated H⁺, NH₃ accepted)

what is the purpose of an indicator and what is an indicator’s nature?

used to measure how acidic or basic something is

• An indicator in any substance that will be different colours in different pH conditions (qualitative change)

• Indicators are weak acids or bases = exist in two forms

what is a universal indicator?

• most commonly used → useful because it is a whole range of different colours depending on the pH - combination of indicators

purpose of the pH

• Describes the acidity of a solution [H⁺]

  • Acids have a high [H⁺] = low pH

  • Bases have a low [H⁺] = high pH

  • pH =7: neutral solution

pH formula

• pH = -log₁₀[H⁺]

how does an indicator work?

• Indicators are weak acids or bases = exist in two forms that are different colours due to its reversibility

HInd ⇌ H⁺ + Ind^-

• H is the hydrogen atom and Ind is the indicator - the rest of the organic molecule

• A chemical equilibrium exists in solution between HInd and Ind^-: and they are different colours

  • Qualitatively shows where the equilibrium lies

  • If a base is added, [H⁺] ↓, so eqm shifts right

  • If acid is added, [H⁺] ↑, so eqm shifts left

methyl orange’s pH range and colour at each pH

Methyl orange

red

3.1- 4.4

yellow

litmus’s pH range and colour at each pH

Litmus (vegetable extract)

red

4.5-8.3

blue

bromothymol blue’s pH range and colour at each pH

Bromothymol blue

yellow

6.0-7.6

blue

phenolphthalein’s pH range and colour at each pH

Phenolphthalein

colourless

8.2-10

pink

how to you prepare and test a natural indicator?

Preparing the Natural Indicator

1. Tear or slice one red cabbage leaf/flower petal/tea/fruit leaf into small pieces and place them in a beaker

2. Cover it in 250mL boiling water for 15 minutes

3. Pour the mixture through a strainer and collect the indicator solution in a clean beaker

Testing the Natural Indicator(observe colour of indicator in solutions with a KNOWN pH)

1. Pour indicator into 3 small test tubes until they are half-full

2. Add 1mL of 0.5 HCl into the first tube and agitate to mix (acid)

3. Add 1mL of 0.5 NaOH into the second tube and agitate to mix (base)

4. Add 1mL of water into the third tube and agitate to mix (neutral)

5. Record the colour of each solution

Acid + base

salt + water

acid + metal hydroxide

salt + water 

Acid + Metal Oxides →

salt + water

Acid + Metal Carbonate →

salt + water + carbon dioxide

Acid + hydrogen carbonate →

salt + water + carbon dioxide

Dilute Acid + reactive metal →

salt + hydrogen gas

• reactive metal is higher than hydrogen on the standard potential table

Acid + ammonia

ammonium salt

what is the nature of the ammonium ion

a weak acid

what is the nature of ammonia

weak base

what does enthalpy mean

• how much energy is found within a substance or groups of substances

  • the change of enthalpy (∆H) is measured in kJ/mol

  • ∆H = H(products) - H(reactants)

what is the enthalpy of neutralisation?

• change in enthalpy per one mole of H₂O formed

what is the net ionic equations of strong acids and bases?

• H⁺_(aq) + OH^-_(aq) → H₂O_(l)   ∆H_n = -57kJ/mol

what is the overall enthalpy of neutralisation of weak acids/bases?

• less exothermic, possibly endothermic

  • They don't completely dissociate in water

  • Heat produced is used to dissociate the weak acid/base into ions

  • ∆H _(observed) = ∆H _{(neutralisation)} + ∆H _{(dissociation of acid/base)}: negative + positive

how do you find the heat produced?

q=mc∆T

q= heat produced (J)

m = mass of reaction (kg)

c= specific heat capacity of water (4.18 x 10³)

∆T = change in temp of water (units don’t matter)

how do you find the change in enthalpy (formula)

∆H = -q/n

∆H=enthalpy change (kJ/mol)

q= heat produced (kJ)

n = moles of water produced

what is acid strength?

the extent of ionisation/dissociation in water (NOT based on the concentration - its the nature of the particle)

• Acids are covalent molecular compounds that split into ions (ionises) when it interacts with water

what do strong acids do?

Dissociation goes to completion/ionises further (→ )

HA(aq) + H2O(l) → H3O+(aq) + A-(aq)  HA(aq) →  H+(aq) + A-(aq)

Strong electrolytes (conductors of electricity in solution)

what do weak acids do?

Dissociation is reversible: reaches a dynamic equilibrium (⇌) % of products/reactants formed depends on the concentration

HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq)  HA(aq) ⇌ H+(aq) + A-(aq)

Weak electrolytes

what are the 5 main strong acids?

HCl, HBr, HI, H₂SO₄, HNO₃

what does a strong base do?

Strong Base ~ ionic

Dissociates completely in water when it dissolves, producing free OH- ions (→)

Strong electrolytes

NaOH(s) → Na(aq)+ + OH(aq)-

what does a weak base do?

Weak Base

Reacts partially with water to form hydroxide ions in solution (⇌) (include water in equation) If ionic: will dissociate into OH- Some weak bases are anions inside an ionic compound First dissolution to produce ions (→) Show anion producing OH- from water (⇌)

Weak electrolytes

Ionic NH3 (aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq)  Anions in an ionic compound Na2CO3 (s) → 2Na+(aq) + CO32-(aq) CO32-(aq) + H2O(l) ⇌ HCO3-(aq) + OH-(aq)

examples of strong bases

Group 1, 2 metal hydroxides: NaOH, KOH, Ca(OH)₂

what is K_a?

• Equilibrium constant for the dissociation (ionisation) of an acid into a hydrogen or hydronium ion and anion

• Indicates the extent of dissociation/strength of an acid

• The stronger the acid, the ↑ K_a (eqm lies on the right, more ionisation)

• Polyprotic acids have more than one K_a value (for each H⁺ they donate)

what is K_b?

Indicates the extent of dissociation/strength of base The stronger the base, the ↑Kb

water is both

amphoteric and amphiprotic (acts as both base and acid and can donate/accept a proton)

• Water can acts as either a B/L acid or a B/L base

give an example of water acting as an acid

• Molecular acids dissolve in water to produce ions -- according to the Bronsted/Lowry theory this occurs because a proton is donated from the acid to the water molecule to produce the hydronium ion

NH_{3 (g)} + H₂O_(l) ⇌ NH₄⁺_(aq) + OH^-_(aq)

give an example of water acting as a base

HCl_(g) + H₂O_(l) → H₃O⁺_(aq) + Cl^-_(aq)

• The HCl = B/L acid as it it a proton donor, water is the B/L base as it is a proton acceptor
  

pH =

-log₁₀[H⁺]

pOH =

-log₁₀[OH^-]

pH + pOH =

14 (for aqueous solutions at 25℃)

what does the pOH range look like?

what did Lavoisier propose for the definition of acids?

• Acids contain oxygen - the presence of oxygen gives acids their properties

• Based on compounds like HNO₃ and H₂SO₄

what is the limitations of Lavoisier’s definition

Couldn't explain acids that didn’t contain oxygen (HCl)

what is Davy’s definition of an acid?

• Stated acidity is due to the presence of hydrogen

• Showed that HCl is strongly acidic

what is Davy’s limitation in his definition?

• Doesn’t define a base

• Lacks an explanation of the chemical behaviour of bases

• Didn’t explain why other compounds that contained H weren’t acidic (CH₄)

what is arrhenius’ definition of acids and bases?

Acids as substances that increases the concentration of H+ ions in aqueous solutions HCl → H+ + Cl- Strongest acids are the most dissociated Bases are substances that  increase the concentration of OH- ions in aqueous solutions NaOH → Na+ + OH- Allowed for an explanation of neutralisation

what is the limitation of arrhenius’ definition?

• Only for aqueous solutions

• Doesn’t explain reactions in non-aqueous solvents or gas phase reactions

• Cannot explain bases that don’t contain OH^- like NH₃ 

• Couldn’t explain why some neutralisation reactions produced unneutral solutions

what is bronsted and Lowry’s definition of acids and bases?

Acids are proton (H⁺) donors

Bases are proton acceptors (H⁺) 

• Explains a wide range of acid-base behaviour

• Explains how ammonia works as a base because it accepts a proton

• Explains that ionic compounds are basic because they dissolve first and then ionise

what is the benefit of bronsted lowry defintions?

Works in both aqueous and non-aqueous environments (any solvent)

• Introduces the idea of conjugate acid-base pairs

NH_{3 (g)} + HCl_(g) → NH₄Cl_(s)

NH₃ + H₂O ⇌ NH₄⁺ + OH^-

what is the B/L base and the B/L acid

• NH₃ acts as a B/L base as it accepts a proton to become NH₄⁺

• H₂O acts as a B/L acid as it donates a proton to become OH^-

what is the limitation of B/L definition

• Doesn’t explain acid-base reactions that don’t involve proton transfer (Lewis acid-base theory) e.g. AlCl₃ and BF₃

• Can’t explain reactions between acidic oxides and basic oxides that occur without a solvent

• HOWEVER, it is still used because it is sufficient in covering most situations

what is the lewis acid-base theory?

Lewis acid: an electron pair acceptor

Lewis base: an electron pair donor

what is the order of acid/base definitions from old to new?

lavoisier, davy, arrhenius, bronsted-lowry, lewis

explain the self-ionisation of water at 25 degrees?

H₂O_(l) + H₂O_(l) ⇌ H₃O⁺_(aq) + OH^-_(aq)

• its pH = 7, so the [H⁺] = 10^-7 mol/L

• K_w = [H₃O⁺][OH^-] = 10^-14 (from data sheet)

• ∴ [OH^-]  = 10^-7

K_w =

[H₃O⁺][OH^-] = 10^-14 (at 25℃)

what is conjugate acid-base pairs?

• If something is an acid, something else is a base 

• When a B/L acid donates its proton to the base, the anion of the acid is proton deficient - this species can act as a base as it could accept a proton and reform the original acid

when writing a conjugate acid base pair what do you write first?

When writing the pairs, write the acid first: so here: H₂O/ OH^- and NH₃/ NH₄⁺

what is the base dissociation formula?

AB —> A⁺ + B^-

B^- + H₂O → HB + OH^-

what is the strength of an acid?

the degree of ionisation of an acid or dissociation of a base → can be determined by calculating the [H⁺] of equal concentration of acids