Acids and redox

Chapter 4

Acids

Acids dissociate in solutions and release H+ ions

Common acids:
HCl
H2SO4
HNO3
CH3COOH (weak organic acid)

Alkali

Alkalis dissociate and release OH+ ions into solutions (soluble base)

Common alkalis:
NaOH
KOH
NH3

Strong verses weak acids

Strong acids → completely dissociate in water to produce H⁺ ions

Weak acids → only partially dissociate in water, so an equilibrium is set up with mostly undissociated molecules present

Strength depends on the degree of dissociation, not concentration

Bases and neutralisations

Neutralisation is the reaction between H⁺ ions and base to form a salt + water but within alkalis it is between H+ and OH⁻ ions to form water:
H⁺ + OH⁻ → H₂O

A base is a substance that reacts with an acid to form a salt and water.
(Alkalis are bases that are soluble in water and produce OH⁻ ions in solution.)

Neutralization reactions

Acid + alkali → salt + water

Acid + metal oxide → salt + water

Acid + carbonate → salt + water + carbon dioxide

Preparing a standard solution

• Accurately weigh solid using balance

• Dissolve in beaker with distilled water

• Transfer solution to volumetric flask using funnel

• Rinse beaker and funnel into flask (quantitative transfer)

• Make up to calibration line with distilled water (bottom of meniscus)

• Stopper and invert to mix thoroughly

Acid–base titration

• Rinse burette with titrant, then fill it

• Remove air bubbles from burette tip

• Record initial burette reading

• Use pipette to transfer fixed volume of analyte into conical flask

• Add suitable indicator

• Swirl flask continuously during titration

• Add titrant slowly near endpoint (drop by drop)

• Stop at colour change (endpoint)

• Record final burette reading

• Repeat until concordant results obtained (0.1cm3)

• Calculate mean titre

Standard rules for oxidation numbers

• Atoms in elements have oxidation number 0

• A simple ion has oxidation number equal to its charge

• Sum of oxidation numbers in a neutral compound = 0

• Sum of oxidation numbers in a polyatomic ion = overall ion charge

Common rules in oxidation numbers

• Group 1 metals = +1

• Group 2 metals = +2

• Fluorine = –1 always

• Oxygen = –2 (except in peroxides where it is –1)

• Hydrogen = +1 (–1 in metal hydrides)

• Chlorine usually = –1, except in compounds with oxygen or fluorine

Roman numerals

• Roman numerals show the oxidation state of an element in a compound or ion when it can have more than one

• Used to distinguish different ions of the same element (e.g. different charges/oxidation numbers)

• For example: Fe(II) and Fe (III) corresponding to Fe2+ and Fe3+

Oxidation and reduction

• Oxidation: loss of electrons and/or increase in oxidation number

• Reduction: gain of electrons and/or decrease in oxidation number

Examples oxidation and reduction from S, P and D block

• s-block:
  Na → Na⁺ + e⁻ (oxidation)
  Mg → Mg²⁺ + 2e⁻ (oxidation)

• p-block:
  Cl₂ + 2e⁻ → 2Cl⁻ (reduction)
  O₂ + 4e⁻ → 2O²⁻ (reduction)

• d-block:
  Fe²⁺ → Fe³⁺ + e⁻ (oxidation)
  Cu²⁺ + 2e⁻ → Cu (reduction)

Redox reactions of metals with acids (forming salts)

Metals react with acids to form a salt + hydrogen gas

Redox reaction:

• Metal is oxidised (loses electrons, oxidation number increases)

• Hydrogen ions are reduced (gain electrons to form H₂)

Mg → Mg²⁺ + 2e⁻
2H⁺ + 2e⁻ → H₂
Mg + 2H⁺ → Mg²⁺ + H₂
Mg + 2HCl → MgCl₂ + H₂

Acid–Base Indicators

🌈 Phenolphthalein

• Acid: colourless

• Neutral: colourless

• Alkali: pink

🟠 Methyl orange

• Acid: red

• Neutral: orange

• Alkali: yellow

🟣 Universal indicator (UI)

• Strong acid: red

• Weak acid: orange/yellow

• Neutral: green

• Weak alkali: blue

• Strong alkali: purple

🔴 Litmus

• Acid: red

• Neutral: purple

• Alkali: blue