Inorganic Compounds: Acidic Oxides

• Inorganic compounds are substances that encompass all elements and mineral compounds, not including carbon, with the exception of CO2, carbonates, carbides like SiC, and are commonly present in silicates, oxides, carbonates, sulphides, sulphates, chlorides, and nitrates.
• Compounds are classified into four groups: oxides, acids, bases, and salts.

Oxides

• Oxides form through direct reactions with nearly all elements except noble gases and inactive metals (gold, platinum, palladium).
• A binary compound containing oxygen and another element (metal, non metal, or metalloid).
• Oxide + Element = Oxide

Types of Oxides

• Acidic oxide
• Basic oxide
• Neural oxide
• Peroxide
• Amphoteric oxide

Acidic Oxides

• Acidic oxides react with oxygen and nonmetals.
• Also known as acid anhydrides, which mean acid without water.
• Non-metal + Oxygen = Non-metal oxide
• Example reactions include N2 + O2 -> NO2 and C + O2 -> CO2.

Chemical Properties of Acidic Oxides

• Acidic oxides dissolve in water and produce an acidic solution.
• Acid anhydride + Water = Acid
  • Example: CO2 + H2O -> H2CO3 (Carbonic acid)
  • Example: SO2 + H2O -> H2SO3 (Sulphurous acid)
• Acidic oxides react with basic or metallic oxides to produce salt.
• Acidic oxide + Basic oxide = Salt
  • Example: CO2 + Na2O -> Na2CO3 (sodium carbonate)
  • Example: SO3 + CaO -> CaSO4 (calcium sulphate)
• Acidic oxides react with bases to produce salt and water in a neutralization reaction.
• Acidic oxide + Base = Salt + Water
  • Example: SO2 + 2NaOH -> Na2SO3 + H2O
  • Example: CO2 + 2LiOH -> Li2CO3 + H2O

Basic Oxides

• Basic oxides react with oxygen and metals and are also known as basic anhydrides.
• Basic anhydride means base without water; examples include CaO, MgO, and Na2O.
• Metal + Oxygen = Metal oxide
• Most metal oxides exhibit basic properties, dissolving in water to form alkaline solutions.
• Some metal oxides (Al2O3 and ZnO) are amphoteric.
• Oxides of active metals in Group IA and Group IIA dissolve in water.
• Examples include Li2O, Na2O, K2O, MgO, CaO, BaO, and CuO.

Chemical properties of basic oxides

• Basic oxides dissolve in water to form alkaline solutions:
  • Basic anhydride + Water = Base (alkali)
    • Example: Li2O + H2O -> 2LiOH (lithium oxide)
    • Example: CaO + H2O -> Ca(OH)2 (calcium hydroxide)
• Basic oxides react with acidic or non-metallic oxides to form salt:
  • Acidic oxide + Basic oxide = Salt
    • Example: BaO + SO3 -> BaSO4
    • Example: Na2O + CO2 -> Na2CO3
• Basic oxides react with acids to form salt and water:
  • Basic oxide + Acid = Salt + Water
    • Example: CaO + 2HCl -> CaCl2 + H2O
    • Example: CuO + H2SO4 -> CuSO4 + H2O

Amphoteric Oxides

• These oxides show both acidic and basic behaviors.
• Examples: Al2O3, ZnO, PbO, PbO2, SnO, and SnO2.
• Reaction with acid: Amphoteric oxide + Acid -> Salt + Water
  • Example: Al2O3 + HCl -> AlCl3 + H2O
• Reaction with base: Amphoteric oxide + base -> Salt + Water
  • Example: Al2O3 + NaOH -> NaAlO2 + H2O
• Acidic oxides, basic oxides, and amphoteric oxides are salts.

Neutral Oxides

• These oxides don't react with acids or bases to form salt and water and do not show basic or acidic properties.
• Examples: H2O, CO, N2O, and NO.

Peroxides

• Oxides contain oxygen with an oxidation number of -1.
• This contains two oxygen atoms linked together with other elements, with a "– O – O –" link.
• Examples: H2O2, Na2O2, CaO2, BaO2, and SrO2.

Chemical Properties of Peroxides

• Peroxides act as powerful oxidizing agents, reacting by losing oxygen.
  • Example: PbS (s) + 4H2O2 (aq) → PbSO4 (s) + 4H2O (l)
  • Example: 2KI (aq) + H2SO4 (aq) + H2O2 (aq) → I2 (s) + K2SO4 (aq) + 2H2O (l)
• Peroxides react with aqueous acids to form hydrogen peroxide.
  • Example: Na2O2 (s) + 2HCl (aq) → 2NaCl (aq) + H2O2 (aq)
  • Example: CaO2 (s) + H2SO4 (aq) → CaSO4 (s) + H2O2 (aq)

Hydrogen Peroxide (H2O2)

• A colorless liquid used as bleach and antiseptic
• Solutions can be used as a bleach, an antiseptic, and for cleansing wounds.
• Decomposes with catalysts (MnO2 or Pt) added
  • The reaction is represented as: 2H2O2 (aq) -> 2H2O (l) + O2 (g)
• A strong oxidising agent
• Hydrogen peroxide added to paint containing lead sulphide turns it black.
  • PbS (s) + 4H2O2 -> PbSO4 + 4H2O (1)

Acids

• The term "acid" originates from the Latin word "acidus".
• Sours are abundant in citrus fruits and some vegetables.
• Examples:
• Acetic acid is found in vinegar
• Lemons have citric acid
• Formic acid is present in ant bites
• Buttric acid is contained in rancid butter
• Gastric juice has HCl acid

Arrhenius Definition of Acids

• Svante Arrhenius, a Swedish Chemist, suggested the definition of acids.
• An acid releases hydrogen ions (H⁺) or hydronium ions (H3O+) in aqueous solution: HA(aq) -> H+(aq) + A– (aq) HA(aq) + H2O(l) -> H3O+(aq) + A–(aq)
• Examples:
• HCl(aq) -> H+(aq) + Cl–(aq)
• H2SO4(aq) -> 2H+(aq) + SO4 2-(aq)
• Arrhenius acids: HNO3, H2SO4, H3PO4, HBr, HI and HF

Classification of Acids

• Acids are classified based on the number of replaceable hydrogen ions (H+) or hydronium ions (H3O+ ions) per molecule: mono-protic, di-protic, and poly-protic.

Mono-protic

• Monoprotic acids contain only one ionizable hydrogen atom per molecule.
  • Example: HCl(aq) -> H+ (aq) + Cl– (aq)

Di-protic

• Diprotic acids contain two ionizable hydrogen atoms per molecule.
  • Examples: H2SO4 (aq) -> 2H+ (aq) + SO42–

Poly-protic

• Polyprotic acids contain three ionizable hydrogen atoms per molecule.
  • Example: H3PO4 (aq) ⇌ 3H+ (aq) + PO4 3– (aq)

Classification of Acids by Constituent Elements

• Based on the number of constituent elements acids can be classified as binary and ternary.

Binary Acids

• Composed of two elements: hydrogen and a non-metal.
  • Examples: HCl, HBr, HF, HI, and H2S.

Ternary Acids (Oxy-acids)

• Composed of three elements: hydrogen, oxygen, and a non-metal.
  • Examples: H2SO4, H2CO3, HClO4, and H3PO4.

General Properties of Acids

I. Acids have a sour taste. II. Acids change the color of certain acid-base indicators. III. Acids react with active metals to yield hydrogen gas.

• Moderately active metals (Mg, Zn, Fe and Al) form salts with H2 gas.
• Example: H2SO4 + Zn → ZnSO4 + H2
• Example: 6HCl + 2Al → 2AlCl3 + 3H2
• Concentrated HNO3 and hot concentrated H2SO4 react with copper, producing NO2 and SO2 gases, respectively.
  • Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O
  • Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O IV. Acids react with MCO3 and MHCO3 to form salt, H2O and CO2 gas.
• Acid + Hydrogen carbonate → Salt + Water + Carbon dioxide
• Acid + Carbonate → Salt + Water + Carbon dioxide
  • Example: 2HCl + CaCO3 → CaCl2 + H2O + CO2
  • Example: HCl + NaHCO3 → NaCl + H2O + CO2 V. Concentrated acids react with sulphites to form salts with liberation of SO2 gas.
• Sulphite + Acid → Sulphur dioxide + Salt + Water
  • Example: CaSO3 + H2SO4 → SO2 + CaSO4 + H2 VI. Aqueous solutions of acids are electrolytes and conduct electrical current.
• Acids are strong electrolytes (ionize completely in water).
  • Examples: HCl, HNO3 and H2SO4
• Acids are weak electrolytes (exist primarily in a non-ionized form when dissolved in water).
  • Examples: H2CO3, H3PO4, CH3COOH.

Strength of Acids

• Based on their degree of dissociation, acids are divided into strong acids and weak acids.

Strong Acid

• Acids completely dissociate in an aqueous solution:
  • Example: HCl (aq) → H+ (aq) + Cl– (aq)
• Examples: HCl, HNO3 and H2SO4

Weak Acid

• Acids partially dissociate in an aqueous solution.
  • Examples: acetic acid (CH3COOH), oxalic acid (H2C2O4), carbonic acid (H2CO3).

Concentrated and Dilute Acids

• Acids are classified based on the number of moles: concentrated and dilute acids.

Concentrated Acid

• The solute is a relatively large amount.
  • Example: 96% H2SO4 and 4% H2O.

Dilute Acid

• The solute is relatively small.
• Example: 96% H2O and 4% H2SO4 acid.

Precautions in Handling Acid

• Wear goggles, gloves and a laboratory coat.
• Wash the affected part if acid is spilled on your body with running water and 10% Na2CO3 solution.
• Be aware of what way to dilute concentrated acids (pour acid in to water)
• Administer weak bases if corrosive acids are swallowed (Mg(OH)2 or Al(OH)3).
• Use bellows to pipette acids.
• Wash eyes thoroughly if acids enter eyes.
• Wash off spilled acid.

pH and pH Scale

• pH measures the acid strength of a solution and ranks solutions of acidity or basicity.
• Acidic pHs have a larger hydrogen ion concentration.
• Kw = [H+][OH–] = 1.0 x 10–14 at 25 °C
• [H+] = [OH–] or [H+]2 = 1.0 × 10–14
• [H+] = 1.0 ×10–7 at 25 °C.
• pH = - log[H+]
• The ranges of pH scale is from 0-14.
• This can be used to calculate the pH of water: pH = - log[H+] = -log (1x10–7) = 7 and similarly, pOH = - log[OH–] = -log(1x10–7) = 7 and thus pH + pOH = 14

The Relationship between pH and Concentration

• A solution of pH 1 has 10 times higher hydrogen ion conc. than a solution of pH 2.
• A solution of pH 2 has 100 times higher hydrogen ion conc. than a solution of pH 3.

Preparation of Acids

• Acids can be prepared:
• Reaction of oxides of non metals (acidic oxides) and water (Acidic oxide + Water = Acid)
  • Example: N2O5 (s) + H2O (l) = 2HNO3 (aq)
  • Example: P4O10 (s) + 6H2O (l) = 4H3PO4 (aq)
• Direct combination of non metals like S and Cl with hydrogen to prepare binary acids (Non-metal + Hydrogen is Binary acid)
  • Example: H2 (g) + Cl2 (g) = 2HCl (g)
  • Example: H2(g) + S (s) = H2S (g)
• Heating acids with a non volatile acid to prepare volatile acids
  • Example: NaCl (s) + H2SO4 (l) → NaHSO4 (s) + HCl (l)
  • Example: NaNO3 (s) + H2SO4 (l) → NaHSO4 (s) + HNO3(l)

Use of Important Acids

Sulphuric Acid (H2SO4)

• Used in the production of Sulphate and phosphate fertilisers etc
• Used to refine petroleum
• Used in the production of metals
• Used in car batteries

Hydrochloric Acid (HCl)

• Is in the gastric juice of our body and helps in the digestion of food
• Used industrially to remove surface impurities before galvanizing
• Used to produce dyes, drugs photographic films and plastics like Polyvinyl chloride (PVC)
• Used to recover magnesium from seawater

Nitric Acid (HNO3)

• Used industrially to manufacture: Explosvies such as trinitrotoluene (TNT) and trinitroglycerine (TNG) Fertilisers eg KNO3 and NH4NO3
• Used in rubber, chemicals, plastics, dyes and drugs.

Bases

• According to the Arrhenius theory, a base produces a hydroxide ion (OH-) when it is dissolved in water., thus are soluble and called alkalis
• Examples include; KOH, NaOH, Ca(OH)2, NH3,
• Bases react with acids and acidic oxides to form salt and water:
  • Example: NaOH (aq) + HCl (aq) -> NaCl (aq) + H2O (I)
  • Example: 2NaOH (aq) + CO2 (aq) -> Na2CO3 (aq) + H2O(l)
• Strong/Weak Bases (depending on the number of solute)
• Concentrated has a relatively large amount of solute dissolved in a small amount of solvent
• Example: In a solution, if there is 96% NaOH and 4% H2O.
• Dilute has a relatively smaller amount of solute dissolved in a large amount of solute
• Example: In a solution, if there is 96% H2O and 4% Mg(OH)2 base.

Salts

• Salts form by the neutralization of an acid with a base.
• (Na+ OH) + (H+Cl-) → (Na+Cl-) + H2O

Classification of Salts

• Types include normal salt, acidic salt, and basic salt

Normal Salts

• Salts that reacts between strong acid and strong base
  • Example
    • HCl(aq) + KOH(aq) → KCl(aq) + H2O(l)
    • H2SO4(aq) + ZnO(aq) → ZnSO4(aq) + H2O(l)
    • HNO3(aq) + NH4OH(aq) → NH4NO3(aq) + H2O(l)

Acidic Salts

• Salts are formed by the partial replacement of replaceable H+ ions of an acid by a positive metal ion.

Basic Salts

• Salts is which not all of the hydroxide ions in a base have been replaced by the anions of the acid

General Methods for the preparation of salts

Preparation of Insoluble Salts Double decomposition reaction, (creates a precipitate) where:

• AgNO3(aq) + NaCl (aq) → AgCl(s) + NaNO3(aq)
• Na2CO3(aq) + CuSO4(aq) → CuCO3 (s) + Na2SO4 (aq)

Salts Usage

1. NaCl usage

• Food and preservation
• Raw material for the manufacture of Na, Cl2, and NaOH ORS component
• Na2CO3 Manufacture

2. Ammonium nitrate,(NH4NO3) usage

• Nitrogen Fertilizer and Explosives

3. Copper (II) sulphate,(CuSO4) usage

• Bordeaux mixture with Ca(OH)2 is created from these and is applied to prevents leaves and vines fungal attacks
• Electroplating and dyeing

4. iron (III) chloride. FeCl3 usage.

• Waste waster treatment and etching printed circuits

5. Potassium nitrate (KNO3) Usage.

• Gunpowder and fertiliser

6. Calcium Sulphate, uses:

• Supports fractured bones, as an agent plastering walls,

7. Barium sulphate "barium meal" before gastrointestinal x-ray photography.

Properties of Salts

• They either absorb water from the atmosphere of release it depending on the type. Type of Salts I. Hygroscopic
  • Salts absorb water from the atmosphere
  • Anhydrous example; copper (II) sulphate is used,

II. Deliquescent

• Water is absorbed and the solid dissolves
• CaCl2 + Sodium nitrate(NaNO3) example

III. Efflorescent

• Loss of crystallisation through solid crystals in the atmosphere

Thermal stability of Salts

Thermal Decomposition:

• MCO3(s) = MO(s) + CO2 Carbonate
• Group IIA's carbonates are: MgCO3, CaCO3, SrCO3, BaCO3,
• Li2CO3(s) Decomposition = Li2O(s)+CO2,

Nitrate

• Metal oxide is created through nitrates decomposition 2Mg(NO32 (s) Decomposition) = 2MgO (s) + 4NO2 (g) + O2 (g)
  2Pb(NO3)2 (s Decomposition) = 2PbO (s) + 4NO2 (g) + O2 (g)
• Sodiums and potassiums decompose but their Nitrates are applied to create gas. 2NaNO3 Decopositon to create 2NaNO2 (s) + O2 (g) 2KNO3 Decomposition to 2KNO2 (s) + O2 (g) Plant decomposition in the presence of lithium varies to this to create product: LiNO3 through decomposition of 2Li2O oxide gas through Li2CO3 carbonate gas.
• Group nitrates and metals composition differs in this senario with the presence of
• 4LINO2s
  • Li2CO2s

Plant based decomposition

Macronutrients are what is needed in elements for a large decomposition composition. These consist of nitrogen, phorphorus, magnesium, sulphur, potassium, carbon.

Fertilizers

• Applied to soil to enhance growth and nutritional composition through the utilization of synthetic and organic forms
• Organic from animal and plant based composition
• Nitrogen phosphorous that composes decomposition NPK - for fertility including elements. ammonium sulphate used in acidic soils that kill the plant's growth and development