Chemical Reactions, Lewis Structures, Polyatomic Ions, and Acid Nomenclature

Chemical Reactions and Balancing Equations

Fundamental Law of Conservation of Mass:
- Chemical balancing is strictly based on the Law of Conservation of Mass.
- In any chemical reaction, mass is neither created nor destroyed.
- The total number of atoms of each element before the reaction (on the reactant side) must match the total number of atoms of each element after the reaction (on the product side).
Anatomy of a Chemical Equation:
- Reactants: The starting substances involved in a chemical reaction, written on the left side of the reaction arrow ( $\rightarrow$ ).
- Products: The new substances formed by the reaction, written on the right side of the reaction arrow ( $\rightarrow$ ).
- Coefficients: Whole numbers placed directly in front of chemical symbols or formulas (schematically indicated as boxes $\square$ ). Coefficients multiply every atom in the attached formula and are adjusted to balance the equation.
- Subscripts: Small numbers written below and to the right of atomic symbols within chemical formulas. Subscripts specify the exact ratio of atoms in a compound and can never be changed when balancing an equation.

Reactants and coefficients layout in a chemical equation

General visual reaction framework: $\square A + \square B \rightarrow \text{Products}$ where $\square$ represents the coefficient box.
- Detailed Worked Examples of Equation Balancing:
Example 1: Synthesis of Ammonia ( $NH_3$ ):
- Initial unbalanced equation: $\square N_2 + \square H_2 \rightarrow \square NH_3$
- Initial atom inventory:
  - Reactant side: Nitrogen ( $N$ ) = $2$ , Hydrogen ( $H$ ) = $2$
  - Product side: Nitrogen ( $N$ ) = $1$ , Hydrogen ( $H$ ) = $3$
- Step 1: Balance Nitrogen by placing a coefficient of $2$ in front of $NH_3$ : $\square N_2 + \square H_2 \rightarrow 2 NH_3$ .
  - Updated product count: Nitrogen ( $N$ ) = $2$ , Hydrogen ( $H$ ) = $2 \times 3 = 6$ .
- Step 2: Balance Hydrogen by placing a coefficient of $3$ in front of $H_2$ : $N_2 + 3 H_2 \rightarrow 2 NH_3$ .
  - Updated reactant count: Hydrogen ( $H$ ) = $3 \times 2 = 6$ .
- Final balanced equation: $N_2 + 3 H_2 \rightarrow 2 NH_3$
Example 2: Formation of Water ( $H_2O$ ):
- Initial unbalanced equation: $\square H_2 + \square O_2 \rightarrow \square H_2O$
- Initial atom inventory:
  - Reactant side: Hydrogen ( $H$ ) = $2$ , Oxygen ( $O$ ) = $2$
  - Product side: Hydrogen ( $H$ ) = $2$ , Oxygen ( $O$ ) = $1$
- Step 1: Balance Oxygen by placing a coefficient of $2$ in front of $H_2O$ : $\square H_2 + \square O_2 \rightarrow 2 H_2O$ .
  - Updated product count: Oxygen ( $O$ ) = $2$ , Hydrogen ( $H$ ) = $2 \times 2 = 4$ .
- Step 2: Balance Hydrogen by placing a coefficient of $2$ in front of $H_2$ : $2 H_2 + O_2 \rightarrow 2 H_2O$ .
  - Updated reactant count: Hydrogen ( $H$ ) = $2 \times 2 = 4$ .
- Final balanced equation: $2 H_2 + O_2 \rightarrow 2 H_2O$
Example 3: Reaction of Magnesium Hydroxide and Hydrochloric Acid:
- Formula equation: $1 Mg(OH)_2 + 2 HCl \rightarrow MgCl_2 + 2 H_2O$
- Reactant side inventory table breakdown:
  - Magnesium ( $Mg$ ): $1$
  - Oxygen ( $O$ ): $2$ (from $(OH)_2$ )
  - Hydrogen ( $H$ ): Initially counted as $3$ , updated to $4$ ( $2$ from $(OH)_2$ plus $2$ from $2 HCl$ )
  - Chlorine ( $Cl$ ): Initially counted as $1$ , updated to $2$ (from $2 HCl$ )

Lewis Structures of Ionic and Covalent Compounds

Comparison of Neutral Atoms and Ions:
- Lithium Atom ( $Li$ ) vs. Lithium Ion ( $[Li]^+$ ):
- Neutral Lithium atom ( $Li$ ) has $1$ valence electron in its outer shell.
- Lithium ion ( $[Li]^+$ ) loses its $1$ valence electron, leaving a positively charged cation enclosed in square brackets with a charge of $+1$ : $[Li]^+$ .
- Oxygen Atom ( $O$ ) vs. Oxygen/Oxide Ion ( $[O]^{2-}$ ):
- Neutral Oxygen atom ( $O$ ) has $6$ valence electrons.
- Oxide ion ( $[O]^{2-}$ ) gains $2$ valence electrons to complete a full octet ( $8$ electrons), forming a negatively charged anion enclosed in square brackets with a charge of $-2$ : $[:\ddot{O}:]^{2-}$ .
Lewis Structures for Ionic Compounds:
- Ionic compounds form via electron transfer from a metal to a non-metal.
- Magnesium and Oxygen ( $Mg + O$ ):
- Neutral Magnesium ( $Mg$ ) has $2$ valence electrons; neutral Oxygen ( $O$ ) has $6$ valence electrons.
- Magnesium transfers both of its valence electrons to Oxygen.
- Resulting ionic structure: Magnesium cation $[Mg]^{2+}$ and Oxide anion $[:\ddot{O}:]^{2-}$ .
- Potassium and Sulfur ( $K + S$ ):
- Neutral Potassium ( $K$ ) has $1$ valence electron; neutral Sulfur ( $S$ ) has $6$ valence electrons.
- Two Potassium atoms are required to donate $1$ valence electron each ( $2$ electrons total) to Sulfur to fill Sulfur's valence shell to an octet.
- Resulting ionic structure: $[K]^+ [:\ddot{S}:]^{2-} [K]^+$ or $2 [K]^+ [:\ddot{S}:]^{2-}$ .
Lewis Structures for Covalent Compounds:
- Covalent bonding involves electron sharing between non-metal elements.
- Valence shell geometry: Each valence level has $4$ sub-orbitals (top, bottom, left, right), holding up to $8$ total valence electrons (octet rule).
- Step-by-Step Methodology for Drawing Covalent Lewis Structures:
1. Count all valence electrons: Sum the total valence electrons supplied by all reacting elements in the compound.
2. Form single bonds: Place $2$ shared electrons (represented as a single line bond) between each pair of bonded elements.
3. Fill valence shells: Use remaining electrons to complete the valence shells (octets) of outer/terminal atoms first, then place left-over electrons on the central atom as lone pairs.
4. Form multiple bonds if necessary: If any atom lacks a full valence octet, share more electron pairs from adjacent atoms to form double or triple bonds.
Detailed Covalent Compound Examples:
- Carbon Tetrabromide ( $CBr_4$ ):
- Total valence electron count: $4 \text{ (from Carbon)} + 4 \times 7 \text{ (from Bromine)} = 4 + 28 = 32$ valence electrons.
- Structure: Central Carbon atom single-bonded to $4$ surrounding Bromine atoms.
- Electron allotment: $4$ single bonds use $8$ electrons. Each Bromine receives $3$ lone pairs ( $6$ non-bonding electrons) to achieve $8$ electrons. Carbon shares $4$ single bonds ( $8$ shared electrons) to complete its octet.
- Water ( $H_2O$ ):
- Total valence electron count: $2 \times 1 \text{ (from Hydrogen)} + 6 \text{ (from Oxygen)} = 8$ valence electrons.
- Structure: Central Oxygen atom single-bonded to $2$ Hydrogen atoms ( $H - O - H$ ).
- Electron allotment: $2$ single bonds use $4$ electrons. Oxygen receives $2$ lone pairs ( $4$ non-bonding electrons) to complete its octet ( $8$ total electrons). Each Hydrogen shares $2$ electrons (duet rule).

Water covalent Lewis structure diagram

Nitrogen Gas ( $N_2$ ):
- Total valence electron count: $2 \times 5 \text{ (from Nitrogen)} = 10$ valence electrons.
- Structure: A single bond uses $2$ electrons, leaving $8$ electrons. To satisfy the octet rule for both Nitrogen atoms, a triple bond is formed by sharing $6$ electrons ( $:N \equiv N:$ ).
- Electron allotment: Each Nitrogen atom possesses $1$ lone pair ( $2$ non-bonding electrons) and $1$ triple bond ( $6$ shared electrons) = $8$ valence electrons per atom.

Polyatomic Ions and Acid Naming

Definition and Fundamentals of Polyatomic Ions:
- A polyatomic ion is a covalently bonded compound (group of non-metal atoms) that carries an overall net electric charge.
- Representative example: Sulphate ion ( $SO_4^{2-}$ ).
- Suffix Conventions:
- MOST polyatomic ions end with the suffix -ate or -ite (e.g., sulphate $SO_4^{2-}$ , sulphite $SO_3^{2-}$ , nitrate $NO_3^-$ , phosphate $PO_4^{3-}$ ).
- EXCEPTIONS ending in -ide: Cyanide ( $CN^-$ ), Hydroxide ( $OH^-$ ).
- Charge Conventions:
- MOST polyatomic ions are negative ions (anions).
- EXCEPTIONS that are positive ions (cations):
  - Ammonium ( $NH_4^+$ )
  - Hydronium ( $H_3O^+$ )
- Rules for Subscripts and Parentheses (Brackets):
- The chemical subscripts inside a polyatomic ion CANNOT be modified or changed under any circumstances.
- To indicate more than one unit of a polyatomic group in a formula, place parenthetical brackets around the WHOLE polyatomic group before adding the outer subscript multiplier.

Lead(IV) cyanide chemical formula bracket representation

Nomenclature and Formula Writing Examples with Polyatomic Ions:
- Formula to Name Conversions:
- $Na_2SO_4$ : Sodium sulphate
- $Mg(OH)_2$ : Magnesium hydroxide
- $FeSO_4$ : Iron(II) sulphate
- $Ru(NO_3)_3$ : Ruthenium(III) nitrate
- $Au_2Cr_2O_7$ : Gold(I) dichromate
- Name to Formula Conversions:
- Sodium Carbonate: $Na_2CO_3$
- Lead(IV) Cyanide: $Pb(CN)_4$
- Ammonium Phosphate: $(NH_4)_3PO_4$
Nomenclature Rules for Acids:
- Definition of Aqueous State Symbol: The designation $(aq)$ stands for aqueous, meaning the compound is dissolved in water.
- Category 1: Binary Acids ( $H + \text{halogen}$ ):
- Formed from Hydrogen and a single halogen element.
- Naming pattern: hydro + root name of halogen + ic acid
- Example: $HCl(aq)$ = Hydrochloric acid
- Category 2: Oxyacids Derived from "-ate" Polyatomics ( $H + \text{Polyatomic ending in -ate with Oxygen}$ ):
- Formed from Hydrogen and an oxygen-containing polyatomic anion ending in -ate.
- Naming pattern: root name of element/polyatomic + ic acid (replace -ate with -ic acid)
- Example: $H_2SO_4(aq)$ = Sulphuric acid (derived from sulphate, $SO_4^{2-}$ )
- Category 3: Oxyacids Derived from "-ite" Polyatomics ( $H + \text{Polyatomic ending in -ite with Oxygen}$ ):
- Formed from Hydrogen and an oxygen-containing polyatomic anion ending in -ite.
- Naming pattern: root name of element/polyatomic + ous acid (replace -ite with -ous acid)
- Example: $H_2SO_3(aq)$ = Sulphurous acid (derived from sulphite, $SO_3^{2-}$ )