Ionic Equilibria Notes

SYJC Science Chemistry Part 01: Ionic Equilibria

Introduction to Ionic Equilibria

• In chemical equilibrium, reactions involve only molecules.

• In ionic equilibrium, reversible reactions involve the formation of ions in water.

• When a polar covalent compound (solute) is dissolved in a solvent (water), it reacts with water to form ions because water has the ability to dissociate the solute into ions.

• $Solute + Solvent \rightleftharpoons H^+ + H<em>2O \rightleftharpoons H</em>3O^+$

Electrical Conductors

• Substances that allow electric current to pass through them are known as conductors or electrical conductors.

• Conductors are divided into two types:

  • Metallic or electronic conductors: Conduct electricity without undergoing any chemical change (e.g., silver, gold, copper, steel).

  • Electrolytic conductors or electrolytes: Undergo decomposition (a chemical change) when an electric current is passed through them.

• Examples:

  • Ethanol: No conductivity

  • KCl: High conductivity

  • Acetic acid solution: Low conductivity

Types of Electrolytes

• Electrolytes: Substances that dissociate or ionize in their aqueous solutions (or in any polar solvents), forming positively charged ions (cations) and negatively charged ions (anions) (e.g., NaCl, HCl).

• Electrolytes are divided into two types based on their strengths:

  • Strong electrolytes:

    • Almost completely ionize into ions in their aqueous solution.

    • Degree of ionization is approximately one ($\alpha \approx 1$).

    • Examples: HCl, H2SO4, NaCl, HNO3, KOH, NaOH, AgNO3, CuSO4 (strong acids, strong bases, and salts).

    • $AX(aq) \rightarrow A^+ + X^-$

  • Weak electrolytes:

    • Ionize to a small extent in their aqueous solution.

    • Degree of ionization is much less than one ($\alpha <<< 1$).

    • Examples: H2O, CH3COOH, NH4OH, HCN, Liq. SO2, HCOOH (weak acids and weak bases).

    • $MZ(aq) \rightleftharpoons M^+ + Z^-$

• Illustrations:

  • Addition of strong electrolyte AX(aq) into water: Ions furnished from strong electrolytes conduct electricity (maximum glow due to complete ionization).

  • Addition of weak electrolyte MZ(aq) into water: Ions furnished from weak electrolytes conduct electricity (less glow due to partial ionization).

Distinguishing Between Acids and Bases

• Acids:

  • Sour to taste.

  • Turn blue litmus to red.

  • Neutralize bases.

  • Liberate H2 gas on reaction with metals.

  • Aqueous solutions are good conductors.

• Bases:

  • Bitter to taste.

  • Turn red litmus to blue.

  • Neutralize acids.

  • Soapy to touch.

  • Aqueous solutions are good conductors.

Concepts of Acids and Bases

• Different concepts have been proposed to characterize acids and bases:

  • Arrhenius Concept

  • Bronsted-Lowry Concept

  • Lewis Concept

Arrhenius Concept

• Acids: Substances that give $H^+$ ions when dissolved in water (e.g., HCl, H2SO4, HNO3).

• Bases: Substances that ionize in water to furnish $OH^-$ ions (e.g., NaOH, KOH).

• Salts: Substances that neither give $H^+$ nor $OH^-$ ions when ionized in water.

• Neutralization: Reaction of an acid and a base to form salt and water.

  • $HCl + NaOH \rightarrow NaCl + H_2O$ (Acid + Base → Salt + Water)

  • $H^+ + OH^- \rightarrow H_2O$

• Types of acids:

  • Monobasic acids: Give a single $H^+$ per molecule (e.g., $HNO<em>3$, $HClO</em>4$, $HCl$).

  • Dibasic acids: Give two $H^+$ per molecule (e.g., $H<em>2SO</em>4$, $H<em>2PO</em>3$).

  • Tribasic acids: Give three $H^+$ per molecule (e.g., $H<em>3PO</em>4$).

• Types of bases:

  • Monoacidic bases: Give a single $OH^-$ per molecule (e.g. $CsOH$, $RbOH$, $NH_4OH$).

  • Diacidic bases: Give two $OH^-$ per molecule (e.g., $Ba(OH)<em>2$, $Ca(OH)</em>2$).

  • Triacidic bases: Give three $OH^-$ per molecule (e.g., $Fe(OH)_3$).

• Acids produce $H^+$ ions, which are the only positive ions in the solution.

  • $HCl \rightarrow H^+ + Cl^-$

• Bases produce $OH^-$ ions, which are the only negative ions in the solution.

  • $NaOH \rightarrow Na^+ + OH^-$

• Neutralization is also a double replacement reaction:

  • Acid + Base → Salt + water

  • $H_x + Metal\ OH \rightarrow Metal\ with\ nonmetal + water$

  • $HI + KOH \rightarrow KI + H_2O$

• Pure water is when pH = 7, Net ionic result: $HOH \rightleftharpoons H_2O$ (Neutral)

Important Points Regarding Arrhenius Concepts

• Strong acids and bases ionize almost completely in solutions.

  • Examples of strong acids: HCl, HNO3, H2SO4, HClO4.

  • Examples of strong bases: NaOH, KOH, Ba(OH)2.

• Weak acids and bases do not ionize almost completely in solutions.

  • Examples of weak acids: HCOOH, CH3COOH.

  • Examples of weak bases: NH4OH.

• Not every hydrogen compound is an acid (e.g., CH4 is not an acid).

• Compounds with -OH groups aren't necessarily bases (e.g., CH3OH, C2H5OH are not bases).

Limitations of the Arrhenius Concept

• Does not explain acidic and basic character of substances in non-aqueous solvents.

• Cannot explain neutralization reactions giving salt in the absence of a solvent.

• Cannot explain the acidic character of certain non-protonic salts like AlCl3, BF3 in aqueous solution.

• Cannot explain the existence of $H^+$ in water ($H_3O^+$, hydronium).

• Does not explain the basic nature in $NH<em>3$, $PH</em>3$, $NH_2$.

• Does not explain the acidic character of non-metal oxides and basic character of metal oxides.

Bronsted-Lowry Concept

• An acid is a substance (molecule or ion) that can donate a proton ($H^+$) to some other substance.

• A base is a substance that can accept a proton from an acid.

• Simply, an acid is a proton-donor and a base is a proton acceptor.

• Example 1: $HCl + H<em>2O \rightleftharpoons H</em>3O^+ + Cl^-$

  • HCl acts as an acid because it donates a proton to the water molecule.

  • Water behaves as a base by accepting a proton from the acid.

• Example 2: $NH<em>3 + H</em>2O \rightleftharpoons NH_4^+ + OH^-$

  • Water acts as an acid as it donates a proton to $NH_3$ molecules.

  • $NH_3$ molecule behaves as a base as it accepts a proton.

• $HCl(aq) + NH<em>3(aq) \rightleftharpoons NH</em>4^+(aq) + Cl^-(aq)$

  • acid + base = addition of a proton

Conjugate Acid & Base Pairs

• When an acid loses a proton, the remaining part has a tendency to regain a proton, thus behaving as a base.

  • $Acid \rightleftharpoons H^+ + base$

• Acid and base which differ by a proton are known to form a conjugate pair.

• General Reaction

  • $HA(aq) + H<em>2O(l) \rightleftharpoons A^-(aq) + H</em>3O^+(aq)$

  • Acid + Base = Conjugate base + Conjugate acid

• Some more examples:

  • $Acid<em>1 + Base</em>2 \rightleftharpoons Acid<em>2 + Base</em>1$

  • $HCl + H<em>2O \rightleftharpoons H</em>3O^+ + Cl^-$

    • remove $H^+$ becomes conjugate base

    • add $H^+$ becomes conjugate acid

  • $H<em>2O + NH</em>3 \rightleftharpoons NH_4^+ + OH^-$

  • $HCN + H<em>2O \rightleftharpoons H</em>3O^+ + CN^-$

  • $HF + CH<em>3COOH \rightleftharpoons CH</em>3COOH_2^+ + F^-$

  • $CH<em>3COOH + NH</em>3 \rightleftharpoons NH<em>4^+ + CH</em>3COO^-$

  • $NH<em>3 + NH</em>3 \rightleftharpoons NH<em>4^+ + NH</em>2^-$

  • $H<em>2O + H</em>2O \rightleftharpoons H_3O^+ + OH^-$

  • $HCO<em>3^- + H</em>2O \rightleftharpoons H<em>3O^+ + CO</em>3^{2-}$

  • $NH<em>4^+ + H</em>2O \rightleftharpoons H<em>3O^+ + NH</em>3$

  • $HNO<em>2(aq) + H</em>2O(l) \rightleftharpoons NO<em>2^-(aq) + H</em>3O^+(aq)$

  • $HF(aq) + H<em>2O(l) \rightleftharpoons F^-(aq) + H</em>3O^+(aq)$

Acid and Base Strength

• Strong acids have weak conjugate bases, while weak acids have strong conjugate bases.

• The strength of an acid depends upon its tendency to lose its proton, and the strength of a base depends upon its tendency to gain the proton.

• Example: Dissociation of HCl

  • $HCl + H<em>2O \rightarrow H</em>3O^+ + Cl^-$

  • Strong acid + Weak conjugate base

• Relationship of acid and based + $H<em>2O$: $CH</em>3COOH(aq) \rightleftharpoons H^+(aq) + CH_3COO^-(aq)$

• Leveling effect: In acid-base strength series, all acids above $H<em>3O^+$ in aqueous solution fall to the strength of ionized $H</em>3O^+$. Similarly, the basic strength of bases below $OH^-$ fall to the strength of $OH^-$ in aqueous solution.

• Acid strength increases: Strong → Weak

• Base strength increases: Weak → Strong

• The relative strength of acids:

  • HCl
    *$H<em>2SO</em>4$

• HNO3
  *$H<em>3O^+(aq)$ *$C</em>2H<em>3O</em>2$

• H2CO3
  *H2S

• H2PO4-

• NH4+
  $HCO_3^-$

• HPO42-
  *H2O
  *OH-
  *H2
  *CH4
  Acid Strength
  Base Strength
  Negligible
  Strong
  Weak
  Negligible
  Strong
  Weak
  Base

• The strength of an acid depends upon the solvent used (e.g., HClO4 > H2SO4 > HCl > HNO3 in acetic acid).

Types of Solvents

• On the basis of proton interaction, solvents can be classified into four types:

  • Protophilic solvents: Solvents with a greater tendency to accept protons (e.g., water, alcohol, liquid ammonia).

  • Protogenic solvents: Solvents with a tendency to produce protons (e.g., water, liquid hydrogen chloride, glacial acetic acid).

  • Amphiprotic solvents: Solvents that act both as protophilic or protogenic (e.g., water, ammonia, ethyl alcohol).

  • Aprotic solvents: Solvents that neither donate nor accept protons (e.g., benzene, carbon tetrachloride ($CCl<em>4$), carbon disulfide ($CS</em>2$)).

Limitations of Bronsted-Lowry Concept

• Failed to explain the acidic nature of BF3.

• Failed to explain the basic nature of metal oxides like MgO, and CaO.

Lewis Concept

• A base is a substance that can donate a pair of electrons to form a coordinate bond, whereas an acid is a substance that can accept a pair of electrons.

• The acid is also known as an electron pair acceptor or electrophile, while the base is an electron pair donor or nucleophile.
  $Acid (electron\ pair\ acceptor) + Base (electron\ pair\ donor) \rightarrow New\ bond (coordinate\ covalent)$

• Lewis acids:

  • Molecules in which the central atom has an incomplete octet (e.g., BF3, BCl3, AlCl3, BeCl2).

  • Simple cations: All cations are expected to act as Lewis acids since they are deficient in electrons ($Na^+$ , $K^+$, $Ca^{2+}$, etc have a very little tendency to accept electrons while the cations like $H^+$, $Ag^+$, $Fe^{2+}$, $CO^{2+}$, $Mn^{2+}$ etc., have greater tendency to accept electrons).

  • Molecules in which the central atom has empty d-orbitals (e.g., SiX4, GeX4, TiCl4, SnX4, PX3, PF5, SF4, SeF4, TeCl, etc.).

• Lewis bases:

  • Neutral species having at least one lone pair of electrons (e.g., ammonia, amines, alcohols, RNH2, R2NH, R3N, H2O, ROH, ROR, etc.).

  • Negatively charged species or anions (e.g., chloride cyanide, hydroxide ions, etc., CN−, Cl−, Br−, I−, OH−, etc.).

• All Bronsted bases are also Lewis bases, but all Bronsted acids are not Lewis acids.

Limitation of Lewis Concept

• Since the strength of Lewis acids and bases depends on the type of reaction, it is not possible to arrange them in any order of their relative strength.