Chemical Reactions and Properties of Solutions | General Chemistry 2

Electrolyte and conductivity:

An electrolyte is a substance that dissolves in water to form solutions that conduct electricity by producing ions. Electrolytes can be classified as strong or weak based on their ability to dissociate into ions in solution.

Electrolytic solutions conduct electricity because mobile ions move and conduct an electric current.
 

Types of electrolytes:

• Strong electrolytes: Substances that completely dissociate into ions in solution. Solutions of strong electrolytes conduct electricity very well due to the high concentration of ions.
• Weak electrolytes: Substances that partially dissociate into ions in solution. Solutions of weak electrolytes conduct electricity poorly because only a small fraction of the molecules dissociate into ions.
• Non-electrolytes: Substances that dissolve to give non-conducting solutions as they do not dissociate into ions.
   

• Strong electrolytes: Strong acids (e.g., HCl), strong bases (e.g., NaOH), and salts (e.g., NaCl).
  Completely ionize in solution: NaCl (s) → Na⁺ (aq) + Cl^- (aq)
• Weak electrolytes: Weak acids (e.g., acetic acid) and weak bases (e.g., ammonia).
  Partially ionize in solution: CH₃COOH (aq) $\rightleftharpoons$ CH₃COO^- (aq) + H⁺ (aq)

Ionic strength I:

Ionic strength is a measure of the electrical intensity of a solution containing ions.
 

Solution:

A solution is a homogeneous mixture consisting of a solvent (usually present in greater quantity) and one or more dissolved species called solutes. A saturated solution is a homogeneous mixture that contains as much solute as possible; additional solute will remain undissolved.

An aqueous solution of NaCl is a homogeneous mixture composed of NaCl (the solute) dissolved in water (the solvent).

Solubility:

The solubility of a solute is the maximum amount of solute that can be dissolved in a specified amount of solvent at a particular temperature. A solid is generally more soluble at higher temperatures and in a solvent with similar types of intermolecular forces.

The solubility of NaCl in water at 25°C is 360 g per kg of water, meaning 360 g of NaCl can be dissolved in 1 kg (1 L) of water. If more NaCl is added, it will remain undissolved.

Factors affecting solubility:

• Nature of solute and solvent: Like dissolves like ⇒ Polar solutes dissolve in polar solvents (e.g., salt in water), and nonpolar solutes dissolve in nonpolar solvents (e.g., oil in hexane).
• Temperature: Generally, the solubility of solids in liquids increases with an increase in temperature, while the solubility of gases in liquids decreases with an increase in temperature.
• Pressure: The solubility of gases in liquids increases with an increase in pressure (Henry's Law).

Solubility rules for ionic compounds:

Soluble compounds:

• Compounds containing alkali metal ions (Li⁺, Na⁺, K⁺, etc.) and ammonium ion (NH₄⁺​).
• Compounds containing nitrate (NO₃^-​), acetate (CH₃​COO^-), and most sulfate (SO₄^2-​) ions.

Insoluble compounds:

• Compounds containing carbonate (CO₃^2-​), phosphate (PO₄^3-​), and hydroxide (OH^-) ions, except when paired with alkali metals or ammonium.
• Most other ionic compounds.

Mole fraction x:

The mole fraction is the number of moles of a component divided by the total number of moles in a mixture.
 

Precipitation reaction:

A precipitation reaction is a chemical reaction in which a precipitate is formed. A solid forms if ions of an insoluble salt are present.

• Ni²⁺ (aq) + S^2- (aq) → NiS (s)
  Aqueous solution of Ni²⁺ and S^2- results in the formation of a precipitate: NiS.
• AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq)
  Homogeneous mixture of AgNO₃ and NaCl results in the formation of a precipitate: AgCl.

Types of equations:

• Molecular equation: Compounds are represented as if none of the reactants or products has dissociated.
• Ionic equation: Strong electrolytes are represented as ions.
• Net ionic equation: An ionic equation from which spectator ions have been eliminated; only ions involved in the reaction are shown.

Molecular equation: AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq)

Ionic equation: Ag⁺ (aq) + NO₃^- (aq) + Na⁺ (aq) + Cl^- (aq) → AgCl (s) + Na⁺ (aq) + NO₃^- (aq)

Net ionic equation: Ag⁺ (aq) + Cl^- (aq) → AgCl (s)

How to determine the net ionic equation for a precipitation reaction:

1. Write and balance the molecular equation.
2. Write the ionic equation by representing strong electrolytes as their constituent ions.
3. Identify spectator ions that appear on both sides of the equation.
4. Write the net ionic equation by removing the spectator ions.

Arrhenius theory:

• Arrhenius acid: A substance which has hydrogen in its formula and dissociates in water to give H⁺ ions, increasing H⁺ concentration when added to water.
• Arrhenius base: A substance which has OH in its formula and dissociates in water to give OH^- ions, increasing OH^- concentration when added to water.

HCl is an Arrhenius acid: HCl (aq) → H⁺ (aq) + Cl^- (aq)

NaOH is an Arrhenius base: NaOH (aq) → Na⁺ (aq) + HO^- (aq)

Brönsted-Lowry theory:

• Brönsted acid: A substance that can donate a proton H⁺ (proton donor).
• Brönsted base: A substance that can accept a proton H⁺ (proton acceptor).

Conjugate acid-base pair: An acid-base pair that differs only in the presence or absence of a proton.
 

NH₃ (aq) + H₂SO₄ (aq) → NH₄⁺ (aq) + HSO₄^- (aq)

NH₃ is a base, H₂SO₄ is an acid.

NH₄⁺ / NH₃ and H₂SO₄ / HSO₄^- are 2 conjugate acid-base pairs.

Lewis theory:

• Lewis acid: A species that can accept a pair of electrons (electron-pair acceptor). Lewis acid must be electron deficient or have a vacant orbital (e.g., BF₃, AlCl₃, SO₂).
• Lewis base: A species that can donate a pair of electrons (electron-pair donor). Lewis base must have a lone pair of electrons (e.g., NH₃, H₂O).

When a Lewis base donates a pair of electrons to a Lewis acid, a covalent bond is formed between the molecules and the product is called an adduct.
 

NH₃ + BF₃ → NH₃BF₃

NH₃BF₃ is an adduct.

Strength of acids and bases:

• Strong acids and bases: Completely dissociate in aqueous solution.
• Weak acids and bases: Partially dissociate in aqueous solution.
   

• Strong acids: HCl, HNO₃​, H₂SO₄
  Strong bases: NaOH, KOH
• Weak acids: Acetic acid (CH₃​COOH), hydrofluoric acid (HF)
  Weak bases: Ammonia (NH₃​), methylamine (CH₃NH₂)

Neutralization reactions:

A neutralization reaction is a reaction between an acid and a base that produces water and a salt. Neutralization reactions are exothermic, releasing heat as the reaction proceeds. The general equation is:

Acid + Base → Salt + Water
 

Reaction between hydrochloric acid and sodium hydroxide:

HCl (aq) + NaOH (aq) → NaCl (aq) + H₂​O (l)

Reaction with carbonates:

A reaction between an acid and a carbonate that produces water, carbon dioxide, and a salt.
 

2 HCl (aq) + CaCO₃ ​(s) → CaCl₂​ (aq) + H₂​O (l) + CO₂ ​(g)

Oxidation number:

The oxidation number is a numerical value assigned to an atom in a compound or ion, representing its loss or gain of electrons relative to its elemental state. It is also known as the hypothetical charge that an atom would have if all bonds were completely ionic.

The oxidation number helps determine how electrons are distributed in molecules and ions, particularly in redox reactions.
 

Rules for assigning oxidation numbers:

• The oxidation number of an atom in its elemental form is 0.
• The oxidation number of a monatomic ion is equal to its charge.
• Alkali Metals (Group 1): The oxidation number is +1.
• Alkaline Earth Metals (Group 2): The oxidation number is +2.
• Halogens (Group 17): The oxidation number is -1.
• Hydrogen: +1 when bonded to nonmetals; -1 when bonded to metals.
• Oxygen: The oxidation number is usually -2 in most compounds. Exception: In peroxides (e.g., H₂O₂​), oxygen has an oxidation number of -1.
• The sum of the oxidation numbers in a neutral compound is zero, and in a polyatomic ion, it equals the ion's charge.
   

Assigning oxidation numbers in H₂O:

• H: +1; O: -2
• Total: 2x (+1) + 1 x (−2) = 0

 Assigning oxidation numbers in KMnO₄​:

• K: +1; O: -2 (each); Mn: +7
• Total: 1 x (+1) + 1 x (+7) + 4 x (−2) = 0

Main-group elements oxidation numbers:

The oxidation numbers of main-group elements are given by the formula: Oxidation state = number of valence electrons in the free atom - number of valence electrons assigned to the atom in the molecule.
 

Oxidation number of S in SOCl₂:
 

 

• O and Cl are more electronegative than S, so the electrons of the bonds are assigned to O and Cl.
• Number of valence electrons in the free S = 6
• Number of valence electrons assigned to S in SOCl₂ = 2
• Oxidation number of S in SOCl₂ = 6 – 2 = +4

Oxidation-reduction reaction (redox reaction):

A redox reaction is a chemical reaction in which electrons are transferred from one reactant to another. The oxidation state of an atom, molecule, or ion changes during a redox reaction.

• Oxidation: The process in which a particle becomes more positively charged through the loss of electrons, resulting in an increase in its oxidation state.
• Reduction: The process in which a particle becomes less positively charged through the gain of electrons, resulting in a decrease in its oxidation state.

• Oxidation: Zn → Zn²⁺ + 2 e⁻
• Reduction: Cu²⁺ + 2 e⁻ → Cu

Oxidizing vs. reducing agent:

• Oxidizing agent: A species that can accept electrons; it is reduced in a redox reaction.
• Reducing agent: A species that can donate electrons; it is oxidized in a redox reaction.

• 2 Fe (s) + 3 Cl₂ (aq) → 2 Fe³⁺ (aq) + 6 Cl^- (aq) is an oxidation-reduction reaction.
• Fe becomes Fe³⁺ during this reaction ⇒ it donates electrons ⇒ it is the reducing agent.
• Cl₂ becomes Cl^- during this reaction ⇒ it gains electrons ⇒ it is the oxidizing agent

Half redox reactions:

Electron-transfer reactions can be separated into 2 half reactions:

• Oxidation half-reaction: Represents the loss of electrons, where electrons appear on the right side.
• Reduction half-reaction: Represents the gain of electrons, where electrons appear on the left side.

2 Fe (s) + 3 Cl₂ (aq) → 2 Fe³⁺ (aq) + 6 Cl^- (aq) can be separated into 2 half reactions:

• Oxidation half-reaction: Fe (s) → Fe³⁺ (aq) + 3 e^-
• Reduction half-reaction: Cl₂ (aq) + 2 e^- → 2 Cl^- (aq)

Steps to balance redox reactions:

1. Divide the reaction into half-reactions.
2. Balance atoms other than O and H.
3. Balance O atoms by adding H₂O molecules.
4. Balance H atoms by adding H⁺ ions (in acidic solution) or OH^- (in basic solution).
5. Balance charge by adding electrons e^- to one side of each half-reaction.
6. Multiply each half-reaction by an integer to ensure the number of electrons lost equals the number of electrons gained.
7. Combine the half-reactions, canceling out the electrons.