CHM 302 - INORGANIC CHEMISTRY II

CHM 302: Inorganic Chemistry II

Course Overview

The CHM 302 course covers a variety of important topics in inorganic chemistry including:

Chemistry of Hydrogen
The Noble Gases
Electronic structure and general properties of Group IA and Group IIA elements
Chemistry of elements such as Boron, Aluminium, Carbon, Silicon, Nitrogen, Phosphorus, Oxygen, Sulphur, and the Halogens
Transition elements
Methods of separation of metals
Introduction to coordination chemistry including ligand and crystal field theories
Introductory organo-metallic chemistry
Introduction to radiochemistry, focusing on radioactivity and its relation to the periodic table
The role of metals in living systems

Chemistry of Hydrogen

General Properties

Hydrogen is a colourless, odourless, and tasteless gas.
It is the most abundant element in the universe and the third most abundant on Earth, following oxygen and silicon.
Discovered by Henry Cavendish in 1766 through reactions of zinc, iron, and tin with various binary acids, exemplified by the reaction:
$2HCl(g) + Zn(s) ightarrow ZnCl_2(aq) + H_2(g)$
The name hydrogen is derived from Greek, meaning "water producer" (with "hydro" meaning water and "genes" meaning to make). Antoine Lavoisier named it in 1783 because it produces water upon ignition in air.
Hydrogen is found in several essential compounds for life, including:
  - Water (H₂O)
  - Fats
  - Petroleum
  - Table sugar (C₆H₁₂O₆)
  - Ammonia (NH₃)
  - Hydrogen peroxide (H₂O₂)

Placement in the Periodic Table

Hydrogen exhibits properties similar to both alkali metals and halogens, leading to a debated placement among chemists.

Similarities with Alkali Metals

Hydrogen shares various characteristics with alkali metals, including:
  - The presence of one electron in its valence shell
  - Ability to show a +1 oxidation state in compounds
  - Formation of compounds with non-metals and halogens
  - Loss of one electron to form unipositive ions, indicating an electropositive character
  - Good reducing agent due to its strong ability to form O–H bonds
Example reaction demonstrating hydrogen as a reducing agent:
$CuO + H_2 ightarrow Cu + H_2O$

Differences from Alkali Metals

Hydrogen has notable differences from alkali metals:
  - Ionization energy: A significantly higher ionization energy of 1312 kJ/mol for hydrogen compared to lithium (520 kJ/mol), sodium (495 kJ/mol), and potassium (418 kJ/mol).
  - Non-metallic character: Hydrogen is a non-metal, while alkali metals are metals.
  - Atomicity: Hydrogen exists as a diatomic molecule (H₂), whereas alkali metals are monoatomic.
  - Nature of oxides: The oxide of hydrogen is neutral (H₂O), while those of alkali metals are basic (e.g., Na₂O).
  - Nature of compounds: Hydrogen halides (HF, HCl, HBr, HI) are low-boiling covalent compounds, unlike the high-melting ionic solids formed by alkali metal halides.

Similarities with Halogens

Hydrogen shares similarities with halogens such as:
  - Capability to gain an electron for a stable electronic configuration
  - Ionization energy comparable to that of halogens (H = 1312 kJ/mol, Li = 520 kJ/mol, F = 1680 kJ/mol)
  - Formation of compounds in -1 oxidation states (Hydrides with metals and halides with halogens)
  - Exist as diatomic molecules (H₂)

Differences from Halogens

Differences include:
  - Electronegativity: Hydrogen is less electronegative than halogens; for example, it has a lesser tendency to form H⁻ ions.
  - Nature of oxides: Oxides of halogens are generally acidic, while that of hydrogen (H₂O) is neutral.
  - Unshared pairs of electrons: Hydrogen (H₂) lacks unshared pairs of electrons, while halogen molecules have such pairs.

Isotopes of Hydrogen

Hydrogen has three isotopes:
  - Protium (¹H): 0 neutrons, 99.985% abundance
  - Deuterium (²H or D): 1 neutron, 0.0156% abundance
  - Tritium (³H or T): 2 neutrons, radioactive with a half-life of 12.33 years

Properties of Hydrogen

Hydrogen is a non-metal placed above Group 1 due to its ns¹ electron configuration.
Unlike alkali metals, it forms cations (H⁺) reluctantly, with an ionization energy of 1312 kJ/mol versus lithium’s 520 kJ/mol.
Hydrogen forms hydride anions (H⁻) and can be viewed as placed above halogens in certain contexts.
It forms strong covalent bonds, acting as a good reducing agent for metals; an example is:
$CuO(s) + H_2(g) ightarrow Cu(s) + H_2O(g)$

Preparation of Dihydrogen (H₂)

Dihydrogen can be prepared through several methods:
 1. Water with active metals:
 - Sodium:
 $Na + H_2O ightarrow NaOH + H_2$
 - Potassium:
 $K + H_2O ightarrow KOH + H_2$

 2. Metal Hydrides with water:
 -
 $CaH_2 + H_2O ightarrow Ca(OH)_2 + H_2$

 3. Reaction of metals with acids:
 -
 $Zn + HCl ightarrow ZnCl_2 + H_2$
 -
 $Fe + HCl ightarrow FeCl_2 + H_2$

 4. Electrolysis of acidified water using platinum electrodes

 5. Steam with hydrocarbons or coke:
 - Example:
 $C_3H_8(g) + 3H_2O(g) ightarrow 3CO(g) + 7H_2(g)$

 6. Water-gas shift reaction:
 -
 $CO(g) + H_2O(g) ightarrow CO_2(g) + H_2(g)$

Reactions of Hydrogen

Hydrogen reacts with various elements, e.g.:
 - With halogens:

 $H_2(g) + X_2(g) ightarrow 2HX(g)$

 - With dioxygen:

 $2H_2(g) + O_2(g) ightarrow 2H_2O(g)$

 - With dinitrogen:

 $N_2(g) + 3H_2(g) ightarrow 2NH_3(g)$

 - With metal ions and metal oxides:

 $H_2(g) + Pd^{2+}(aq) ightarrow Pd(s) + 2H^+(aq)$

 - With organic compounds: - Example: Hydrogenation of vegetable oils using a nickel catalyst to produce fats:

 $CH_3CH_2CH=CH_2 + H_2 ightarrow CH_3CH_2CH_2CH_3$

Uses of Dihydrogen

Major uses include its application in:
  - Synthesis of ammonia for nitric acid and nitrogenous fertilizers
  - Manufacture of vanaspati fat through hydrogenation of vegetable oils
  - Production of bulk organic chemicals like methanol
  - Preparation of hydrogen chloride
  - Metallurgical processes to reduce heavy metal oxides
  - Atomic hydrogen and oxy-hydrogen torches for cutting and welding
  - Rocket fuel in space research
  - Fuel cells for generating electrical energy, offering pollution-free energy with greater energy release per unit mass than fossil fuels

### Hydrides

Dihydrogen forms binary compounds (hydrides) with almost all elements besides noble gases.

There are three categories of hydrides:
  - Ionic (saline) hydrides: Formed with most s-block elements, e.g., LiH, NaH, they are stoichiometric and conduct electricity in the molten state.
  - Covalent (molecular) hydrides: Formed with p-block elements, e.g., CH₄, NH₃, H₂O, can be classified as electron-deficient, electron-precise, or electron-rich based on Lewis structure.
  - Metallic (non-stoichiometric) hydrides:

The Noble Gases

General Information

Noble gases are a group of chemically inert elements, all odorless, colorless, monatomic gases with low chemical reactivity in standard conditions.
The six noble gases include:
  - Helium (He)
  - Neon (Ne)
  - Argon (Ar)
  - Krypton (Kr)
  - Xenon (Xe)
  - Radon (Rn)

Historical Background

The noble gases were discovered by Henry Cavendish in the late 18th century by removing oxygen and nitrogen from air.

Electron Configurations

The electron configurations for noble gases are as follows:
  - Helium: 1s²
  - Neon: [He] 2s² 2p⁶
  - Argon: [Ne] 3s² 3p⁶
  - Krypton: [Ar] 3d¹⁰ 4s² 4p⁶
  - Xenon: [Kr] 4d¹⁰ 5s² 5p⁶
  - Radon: [Xe] 4f¹⁴ 5d¹⁰ 6s² 6p⁶

Properties of Noble Gases

The following properties characterize noble gases:
  - Atomic mass, boiling point, and atomic radius increase down the group.
  - The first ionization energy decreases down the group.
  - Noble gases possess the largest ionization energies, illustrating their inertness.
  - As atomic size increases, so do atomic radius and interatomic forces, leading to increased boiling/melting points and enthalpy of vaporization.
  - There is a positive correlation between density and atomic mass down the group.