Electronic Configuration and Chemical

EElectronic Configuration and Chemical Combination

Electronic Configuration

Electronic configuration refers to the arrangement of electrons in an atom's orbitals. Electrons are distributed in shells (energy levels) around the nucleus, following specific rules:

1. Aufbau Principle – Electrons fill orbitals starting from the lowest energy level (1s) before filling higher levels.
2. Pauli Exclusion Principle – No two electrons in an atom can have the same set of four quantum numbers.
3. Hund's Rule – Electrons occupy degenerate orbitals singly before pairing.

Energy Levels and Subshells

• Principal Quantum Number (n): Defines the main energy level (K, L, M, N, etc.).
• Subshells (s, p, d, f): Orbitals within each energy level.
• Maximum Electrons per Subshell:
  • s = 2
  • p = 6
  • d = 10
  • f = 14

Electronic Configuration Notation

The general format follows 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶, etc.

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Electronic Configuration of the First 30 Elements

Atomic Number	Element	Electronic Configuration
1	Hydrogen (H)	1s¹
2	Helium (He)	1s²
3	Lithium (Li)	1s² 2s¹
4	Beryllium (Be)	1s² 2s²
5	Boron (B)	1s² 2s² 2p¹
6	Carbon (C)	1s² 2s² 2p²
7	Nitrogen (N)	1s² 2s² 2p³
8	Oxygen (O)	1s² 2s² 2p⁴
9	Fluorine (F)	1s² 2s² 2p⁵
10	Neon (Ne)	1s² 2s² 2p⁶
11	Sodium (Na)	1s² 2s² 2p⁶ 3s¹
12	Magnesium (Mg)	1s² 2s² 2p⁶ 3s²
13	Aluminium (Al)	1s² 2s² 2p⁶ 3s² 3p¹
14	Silicon (Si)	1s² 2s² 2p⁶ 3s² 3p²
15	Phosphorus (P)	1s² 2s² 2p⁶ 3s² 3p³
16	Sulfur (S)	1s² 2s² 2p⁶ 3s² 3p⁴
17	Chlorine (Cl)	1s² 2s² 2p⁶ 3s² 3p⁵
18	Argon (Ar)	1s² 2s² 2p⁶ 3s² 3p⁶
19	Potassium (K)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹
20	Calcium (Ca)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
21	Scandium (Sc)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹
22	Titanium (Ti)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d²
23	Vanadium (V)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
24	Chromium (Cr)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵
25	Manganese (Mn)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁵
26	Iron (Fe)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
27	Cobalt (Co)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷
28	Nickel (Ni)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸
29	Copper (Cu)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰
30	Zinc (Zn)	1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰

Chemical Combination

Chemical combination refers to how atoms bond to form molecules. The main types of chemical bonds are: 

1. Electrovalent (Ionic) Bond
2. Covalent bond 
3. Coordinate bonding 

Atoms combine to achieve stability by either losing, gaining, or sharing electrons to attain a stable electronic configuration (usually the noble gas configuration). This leads to the formation of chemical bonds, mainly electrovalent (ionic) bonding and covalent bonding.

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Electrovalent Bonding (Ionic Bonding)

Definition:

Electrovalent bonding (also called ionic bonding) is formed when one atom transfers one or more electrons to another atom. This occurs between metals and non-metals, leading to the formation of positive and negative ions that attract each other due to electrostatic forces.

Formation Process:

1. Metal atoms lose electrons to form positively charged ions (cations).
2. Non-metal atoms gain electrons to form negatively charged ions (anions).
3. The oppositely charged ions attract each other strongly, forming a stable ionic compound.

Examples of Electrovalent Bonding:

1. Sodium Chloride (NaCl) – Table Salt

• Sodium (Na) has one electron in its outer shell (1s² 2s² 2p⁶ 3s¹).
• Chlorine (Cl) has seven electrons in its outer shell (1s² 2s² 2p⁶ 3s² 3p⁵).
• Sodium donates its one outer electron to chlorine, forming Na⁺ and Cl⁻.
• The electrostatic attraction between Na⁺ and Cl⁻ forms sodium chloride (NaCl)
  
  
  .

Equation:

$$\text{Na} \to \text{Na}^+ + e^-$$ $$\text{Cl} + e^- \to \text{Cl}^-$$ $$\text{Na}^+ + \text{Cl}^- \to \text{NaCl}$$

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2. Magnesium Oxide (MgO)

• Magnesium (Mg) has two valence electrons (1s² 2s² 2p⁶ 3s²).
• Oxygen (O) has six valence electrons (1s² 2s² 2p⁴).
• Magnesium donates two electrons to oxygen, forming Mg²⁺ and O²⁻.
• The strong attraction between Mg²⁺ and O²⁻ results in the formation of MgO.

Equation:

$$\text{Mg} \to \text{Mg}^{2+} + 2e^-$$ $$\text{O} + 2e^- \to \text{O}^{2-}$$ $$\text{Mg}^{2+} + \text{O}^{2-} \to \text{MgO}$$

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Properties of Ionic Compounds:

✔ High melting and boiling points due to strong ionic bonds.
✔ Conduct electricity in molten or dissolved (aqueous) state but not in solid state.
✔ Soluble in polar solvents like water.

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Covalent Bonding

Definition:

Covalent bonding occurs when two or more non-metal atoms share electrons to achieve a stable configuration. Unlike ionic bonds, covalent bonds do not involve electron transfer but electron sharing.

Formation Process:

1. Each atom contributes one or more electrons for sharing.
2. The shared electrons form a covalent bond.
3. The resulting compound is a molecule rather than an ionic structure.
   
   
   

Types of Covalent Bonds:

1. Single Covalent Bond – One pair of shared electrons (H₂, Cl₂).
2. Double Covalent Bond – Two pairs of shared electrons (O₂, CO₂).
3. Triple Covalent Bond – Three pairs of shared electrons (N₂).

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Examples of Covalent Bonding:

1. Hydrogen Molecule (H₂) – Single Bond

• Each hydrogen atom (H) has one valence electron.
• Both hydrogen atoms share their electrons, forming a single bond (H—H).

Equation:

$$\text{H} \cdot + \cdot \text{H} \to \text{H}_2$$

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2. Oxygen Molecule (O₂) – Double Bond

• Each oxygen atom (O) has six valence electrons.
• They share two pairs of electrons, forming a double bond (O=O).

Equation:

$$\text{O} \cdot \cdot + \cdot \cdot \text{O} \to \text{O} = \text{O}$$

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3. Carbon Dioxide (CO₂) – Double Bond

• Carbon (C) has four valence electrons, and oxygen (O) has six.
• Carbon shares two electrons with each oxygen, forming two double bonds (O=C=O).

Equation:

$$\text{O} \cdot \cdot + \cdot \cdot \text{C} \cdot \cdot + \cdot \cdot \text{O} \to \text{O} = \text{C} = \text{O}$$

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Properties of Covalent Compounds:

✔ Low melting and boiling points due to weak intermolecular forces.
✔ Do not conduct electricity (except some like graphite).
✔ Insoluble in water but soluble in organic solvents.

Comparison of Electrovalent (Ionic) Bond and Covalent Bond

Property	Electrovalent (Ionic) Bond	Covalent Bond
Bond Formation	Electrons are transferred from a metal to a non-metal.	Electrons are shared between two non-metal atoms.
Type of Elements Involved	Occurs between a metal and a non-metal.	Occurs between two or more non-metals.
Example	NaCl (Sodium Chloride), MgO (Magnesium Oxide)	H₂ (Hydrogen), O₂ (Oxygen), CO₂ (Carbon Dioxide)
Type of Structure	Forms a crystal lattice structure.	Forms individual molecules.
Electrical Conductivity	Conducts electricity in molten or dissolved (aqueous) state.	Does not conduct electricity (except graphite).
Solubility	Soluble in water (polar solvent).	Soluble in organic solvents, but mostly insoluble in water.
Melting & Boiling Points	High due to strong electrostatic forces.	Low due to weak intermolecular forces.
Bond Strength	Strong due to ionic attraction.	Relatively weaker due to shared electrons.

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3. Coordinate Covalent (Dative) Bonding

• A special type of covalent bond where both shared electrons come from the same atom.
• Example: Ammonium ion (NH₄⁺) – The nitrogen atom donates a lone pair to bond with H⁺.

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Coordinate Bonding and Van der Waals Forces

Coordinate (Dative) Bonding

A coordinate bond (also called a dative covalent bond) is a type of covalent bond where both electrons in the bond are donated by one atom. This happens when one atom has a lone pair of electrons and donates it to another atom that has an empty orbital.

Examples of Coordinate Bonding:

• Ammonium Ion (NH₄⁺): NH₃ donates a lone pair to H⁺.
• Hydronium Ion (H₃O⁺): H₂O donates a lone pair to H⁺.
• Carbon Monoxide (CO): Oxygen donates a lone pair to Carbon.

Van der Waals Forces

Van der Waals forces are weak intermolecular forces that exist between molecules. These forces do not involve the transfer or sharing of electrons but are responsible for interactions between neutral molecules.

Types of Van der Waals Forces:

• London Dispersion Forces: Temporary forces due to the movement of electrons.
• Dipole-Dipole Interaction: Attraction between polar molecules.
• Hydrogen Bonding: A strong dipole-dipole interaction, seen in water molecules.

JAMB Chemistry Quiz - Bonding