Electrovalency in Chemistry

Electrovalency

• Positive Electrovalency: Represents the number of electrons lost by an atom.

• Sodium (Na) loses 1 electron to form Na+, with a configuration change from (2,8,1) to (2,8).

• Magnesium (Mg) loses 2 electrons to become Mg2+, changing from (2,8,2) to (2,8).

• Calcium (Ca) also loses 2 electrons, forming Ca2+, altering from (2,8,8,2) to (2,8,8).

• Negative Electrovalency: Indicates the number of electrons gained to complete the outer shell.

• Chlorine (Cl) gains 1 electron to form Cl-, transitioning from (2,8,7) to (2,8,8).

• Oxygen (O) gains 2 electrons to form O2-, with a shift from (2,6) to (2,8).

• Nitrogen (N) gains 3 electrons to form N3-, changing from (2,5) to (2,8).

Chemical Bonds

• A chemical bond is the attractive force between atoms in a molecule.

• Types of bonds include:

  • Ionic bond (electrovalent bond)
  • Covalent bond
  • Coordinate (dative) bond
  • Hydrogen bond
  • Metallic bond

• Ionic Bond: Formed by the transfer of electrons. Examples are NaCl, MgO, and CaO.

• Characteristics of Ionic Compounds:

  • Soluble in water
  • Conduct electricity
  • High melting and boiling points
  • Non-directional

Covalency

• Describes how many electrons are shared to fill the outer shell.
• Single Covalency: Sharing one pair of electrons, e.g., H2, Cl2.
• Double Covalency: Sharing two pairs, e.g., O2.
• Triple Covalency: Sharing three pairs, e.g., N2.

Structure of an Atom

• An atom is the smallest unit of an element, consisting of protons, neutrons, and electrons.
• Atoms are electrically neutral due to equal numbers of protons and electrons.
• Fundamental particles:
  • Proton (P): +1 charge, located in the nucleus, discovered by Rutherford.
  • Neutron (n): Neutral, also in the nucleus, discovered by Chadwick.
  • Electron (e-): -1 charge, located outside the nucleus, discovered by Thomson.
• Atomic structure influences physical and chemical properties.

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in the nucleus; indicates element's position in the periodic table.
• Mass Number (A): Total of protons and neutrons; does not determine the position in the periodic table.
• Relation: Number of Neutrons (n) = Mass number (A) - Atomic number (Z).

Electronic Configuration

• Distribution of electrons in orbits around the nucleus.
• Examples:
  • Sodium (Na): 11 electrons represented as (2, 8, 1).
  • Magnesium (Mg): 12 electrons as (2, 8, 2).

Bond Types

• Coordinate Bond (Dative Bond): Shared electron pair originating from one atom, indicated by an arrow (→). Examples include NH4+ and H2SO4.
• Hydrogen Bond: A weak bond formed between hydrogen and electronegative atoms (e.g., F, O, N).
• Intramolecular Hydrogen Bonds involve attractions within a single molecule, while Intermolecular bonds occur between different molecules.

Properties of Hydrogen Bonding

• Weaker than ionic and covalent bonds.
• Compounds with hydrogen bonding often have higher melting and boiling points.
• Essential for the stability of certain crystalline forms.

Types of Solids

• Crystalline Solids: Highly ordered structures with a specific geometric arrangement. Defined by unit cells.
• Amorphous Solids: Random particle arrangement, lack sharp melting points; examples include glass and plastic.

Crystal Lattice and Unit Cell

• Crystal Lattice: 3D structure formed by the arrangement of atoms or molecules.
• Unit Cell: Smallest repeating unit in a crystal; contributes to the properties of the entire crystal.
• Coordinate Number: The total number of atoms or ions surrounding a given atom in a crystal lattice.

Types of Crystal Structures

• Simple Cubic (SC): Atoms at corners of a cube. Coordination number = 6.
• Body Centered Cubic (BCC): Atoms at the corners and one in the center. Coordination number = 8.
• Face Centered Cubic (FCC): Atoms at corners and centers of each face. Coordination number = 12.
• Hexagonal Close Packed (HCP): Compact arrangement with a different geometric configuration.