Chemical Bonds and Valence Electrons

Chemical Bonds

• Chemical bonds significantly influence the physical and chemical properties of compounds.
• These bonds involve valence electrons, which are electrons in the outermost energy level of an atom.

Valence Electrons

• Valence electrons are involved in bond formation.
• The number of valence electrons can be determined by an element's position in the periodic table.
  • Hydrogen (H) has 1 valence electron.
  • Oxygen (O) has 6 valence electrons.
  • Aluminum (Al) has 3 valence electrons.
  • Nitrogen (N) has 5 valence electrons.

Lewis Structure

• A Lewis structure, or electron dot diagram, displays an atom's valence electrons.
• The structure shows the atom's symbol surrounded by dots representing the number of valence electrons.

Electronic Theory of Valence

• Atoms interact by gaining, losing, or sharing electrons to achieve a stable noble gas configuration.
• This theory was developed by G.N. Lewis and W. Kossel in 1916.

Octet Rule

• Atoms interact to achieve a stable outer shell of 8 electrons.
• Atoms with less than 4 valence electrons tend to lose electrons.
• Atoms with more than 4 valence electrons tend to gain electrons.
• There are exceptions to this rule.

Oxidation Number

• The oxidation number represents the charge an atom would have if it gained or lost electrons.
• It is used to determine which atoms will interact or bond.
  • Magnesium (Mg) has an oxidation number of +2
  • Oxygen (O) has an oxidation number of -2
  • Hydrogen (H) has an oxidation number of +1 or -1

Chemical Bonding

• Chemical bonding is the force holding atoms together in a molecule.
• There are three main types of bonding:
  • Ionic bonding
  • Metallic bonding
  • Covalent bonding

Ionic Bonding

• Involves the transfer of valence electrons between atoms.
• Creates ions with opposite charges, resulting in a stable molecule.
• Ionic bonds form between metal cations and non-metal anions.
  • Examples include Sodium Chloride (NaCl) , Magnesium Chloride (MgCl2), Calcium Oxide (CaO), and Aluminium Oxide (Al2O3).

Conditions for Ionic Bond Formation

• Atoms that lose electrons should have 1, 2, or 3 valence electrons.
• Atoms that gain electrons should have 5, 6, or 7 valence electrons.
• The formation of an ionic compound results in a net release of energy.
• The electronegativity difference between the two atoms should be greater than 2.

Factors Governing Ionic Bond Formation

• Ionization energy: Energy needed to remove an electron from an atom.
• Electron affinity: Energy released when an electron is added to an atom.
• Lattice energy: Energy released when ions combine to form a solid ionic compound.

Covalent Bonding

• Atoms share one or more pairs of electrons to form a covalent bond, achieving a stable molecule.

Definition of Covalent Bond

• Covalent bonding involves the sharing of one or more electron pairs between two atoms to form a stable molecule.

Conditions for Covalent Bond Formation

• Atoms should have 5, 6, or 7 valence electrons.
• Atoms should have equal or near equal electronegativity to facilitate electron sharing.
• Sharing of electrons should be equal.

Types of Covalent Bonds

• Single bond: Sharing of two electrons (one pair).
• Double bond: Sharing of four electrons (two pairs).
• Triple bond: Sharing of six electrons (three pairs).

Sigma (σ) and Pi (π) Bonds

• Sigma (σ) bonds: Formed by end-to-end overlapping of atomic orbitals along the internuclear axis
• Pi (π) bonds: Formed by side-to-side overlapping of atomic orbitals perpendicular to the internuclear axis.

Examples of Covalent Compounds

• Hydrogen (H₂): Sharing of one electron pair between two hydrogen atoms.
• Methane (CH₄): Sharing of four electron pairs between carbon and four hydrogen atoms.
• Water (H₂O): Sharing of two electron pairs between oxygen and two hydrogen atoms.

Unequal Sharing of Electrons (Polar Covalent Bonds)

• When atoms with different electronegativities share electrons, the sharing is unequal.
• This creates a partial positive charge on one atom and a partial negative charge on the other atom.
• Examples include water (H₂O), ammonia (NH₃), and hydrogen fluoride (HF).

Nonpolar and Polar Molecules

• Nonpolar molecules: Symmetrical shapes, no net dipole moment.
• Polar molecules: Asymmetrical shapes, net dipole moment.

Dipole-Dipole Interactions

• Weak intermolecular forces between polar molecules.
• Responsible for physical properties like boiling and melting points.

Coordinate Covalent Bonds

• Both electrons in the shared electron pair come from one atom.

Hydrogen Bonding

• Strong intermolecular force between hydrogen atom bonded to a highly electronegative atom (O, N, or F) and another electronegative atom.
• Responsible for unique properties of certain compounds and molecules.

Characteristics of Covalent Compounds

• Low melting and boiling points due to weaker intermolecular forces.
• Solubility in organic solvents due to similar intermolecular forces.
• Generally insoluble in water due to lack of hydrogen bonding.
• Non-conducting electricity due to the absence of free ions.

Electron-Deficient Compounds

• Some compounds have fewer than 8 electrons around the central atom.
  • Beryllium Chloride (BeCl₂)
  • Boron Trifluoride (BF₃)

تتناول هذه الاختبار أهمية الروابط الكيميائية وسلوك الإلكترونات التكافؤية. يتم استكشاف كيفية تأثير هذه الروابط على خصائص المركبات، بالإضافة إلى فهم البنية لويس و قاعدة الثمانية. هذا الاختبار مناسب للطلاب الذين يدرسون الكيمياء.