Covalent Bonding Notes

Chemical Bonds

  • A chemical bond is a lasting attraction between atoms that enables the formation of chemical compounds.

  • It can result from:

    • Electrostatic force of attraction between atoms with opposite charges.

    • Sharing of electrons (covalent bonds).

  • Bond strengths vary:

    • Strong bonds: covalent, ionic.

    • Weak bonds: dipole-dipole interaction, London dispersion force, hydrogen bonding.

Summary of Bonding

Force

Model

Bonding

Basis of Attraction

Energy (kJ/mol)

Example

Ionic

Cation-anion

Ionic

Cation-anion

400−4000400-4000

NaCl

Covalent

Nuclei-shared electron pair

Covalent

Nuclei-shared electron pair

150−1100150-1100

H-H

Metallic

Cations-delocalized electrons

Metallic

Cations-delocalized electrons

75−100075-1000

Fe

Covalent Bonding

  • Occurs between non-metals.

  • Atoms share electrons to achieve a stable, noble gas configuration.

  • Molecules formed are neutral.

  • Covalent Bond: A strong bond between two atoms where each atom shares one or more electrons.

  • Held together by electrostatic force.

Formation of Covalent Bonds

  • When two non-metal elements react, a covalent compound forms.

  • Covalent bond occurs because two orbitals, each containing one electron, overlap.

  • Strength of the bond is proportional to the amount of overlap of the orbitals.

  • Atoms are held together by the attraction of the positive nuclei to the pair of negative electrons in the bonding orbital.

Multiple Bonds

  • Occur when more than one pair of orbitals overlap, resulting in more than one pair of electrons being shared.

  • Multiple bonds are stronger than single bonds.

Covalent Bonding in Elements

  • Non-metal atoms share electrons to fill their outer shell and become stable.

  • The strong force that joins the atoms together is called a covalent bond.

  • Many elements exist as molecules (two or more atoms joined by a covalent bond).

  • Each atom has a full outer electron shell and is stable.

  • Only the outer shell of electrons is involved in covalent bonding.

  • Two common ways to indicate a covalent bond:

    • Dot and cross diagram.

    • Solid line.

Specific Examples of Covalent Bonding in Elements

  • Hydrogen (electron configuration: 1) needs 1 more electron.

    • Forms H2 or H–H.

    • Diatomic molecule (two atoms).

    • Shares an electron with another hydrogen atom, creating a single bond.

  • Chlorine (2.8.7) needs 1 more electron.

    • Forms Cl2 or Cl–Cl.

    • Shares an electron with another chlorine atom, creating a single bond.

  • Oxygen (2.6) needs 2 more electrons.

    • Forms O2 or O=O.

    • Shares two electrons with another oxygen atom, creating a double bond.

  • Nitrogen (2.5) needs 3 more electrons.

    • Forms N2 or N≡N.

    • Shares three electrons with another nitrogen atom, creating a triple bond.

Covalent Bonding in Compounds

  • Covalent bonding can occur between atoms of different elements.

  • Can be single, double, or triple bonds.

  • Hydrogen (1) and chlorine (2.8.7) each need 1 more electron.

    • Share one electron each to fill their outer shells and become stable.

    • Forms HCl or H–Cl.

  • Oxygen (2.6) needs 2 more electrons, hydrogen (1) needs 1 more.

    • Oxygen atom shares 1 electron with 1 hydrogen atom, and a second electron with another hydrogen atom.

    • Forms H2O or H–O–H.

  • Nitrogen and hydrogen in ammonia (NH3).

    • Nitrogen (2.5) needs 3 electrons.

    • Hydrogen (1) needs 1 electron.

    • Ratio of atoms: 1 N to 3 H.

  • Carbon and hydrogen in methane (CH4).

    • Carbon (2.4) needs 4 electrons.

    • Hydrogen (1) needs 1 electron.

    • Ratio of atoms: 1 C to 4 H.

Covalent Bonding Diagrams

  • Examples provided for:

    • Hydrogen sulfide (H2S).

    • Carbon dioxide (CO2).

    • Ethane (C2H6).

True or False Statements

  • Atoms bond to complete their outer shell, not half complete it.

  • Non-metals commonly form covalent bonds.

  • A covalent bond is a shared pair of electrons (TRUE).

  • Oxygen molecules contain a double bond (TRUE).

  • Sodium chloride contains ionic bonds, not covalent bonds.

  • Covalent bonds can occur in elements and compounds, not only in compounds.

Properties of Covalent Compounds

  • Intermolecular attraction between covalently bonded molecules is generally weak.

  • Low melting and boiling points.

  • Often poor conductors of electricity.

Bond Polarity

  • Unequal sharing of electrons in a covalent bond, leading to a separation of electric charge along the bond.

Electronegativity

  • The ability of an element to attract the bonding electron pair in a covalent bond.

  • Measured on the Pauling scale.

Factors Affecting Electronegativity

  1. Atomic Radius

    • As the radius increases, the bonding pair of electrons becomes further from the nucleus.

    • Lower electronegativity with larger radius.

  2. Number of Unshielded Protons

    • Greater number of protons in a nucleus leads to a greater attraction to the electrons in the covalent bond.

    • Full energy levels of electrons shield the electrons in the bond from the increased attraction of the greater nuclear charge, reducing electronegativity.

Electronegativity Trends

  • Across a Period

    • Electronegativity increases.

      1. Atomic radius decreases.

      2. Nuclear charge increases without significant extra shielding.

  • Down a Group

    • Electronegativity decreases.

      1. Atomic radius increases.

      2. Shielding increases significantly as electrons fill new principal energy levels.

Non-Polar Bonds

  • Electronegativity of both atoms is identical.

  • Electrons are equally attracted, leading to a symmetrical distribution of electron density.

  • Bonding in elements (e.g., O2, Cl2) is always non-polar.

Polar Bonds

  • Occur when a covalent bond is formed between two different atoms.

  • Atoms have different electronegativities, leading to unequal attraction of electrons.

  • Asymmetrical distribution of electron density results, creating a dipole.

  • The atom with the greater share of electron density has a slight negative charge (δ−\delta^-), and the atom with a lesser share has a slight positive charge (δ+\delta^+).

Electronegativity Difference and Polarization

  • The greater the electronegativity difference, the greater the polarization of the bond.

  • Example: Hydrogen halides (HF, HCl, HBr, HI) show decreasing polarization with decreasing electronegativity difference.

Electronegativity values

  • H = 2.2

  • F = 4.0

  • Cl = 3.2

  • Br = 3.0

  • I = 2.7

Electronegativity Difference

  • H-F = 1.8.

  • H-Cl = 1.0

  • H-Br = 0.8

  • H-I = 0.5

Ionic vs. Covalent Bonds

  • Ionic and covalent bonds are at the extremes of a scale.

  • Less polar bonds have more covalent character.

  • More polar bonds have more ionic character.

  • If the difference in electronegativity is large enough, the more electronegative atom can ionize the other atom.

Electronegativity and Bond Type

  • 1. 7 to 4.0: Ionic Bond

  • 0. 3 to 1.7: Polar Covalent Bond

  • 0. 0 to 0.3: Non-Polar Covalent Bond

Polar Molecules

  • Molecules containing polar bonds are not always polar.

  • If polar bonds are arranged symmetrically, the partial charges cancel out, and the molecule is non-polar.

  • If polar bonds are arranged asymmetrically, the partial charges do not cancel out, and the molecule is polar.

Examples of Polar and Non-Polar Molecules

  • Hydrogen sulfide (H2S): polar

  • Ammonia (NH3): polar

  • Boron trifluoride (BF3): non-polar

  • Carbon dioxide (CO2): non-polar

  • Chloromethane (CH3Cl): polar

  • Methane (CH4): non-polar

  • Tetrachloromethane (CCl4): non-polar

  • Water (H2O): polar

  • Methanol (CH3OH): polar

  • Carbon monoxide (CO): polar

Simple Covalent Structures

  • Atoms joined by covalent bonding can form different types of structures.

  • Molecules (e.g., oxygen, water, carbon dioxide) have a simple structure because they contain only a few atoms.

  • Most molecular substances are gas or liquid at room temperature; some are solid (molecular solids).

Molecular Solids

  • Iodine is a molecular solid.

  • Millions of iodine molecules are held together by weak forces of attraction.

  • Two iodine atoms form a single covalent bond to become an iodine molecule.

Properties of Molecular Solids

  • Low melting and boiling points.

  • Usually soft and brittle.

  • Insoluble in water but soluble in other solvents (e.g., petrol).

  • Cannot conduct electricity.

Giant Covalent Structures

  • Millions of atoms join together by covalent bonding, forming giant covalent structures (not molecules).

  • Very high melting and boiling points.

  • Usually hard.

Allotropes of Carbon

  • Different forms of carbon in which atoms bond in different ways (e.g., diamond and graphite).

  • Same chemical properties (same number of electrons) but different physical properties (electrons are shared differently).

Structure and Properties of Diamond

  • Each carbon atom is covalently bonded to four others, creating a giant lattice.

  • Very hard – the hardest natural substance on Earth.

  • Very high melting and boiling point.

  • Cannot conduct electricity.

Structure and Properties of Graphite

  • Each carbon atom is covalently bonded to three others, forming rings of six atoms in layers.

  • Layers are held together by weak forces of attraction.

  • Soft and slippery – layers can easily slide over each other.

  • Can conduct electricity (only non-metal to do so).

Other Allotropes of Carbon

  • Buckminsterfullerene (C60).

    • 60 carbon atoms, each bonded to three others (two single bonds and one double bond).

    • Arranged in 12 pentagons and 20 hexagons to form spheres (bucky balls).

Sand

  • Impure form of silicon dioxide (quartz).

  • Giant covalent structure similar to diamond.

  • Each silicon atom (2.8.4) is bonded to four oxygen atoms, and each oxygen atom (2.6) is bonded to two silicon atoms.

Glossary (part 1)

  • Allotrope: A structurally different form of an element with different physical properties.

  • Covalent Bond: A strong bond between two atoms in which each atom shares one or more electrons with the other.

  • Covalent Compound: A compound containing atoms joined by covalent bonds.

  • Double Bond: A covalent bond in which each atom shares two of its electrons.

  • Giant Structure: A structure containing millions of atoms or ions bonded together, extending in three dimensions.

Glossary (part 2)

  • Molecule: A simple structure containing two or more atoms covalently bonded together.

  • Molecular Solid: A solid substance made up of molecules held together by weak forces of attraction, forming a lattice.

  • Single Bond: A covalent bond in which each atom shares one of its electrons.

  • Triple Bond: A covalent bond in which each atom shares three of its electrons.