Covalent Bonds and Molecular Bonding – Study Notes (Pages 1-19)

Covalent Bonds

  • Covalent bonds are the sharing of electrons between atoms to form molecules. This sharing allows atoms to achieve a full outer shell (the octet rule) without transferring electrons.
  • Ionic bonds (electron transfer) are different from covalent bonds (electron sharing). In covalent bonding, there are no free ions or free charges in the bond.
  • Covalent bonding primarily involves non-metals; these elements form the most common compounds found on Earth and in living things.

Non-Metals

  • Non-metals make up a very small percentage of the periodic table, yet they are the most common constituents of compounds found on Earth and in living things.

Properties of Covalent Molecules

  • Covalent molecules are generally soft and have low melting points (often gases or liquids).
  • They do not conduct electricity because there are no free electrons or moving charges.
  • Bonding is relatively weak compared to ionic bonds.
  • When covalent compounds dissolve in water, the entire molecules separate, but each individual molecule remains intact (unlike ionic compounds, which dissociate into ions).

Molecules

  • Molecules are groups of 2 or more atoms covalently bonded together.
  • Examples:
    • One type of atom (diatomic molecules) e.g. O_2 (oxygen), He (helium, though He does not form covalent bonds with itself in the same way as diatomic nonmetals).
    • More than one type of atom (molecular compounds): H2O (water), CO2 (carbon dioxide).
  • Notation examples given in the slide: Oxygen: O2; Water: H2O; Carbon dioxide: CO_2.

Molecular Bonds

  • Covalent molecules have two forms of bonding:
    • Intramolecular bonds: the bonds that hold the atoms together within a molecule (e.g., the O–O bond in O2, the H–O bonds in H2O).
    • Intermolecular bonds: bonds between separate molecules; this topic is typically covered later (Year 11).

How Do Molecular Bonds Form

  • Octet rule: atoms aim to have a full outer shell of electrons.
  • Covalent bonding forms when atoms share pairs of electrons (one electron from each atom).
  • Shared electron pairs spend time orbiting both nuclei involved in the bond.
  • The two nuclei repel each other, but the shared electrons create an attraction that holds the nuclei together through electrostatic forces.

How Do Molecular Bonds Form (Water Example)

  • Hydrogen has 1 valence electron and needs 1 more to achieve a full outer shell (valence shell of 2 for hydrogen).
  • Oxygen has 6 valence electrons and needs 2 more to achieve a full octet.
  • In H_2O, two hydrogens share one electron each with oxygen; oxygen shares one electron with each hydrogen.
  • Result:
    • Water molecule has two O–H bonds.
    • Hydrogen atoms each achieve a full 2-electron arrangement (a pair) via sharing with oxygen.
    • Oxygen achieves a full octet (8 electrons around O in the molecule).

Covalent Bonds – Bond Types

  • Covalent bonds can be:
    • Single bonds (one pair of shared electrons): e.g., the H–H bond in H_2.
    • Double bonds (two pairs of shared electrons): e.g., the O=O bond in O_2.
    • Triple bonds (three pairs of shared electrons): e.g., the N≡N bond in N_2.
  • In water, we observe single covalent bonds (O–H).

Displaying Bonds

  • Shell Diagram: shows the overlapping or shared electrons along with all other electrons around the atoms.
  • Valence Structure Diagram (often called a Lewis structure): shows valence electrons and shared electron pairs with lines representing bonds.
  • Electron Dot Diagram (Lewis Dot Structure): draws only the valence electrons and the shared electron pairs.

Examples – Hydrogen Chloride

  • Formula: HCl
  • Most basic display (valence): H-Cl
  • Shell diagram: (preferred bottom)
  • Electron dot (Lewis) diagram: (preferred top)
  • Note: There is only one pair of shared electrons in the H–Cl bond.

Examples – Carbon Dioxide

  • Formula: CO_2
  • Most basic display (valence): O=C=O
  • Shell diagram: (preferred)
  • Electron dot (Lewis): (also preferred)
  • Note: There are two electron pairs shared between each carbon–oxygen bond (each C=O is a double bond).

Your Turn (Practice)

  • Draw the following in at least two different formats (shell and Lewis Dot Structure) and explain, using valence electrons, the number of bonds formed:
    • Fluorine – F_2
    • Ammonia – NH_3
    • Methane – CH_4
    • Ethene – C2H4
    • Hydrogen Cyanide – HCN
    • Ethyne – C2H2

Answers

Fluorine – F_2

  • Fluorine has 7 valence electrons, so each fluorine atom needs 1 more electron to complete its outer shell.
  • When two fluorine atoms come together, they share one electron each to form a single covalent bond: F_2.

Ammonia – NH_3

  • Nitrogen has 5 valence electrons and needs 3 more electrons to complete its octet (3 covalent bonds).
  • Hydrogen has 1 valence electron and needs 1 more electron to form a bond.
  • With three hydrogens, nitrogen forms three N–H covalent bonds, satisfying the octet for nitrogen and giving each hydrogen a filled shell.

Methane – CH_4

  • Carbon has 4 valence electrons and needs 4 more electrons to complete its octet (4 covalent bonds).
  • Hydrogen has 1 valence electron and needs 1 more electron to form a bond.
  • Four hydrogens bond with a central carbon to form methane: CH_4.

Ethene – C2H4

  • Carbon has 4 valence electrons and needs 4 more electrons to complete octets.
  • Hydrogen has 1 valence electron and needs 1 more electron per bond.
  • The two carbon atoms share a double bond (two shared electron pairs) between them, and each carbon forms two single bonds to hydrogens, giving the structure C2H4 with a C=C double bond and two C–H bonds on each carbon.
  • Explanation notes: two carbons share a double bond, so each carbon only needs two more bonds to satisfy its octet.

Hydrogen Cyanide – HCN

  • Nitrogen has 5 valence electrons and usually forms three bonds or has a lone pair; here, it is involved in a triple bond to carbon in the cyanide group, plus a single bond to hydrogen in HCN.
  • Carbon has 4 valence electrons and forms a bond to hydrogen and a triple bond to nitrogen (effectively three pairs to N), totalling four bonds around carbon.
  • Hydrogen has 1 valence electron and forms one bond to carbon.
  • The given explanation: Nitrogen has 5 valence electrons and needs 3 electron pairs; Carbon has 4 valence electrons and needs four pairs (three to nitrogen); Hydrogen has 1 valence electron and needs one pair (one to carbon). The resulting structure is H–C≡N.

Ethyne – C2H2

  • Each carbon has 4 valence electrons and needs four more to complete octets; each hydrogen has 1 valence electron and needs one more electron to bond.

  • The two carbon atoms form a triple bond between them, and each carbon forms one bond to a hydrogen atom, giving the structure H-C\equiv C-H.

  • Summary of bonding patterns:

    • Single bonds: display as H–H, H–O, etc. (one shared electron pair).
    • Double bonds: display as O=O, C=O, etc. (two shared electron pairs).
    • Triple bonds: display as N≡N, C≡C, etc. (three shared electron pairs).

  • Key takeaways:

    • Covalent bonds involve electron sharing to satisfy the octet rule.
    • Bond type (single/double/triple) affects bond strength and molecular geometry.
    • Display methods (shell diagrams, Lewis structures) help visualize bonding and valence electrons.