Simple molecules and covalent bonds Revision Notes

A covalent bond is a chemical bond formed when two atoms share one pair of electrons. This shared pair of electrons is attracted by the nuclei of both atoms, holding them together in a molecule.

Unlike ionic bonding, where electrons are transferred to form ions, covalent bonding involves electron sharing, which allows each atom to count the shared electrons towards completing its outer shell.

The shared electrons occupy the overlapping region between the two atoms, creating a strong bond that holds the atoms together.

• Covalent bond = sharing of one pair of electrons between two atoms.
• Shared electrons are attracted by both nuclei.
• Bond formation leads to molecule creation.

Atoms tend to be more stable when they have a full outer shell of electrons, similar to the electronic configuration of noble gases (Group 0 elements).

Noble gases have complete outer electron shells, making them chemically inert and stable.

By sharing electrons in covalent bonds, atoms can effectively fill their outer shells, achieving a noble gas configuration without losing or gaining electrons permanently.

• Noble gases have full outer electron shells (e.g., 2 electrons for helium, 8 for neon).
• Atoms share electrons to mimic noble gas configurations.
• This leads to chemical stability in molecules.

Hydrogen molecules (H₂) form when two hydrogen atoms share one pair of electrons, giving each hydrogen atom the electronic configuration of helium (1s²).

Chlorine molecules (Cl₂) form by sharing one pair of electrons between two chlorine atoms, each achieving the electron configuration of argon.

Water (H₂O) molecules form when oxygen shares electrons with two hydrogen atoms, allowing oxygen to complete its outer shell with 8 electrons, similar to neon.

• H₂: each H shares 1 electron pair → helium configuration.
• Cl₂: each Cl shares 1 electron pair → argon configuration.
• H₂O: O shares electrons with 2 H atoms → oxygen achieves neon configuration.

Covalent bonds are commonly represented by a single line between atoms in structural formulae, indicating a shared pair of electrons.

For example, H–H represents a single covalent bond between two hydrogen atoms.

Dot-and-cross diagrams can also illustrate covalent bonding by showing the electrons from each atom and the shared pair between them, but structural formulae are simpler and widely used.

• Single line (–) represents one covalent bond (one shared electron pair).
• H–H shows a covalent bond in hydrogen molecule.
• Dot-and-cross diagrams show electron sharing explicitly.

When two atoms approach each other, their outer electrons interact.

If both atoms need electrons to complete their outer shells, they can share electrons, forming a covalent bond.

This sharing allows each atom to count the shared electrons towards a full outer shell, achieving a stable electronic configuration similar to noble gases, which explains the stability of many molecules.

• Atoms share electrons to complete outer shells.
• Shared electrons count for both atoms.
• Resulting molecules have atoms with noble gas-like configurations.

A covalent bond forms when two non-metal atoms share one or more pairs of electrons. This sharing allows each atom to attain a full outer shell, often resembling the electronic configuration of the nearest noble gas, which is a more stable arrangement.

Each shared pair of electrons constitutes a single covalent bond. Atoms can share multiple pairs to form double or triple bonds, but this objective focuses on single bonds in simple molecules.

Dot-and-cross diagrams are a visual tool to represent the electrons from each atom: dots represent electrons from one atom, crosses from the other. Shared pairs appear between the atoms, showing the covalent bond.

• Covalent bonds involve sharing electron pairs.
• Atoms share electrons to complete outer shells.
• Dot-and-cross diagrams show electron sharing and bonding.

Hydrogen atoms have one electron each and need two electrons to fill their outer shell (1s orbital). Two hydrogen atoms share their single electrons, forming one shared pair and a single covalent bond.

This bond holds the two hydrogen atoms together in an H₂ molecule.

Dot-and-cross diagram: each H atom contributes one electron (one dot and one cross), sharing one pair between them.

• H atoms share one electron each.
• One single covalent bond forms.
• H₂ molecule has a stable outer shell for both atoms.

Carbon has four electrons in its outer shell and needs four more to complete the octet.

Each hydrogen atom has one electron and needs one more to fill its shell.

Carbon shares one electron with each of four hydrogen atoms, forming four single covalent bonds, resulting in CH₄ (methane). Each bond is a shared pair of electrons between carbon and hydrogen atoms.

• Carbon shares four electrons with four hydrogens.
• Four single covalent bonds form.
• Each atom achieves a full outer shell.

Oxygen has six electrons in its outer shell and needs two more to complete the octet.

Each hydrogen atom shares one electron with oxygen, forming two single covalent bonds.

The oxygen atom shares one electron with each hydrogen atom, resulting in H₂O (water). The oxygen also has two lone pairs of electrons not involved in bonding.

• Oxygen shares electrons with two hydrogens.
• Two single covalent bonds form.
• Oxygen completes its octet; hydrogens fill their shells.

Iodine atoms each have seven electrons in their outer shell and need one more to complete the octet.

Two iodine atoms share one pair of electrons, forming a single covalent bond in I₂.

Nitrogen has five electrons in its outer shell and needs three more to complete the octet. It shares electrons with three hydrogen atoms, each sharing one electron, forming three single covalent bonds in NH₃ (ammonia). Nitrogen also has one lone pair of electrons.

• I₂: two iodine atoms share one electron pair.
• NH₃: nitrogen shares three electron pairs with hydrogens.
• Atoms achieve full outer shells through sharing.

Carbon has four outer electrons and needs four more to complete its octet.

Each chlorine atom has seven outer electrons and needs one more to complete its octet.

Carbon shares one electron with each of four chlorine atoms, forming four single covalent bonds in CCl₄ (carbon tetrachloride). Each chlorine atom shares one electron with carbon, completing their octets.

• Carbon shares four electrons with four chlorines.
• Four single covalent bonds form.
• All atoms achieve stable electronic configurations.

Simple molecular compounds consist of discrete molecules formed by atoms sharing pairs of electrons through covalent bonds. Each molecule is electrically neutral and held together internally by strong covalent bonds, which involve the sharing of electrons between atoms to achieve stable electronic configurations.

Between these molecules, there are weak intermolecular forces, often van der Waals forces (induced dipole-induced dipole attractions), dipole-dipole interactions, or hydrogen bonds depending on the molecules involved. These forces are much weaker than the covalent bonds within the molecules.

The overall structure is therefore a collection of molecules held together by weak forces, not a continuous network of atoms or ions.

• Molecules held by strong covalent bonds internally.
• Weak intermolecular forces between molecules.
• Discrete molecules, electrically neutral overall.

The melting and boiling points of a substance depend on the strength of the forces that must be overcome to change state. In simple molecular compounds, the strong covalent bonds within molecules do not break during melting or boiling; only the weak intermolecular forces between molecules are overcome.

Because these intermolecular forces are weak, relatively little energy is required to separate the molecules from each other. This results in low melting and boiling points compared to ionic or metallic substances, where strong ionic or metallic bonds must be broken.

For example, iodine (I₂) is a simple molecular solid with a melting point of 114 °C, which is low compared to ionic compounds like sodium chloride (NaCl), which melts at 801 °C. Water (H₂O) has a higher boiling point than many simple molecules due to hydrogen bonding, a stronger type of intermolecular force, but still much lower than ionic compounds.

• Melting and boiling involve overcoming weak intermolecular forces, not covalent bonds.
• Low melting and boiling points compared to ionic or metallic compounds.
• Hydrogen bonding can increase boiling points but still lower than ionic compounds.

Electrical conductivity requires the presence of charged particles that are free to move. In simple molecular compounds, the molecules are neutral and do not dissociate into ions in solid or liquid states.

Because there are no free ions or electrons to carry charge, simple molecular compounds do not conduct electricity when solid or molten.

For example, solid iodine and liquid water do not conduct electricity. In contrast, ionic compounds conduct electricity when molten or dissolved because ions are free to move.

• No free ions or electrons in simple molecular compounds.
• Molecules are neutral and remain intact in solid and liquid states.
• No electrical conductivity in solid or liquid forms.

Water is a simple molecular compound with polar covalent bonds and hydrogen bonding between molecules.

Its melting point is 0 °C and boiling point is 100 °C, higher than many simple molecules due to hydrogen bonds, which are stronger intermolecular forces than van der Waals forces.

Water does not conduct electricity in pure form because it contains no free ions; however, it can conduct when impurities or dissolved ions are present.

• Strong covalent bonds within H₂O molecules.
• Hydrogen bonding causes relatively high melting and boiling points.
• No free ions in pure water, so poor electrical conductivity.

The table below summarises the key chemical details, evidence, and exam-useful points relating to the properties of simple molecular compounds.

Covalent bonds form when two non-metal atoms share one or more pairs of electrons. This sharing allows each atom to attain a full outer shell, often resembling the electronic configuration of the nearest noble gas, which increases stability. Unlike ionic bonds, which involve electron transfer, covalent bonds involve electron sharing between atoms.

Dot-and-cross diagrams are a useful way to represent covalent bonding. In these diagrams, electrons from one atom are shown as dots, and electrons from the other atom as crosses. Shared pairs of electrons between atoms represent covalent bonds. Each shared pair counts as one bond, which can be single, double, or triple depending on the number of shared electron pairs.

• Covalent bonds involve sharing of electron pairs.
• Atoms share electrons to complete their outer shells.
• Dot-and-cross diagrams differentiate electrons from each atom.
• Single bond = one shared pair; double bond = two; triple bond = three.

Methanol consists of one carbon (C), four hydrogen (H), and one oxygen (O) atom. Carbon has four valence electrons, oxygen has six, and hydrogen has one each.

Carbon forms four single covalent bonds: three with hydrogen atoms and one with oxygen. Oxygen forms two single bonds: one with carbon and one with hydrogen, and has two lone pairs of electrons.

The dot-and-cross diagram shows: each C–H bond as a shared pair (one dot and one cross), the C–O bond as a shared pair, and the O–H bond as a shared pair. Lone pairs on oxygen are shown as pairs of dots or crosses not involved in bonding.

• C shares electrons with 3 H atoms and 1 O atom.
• O shares electrons with C and 1 H, plus has 2 lone pairs.
• All bonds in CH₃OH are single covalent bonds.
• Lone pairs on oxygen are important for molecular shape.

Ethylene contains two carbon atoms and four hydrogen atoms. Each carbon has four valence electrons.

The two carbon atoms share two pairs of electrons, forming a double bond between them. Each carbon also forms single bonds with two hydrogen atoms.

The diagram shows each C–H bond as a single shared pair and the C=C double bond as two shared pairs of electrons (two dots and two crosses between the carbons). This double bond is stronger and shorter than a single bond.

• C atoms share two pairs of electrons (double bond).
• Each C atom forms two single bonds with H atoms.
• Double bond represented by two shared electron pairs.
• Double bonds affect molecule reactivity and shape.

Oxygen (O₂) consists of two oxygen atoms sharing two pairs of electrons, forming a double bond. Each oxygen atom also has two lone pairs.

Carbon dioxide (CO₂) has one carbon atom double bonded to two oxygen atoms. Carbon shares two pairs of electrons with each oxygen atom, forming two double bonds. Oxygen atoms have two lone pairs each.

Nitrogen (N₂) consists of two nitrogen atoms sharing three pairs of electrons, forming a triple bond. Each nitrogen atom has one lone pair. The triple bond is very strong and short, making N₂ a very stable molecule.

• O₂ has a double bond (two shared pairs) between oxygen atoms.
• CO₂ has two double bonds between carbon and each oxygen.
• N₂ has a triple bond (three shared pairs) between nitrogen atoms.
• Lone pairs are shown on oxygen and nitrogen atoms.
• Bond type affects bond strength and molecular properties.

Simple molecular compounds consist of molecules formed by atoms joined together by covalent bonds. A covalent bond is the sharing of a pair of electrons between two atoms, resulting in a stable electronic configuration similar to noble gases.

Within each molecule, atoms are held together by strong covalent bonds. However, the forces between separate molecules are much weaker and are called intermolecular forces (also known as van der Waals forces or London dispersion forces).

Examples of simple molecular compounds include water (H₂O), carbon dioxide (CO₂), ammonia (NH₃), and iodine (I₂). Their molecules are discrete units with fixed numbers of atoms.

• Covalent bonds: strong bonds within molecules formed by shared electron pairs.
• Molecules: discrete groups of atoms bonded covalently.
• Intermolecular forces: weak forces between molecules, much weaker than covalent bonds.

The melting point and boiling point of a substance depend on the strength of the forces holding its particles together in the solid or liquid state.

In simple molecular compounds, the strong covalent bonds hold atoms within molecules, but to melt or boil the substance, the weak intermolecular forces between molecules must be overcome.

Because these intermolecular forces are weak, little energy is required to separate the molecules, resulting in low melting and boiling points compared to ionic or metallic substances where strong ionic or metallic bonds must be broken.

• Melting and boiling involve overcoming intermolecular forces, not covalent bonds.
• Weak intermolecular forces mean low melting and boiling points.
• Example: Iodine (I₂) sublimes at 114 °C, much lower than ionic compounds.

Electrical conductivity requires the presence of charged particles that are free to move, such as ions or delocalised electrons.

Simple molecular compounds do not contain free ions or delocalised electrons; their electrons are localised in covalent bonds or lone pairs within molecules.

As a result, they cannot conduct electricity in solid or liquid states because there are no charged particles available to carry electric current.

• No free ions or electrons in simple molecular compounds.
• Electrons are localised in covalent bonds or lone pairs.
• No electrical conductivity in solid or liquid states.

Water is a simple molecular compound with molecules formed by covalent bonds between hydrogen and oxygen atoms.

Within each H₂O molecule, two O–H covalent bonds hold the atoms together strongly.

Between water molecules, hydrogen bonding (a type of strong intermolecular force) exists but is still much weaker than covalent bonds, giving water a relatively higher boiling point compared to other simple molecules like CO₂ but still low compared to ionic compounds like NaCl (melting point 801 °C).

• Strong covalent O–H bonds within molecules.
• Hydrogen bonds between molecules are weaker than covalent bonds.
• Water boils at 100 °C due to hydrogen bonding, but this is low compared to ionic compounds.
• Water does not conduct electricity because it lacks free ions (pure water is a poor conductor).

The table below summarises the key chemical details, evidence, and exam-useful points related to the properties of simple molecular compounds.

To answer exam questions on this topic, ensure you:

• Define simple molecular compounds and covalent bonds clearly.
• Explain that strong covalent bonds hold atoms within molecules, but weak intermolecular forces act between molecules.
• State that low melting and boiling points result from weak intermolecular forces requiring little energy to overcome.
• Explain that poor electrical conductivity is due to the absence of free-moving charged particles.
• Use examples such as H₂O, CO₂, or I₂ to support explanations.
• Include relevant chemical terminology such as 'intermolecular forces', 'covalent bonds', 'electrons localised', and 'no free ions'.