Chemical Bonding
Covalent Bonding
5th Year · 6th Year (Leaving Cert)
- ✓By the end of this lesson students will be able to define covalent bonding and explain its formation.
- ✓By the end of this lesson students will be able to draw Lewis structures for simple covalent molecules and polyatomic ions.
- ✓By the end of this lesson students will be able to distinguish between polar and non-polar covalent bonds and molecules.
- ✓By the end of this lesson students will be able to predict molecular shapes using VSEPR theory (Higher Level).
- ✓By the end of this lesson students will be able to describe and explain the different types of intermolecular forces.
Key concepts
Covalent bonding is a type of chemical bond formed when two non-metal atoms share one or more pairs of valence electrons to achieve a stable electron configuration, typically an octet (eight valence electrons). This sharing leads to a strong electrostatic attraction between the shared electrons and the positively charged nuclei of the bonded atoms.
Lewis structures (also known as Lewis dot diagrams or electron dot structures) are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. They are used to visualise the distribution of valence electrons and predict the number of bonds an atom will form.
Electronegativity is a measure of the relative attraction an atom has for a shared pair of electrons in a covalent bond. The Pauling scale is commonly used, with fluorine being the most electronegative element (4.0). Differences in electronegativity determine the polarity of a bond.
A polar covalent bond forms when there is an unequal sharing of electrons between two atoms due to a significant difference in their electronegativities (typically between 0.4 and 1.7 on the Pauling scale). This unequal sharing results in partial positive (δ+) and partial negative (δ-) charges on the atoms.
A non-polar covalent bond forms when electrons are shared equally between two atoms. This occurs either when the two bonded atoms are identical (e.g., H₂ or Cl₂) or when their electronegativity difference is very small (typically less than 0.4).
The overall polarity of a molecule depends on both the polarity of its individual bonds and its molecular geometry. Even if a molecule contains polar bonds, it can be non-polar overall if the bond dipoles cancel each other out due to the molecule's symmetrical shape (e.g., CO₂). If the bond dipoles do not cancel, the molecule is polar.
VSEPR theory is used to predict the three-dimensional geometry of molecules. It states that electron pairs (both bonding pairs and lone pairs) in the valence shell of a central atom repel each other and will arrange themselves as far apart as possible to minimise repulsion, thus determining the molecular shape.
Intermolecular forces are attractive forces that exist between molecules. They are much weaker than the intramolecular covalent bonds within molecules but significantly influence the physical properties of substances, such as boiling points, melting points, and solubility.
London dispersion forces are the weakest type of intermolecular force and are present in all molecules, whether polar or non-polar. They arise from temporary, instantaneous dipoles created by the momentary uneven distribution of electrons around an atom or molecule.
Dipole-dipole forces are attractive forces that occur between the permanent dipoles of polar molecules. The partially positive end of one polar molecule is attracted to the partially negative end of an adjacent polar molecule.
Hydrogen bonding is a particularly strong type of dipole-dipole interaction. It occurs when a hydrogen atom, covalently bonded to a highly electronegative atom (nitrogen, oxygen, or fluorine), is attracted to a lone pair of electrons on another highly electronegative atom (N, O, or F) in an adjacent molecule.
Key facts to remember
- 1Covalent bonds involve the sharing of valence electrons between non-metal atoms.
- 2Lewis structures are used to represent valence electrons and bonding in molecules.
- 3The difference in electronegativity between bonded atoms determines bond polarity.
- 4Molecular polarity depends on both bond polarity and the overall molecular shape (symmetry).
- 5VSEPR theory (HL) states that electron pairs around a central atom repel each other and arrange to minimise repulsion, determining molecular geometry.
- 6Intermolecular forces (IMFs) are attractive forces between molecules, much weaker than intramolecular covalent bonds.
- 7London dispersion forces are present in all molecules and arise from temporary dipoles.
- 8Hydrogen bonding is the strongest type of intermolecular force, occurring when H is bonded to N, O, or F.
Worked examples
Example 1
Draw the Lewis structure for carbon dioxide (CO₂).
Answer
The Lewis structure for CO₂ is: O=C=O, with two lone pairs of electrons on each oxygen atom.
Remember to always check that all atoms (except hydrogen) have an octet and that the total number of electrons used matches the initial count of valence electrons.
Example 2
Predict the molecular shape and overall polarity of ammonia (NH₃) using VSEPR theory.
Answer
Ammonia (NH₃) has a trigonal pyramidal molecular shape and is a polar molecule.
Lone pairs are crucial for determining molecular shape and often contribute to molecular polarity, even though they are not 'bonds'.
Example 3
Explain the significant difference in boiling points between methane (CH₄, boiling point -161.5 °C) and water (H₂O, boiling point 100 °C).
Answer
Methane (CH₄) is a non-polar molecule, so it only experiences weak London dispersion forces between its molecules. Water (H₂O) is a polar molecule and experiences London dispersion forces, dipole-dipole forces, and crucially, strong hydrogen bonding between its molecules. Hydrogen bonds are much stronger than London dispersion forces. Therefore, significantly more energy is required to overcome the strong hydrogen bonds in water to cause it to boil, compared to the weak London dispersion forces in methane, leading to water's much higher boiling point.
Always identify all types of IMFs present, but focus on the strongest type when comparing properties like boiling point.
Common mistakes
- ✗Confusing intramolecular covalent bonds (within a molecule) with intermolecular forces (between molecules).
- ✗Incorrectly counting valence electrons or failing to account for lone pairs when drawing Lewis structures.
- ✗Forgetting to consider lone pairs when determining molecular geometry using VSEPR theory.
- ✗Assuming a molecule with polar bonds is always polar; molecular symmetry can lead to overall non-polarity (e.g., CO₂).
- ✗Incorrectly identifying hydrogen bonding; it only occurs when hydrogen is directly bonded to nitrogen, oxygen, or fluorine.
- ✗Not explaining the link between the strength of intermolecular forces and physical properties like boiling point.
Exam tips
- ★Always show all valence electrons (bonding and lone pairs) clearly in Lewis structures, especially for polyatomic ions where charge is important.
- ★For VSEPR questions, systematically determine electron domains, electron geometry, and then molecular geometry. Clearly state the name of the shape.
- ★When comparing physical properties (e.g., boiling points), explicitly state the types of intermolecular forces present in each substance and compare their relative strengths.
- ★Remember that VSEPR theory applies to the electron pairs around the *central* atom, not necessarily the entire molecule.
- ★For molecular polarity, consider both the polarity of individual bonds and how these bond dipoles add up or cancel out due to the molecule's 3D shape.
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