Cambridge AS & A Level9701

Electronegativity and bonding

Chemistry 9701 Chapter Notes

What this chapter covers

Electronegativity and bondingIonic bondingMetallic bondingCovalent bonding and coordinate (dative covalent) bondingShapes of moleculesIntermolecular forces, electronegativity and bond propertiesDot-and-cross diagrams
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In this chapter

  1. 1Understanding Electronegativity
  2. 2Bond Polarity and Dipoles
  3. 3Polar vs. Non-polar Molecules
  4. 4Intermolecular Forces (IMFs)
  5. 5Special Case: Hydrogen Bonding
  6. 6Predicting Properties from Bonding

1. Understanding Electronegativity

Electronegativity is a fundamental concept that describes how strongly an atom pulls on the electrons in a chemical bond. Think of it as a 'tug-of-war' for electrons between two bonded atoms. The Pauling scale is commonly used to assign a numerical value for the electronegativity of each element. Trends in the periodic table are crucial: electronegativity increases from left to right across a period because the nuclear charge increases while the number of electron shells stays the same, pulling bonding electrons more strongly. Electronegativity decreases down a group because the bonding electrons are in shells further from the nucleus, and there is more 'shielding' by inner electrons, which weakens the nucleus's pull.

Key term

Electronegativity: The measure of the ability of a bonded atom to attract the pair of electrons in a covalent bond towards itself.

Examiner insight

Examiners require a precise definition of electronegativity that specifies it applies to a 'bonded atom' within a 'covalent bond'.

Fun fact

The concept of electronegativity was developed by Linus Pauling, who won a Nobel Prize in Chemistry and also a Nobel Peace Prize, making him one of only two people to have won Nobel Prizes in two different fields.

Worked example 13 marks

Explain why a chlorine atom (Cl) is more electronegative than a sulfur atom (S). [3]

  1. 1
    1. Both chlorine and sulfur are in Period 3 of the periodic table.
  2. 2
    1. A chlorine atom has 17 protons in its nucleus, while a sulfur atom has 16 protons. This means chlorine has a greater nuclear charge.
  3. 3
    1. The outer electrons are in the same shell (n=3) for both atoms, so they experience similar shielding.
  4. 4
    1. The stronger nuclear charge of chlorine attracts the bonding pair of electrons more strongly than sulfur does, making it more electronegative.

Recap

  • Electronegativity measures an atom's pull on bonding electrons in a covalent bond.
  • It increases across a period due to increasing nuclear charge.
  • It decreases down a group due to increased atomic radius and shielding.
  • Fluorine is the most electronegative element.

Quick check

  1. Which element is more electronegative: oxygen (O) or sulfur (S)? Explain your reasoning. [2]2 marks

2. Bond Polarity and Dipoles

When two atoms in a covalent bond have different electronegativities, the electrons are not shared equally. The more electronegative atom pulls the electron pair closer, gaining a slight negative charge (written as δ−, 'delta negative'). The less electronegative atom is left with a slight positive charge (δ+, 'delta positive'). This separation of charge is called a dipole, and the bond is known as a polar covalent bond. If the electronegativity difference is zero (e.g., in Cl₂ or O₂), the electrons are shared perfectly equally, forming a non-polar covalent bond. If the difference is very large (typically >1.7), the 'sharing' is so unequal that the electron is effectively transferred from one atom to the other, forming an ionic bond.

Key term

Polar Covalent Bond: A covalent bond in which the bonding electrons are shared unequally between two atoms with different electronegativities, resulting in a permanent bond dipole.

Common pitfall

Assuming that any bond between two different non-metal atoms is highly polar. Bonds like C-H have a very small electronegativity difference and are often treated as non-polar.

Worked example 13 marks

The electronegativity values for hydrogen, carbon, and fluorine are 2.2, 2.5, and 4.0 respectively. For the C-H and C-F bonds, state which is more polar and add δ+ and δ- symbols to show the polarity of each bond. [3]

  1. 1
    1. Calculate the electronegativity difference (ΔEN) for each bond.
  2. 2

    ΔEN (C-H) = 2.5 - 2.2 = 0.3

  3. 3

    ΔEN (C-F) = 4.0 - 2.5 = 1.5

  4. 4
    1. The C-F bond has a much larger electronegativity difference, so it is significantly more polar than the C-H bond.
  5. 5
    1. For the C-H bond, carbon is slightly more electronegative: C(δ-)—H(δ+). (Note: this bond is often considered effectively non-polar).
  6. 6
    1. For the C-F bond, fluorine is much more electronegative: C(δ+)—F(δ-).

Recap

  • Unequal sharing of electrons in a covalent bond creates a polar bond.
  • The more electronegative atom gets a partial negative charge (δ−).
  • The less electronegative atom gets a partial positive charge (δ+).
  • A large electronegativity difference leads to an ionic bond.

Quick check

  1. Is the bond in a nitrogen molecule (N₂) polar or non-polar? Explain why. [1]1 mark
  2. Place a δ+ and a δ- symbol on the appropriate atoms in a molecule of hydrogen bromide, HBr. [1]1 mark

3. Polar vs. Non-polar Molecules

A molecule can have polar bonds but be non-polar overall. This happens when the molecule has a symmetrical shape, causing the individual bond dipoles to cancel each other out. Think of it like a tug-of-war where two equally strong teams are pulling in opposite directions – there is no net movement. For example, in carbon dioxide (CO₂), the two C=O bond dipoles are equal and point in opposite directions in a linear molecule, so they cancel out, making CO₂ non-polar. In tetrachloromethane (CCl₄), the four C-Cl dipoles point to the corners of a regular tetrahedron, and their effects also cancel out. However, if the molecule's shape is asymmetrical, the dipoles do not cancel, and the molecule has an overall dipole moment, making it a polar molecule. Examples include water (H₂O), which is bent, and ammonia (NH₃), which is trigonal pyramidal.

Key term

Dipole Moment: A measure of the overall polarity of a molecule, which is determined by the magnitude and direction of all its bond dipoles.

Examiner insight

To earn full marks when explaining molecular polarity, you must discuss both bond polarity (due to electronegativity differences) and molecular shape (symmetry).

Worked example 14 marks

Tetrachloromethane (CCl₄) is a non-polar molecule, but trichloromethane (CHCl₃) is polar. Explain this difference, referring to bond polarity and molecular shape. [4]

  1. 1
    1. Both molecules have a tetrahedral shape around the central carbon atom.
  2. 2
    1. In CCl₄, all four bonds are C-Cl bonds. These bonds are polar because chlorine is more electronegative than carbon.
  3. 3
    1. Because the tetrahedral shape is perfectly symmetrical, the four equal C-Cl bond dipoles point in opposing directions and cancel each other out. This results in no overall dipole moment, so CCl₄ is non-polar.
  4. 4
    1. In CHCl₃, there are three polar C-Cl bonds and one C-H bond (which is almost non-polar). The molecule is not symmetrical. The dipoles from the three C-Cl bonds do not cancel out, creating a net dipole moment. Therefore, CHCl₃ is a polar molecule.

Recap

  • A non-polar molecule has no overall dipole moment.
  • A polar molecule has an overall dipole moment.
  • Symmetrical molecules with polar bonds can be non-polar if the bond dipoles cancel.
  • Asymmetrical molecules with polar bonds are always polar.
  • Examples of non-polar molecules with polar bonds include CO₂, BF₃, and CCl₄.

Quick check

  1. Is a molecule of beryllium chloride, BeCl₂, which is linear, polar or non-polar? [1]1 mark

4. Intermolecular Forces (IMFs)

Intermolecular forces (IMFs) are the weak attractive forces that exist *between* molecules. They are much weaker than the strong covalent, ionic, or metallic bonds *within* a substance. The strength of these forces determines a simple molecular substance's physical properties, like its melting and boiling point. There are three main types:

  1. Van der Waals' forces (induced dipole-induced dipole): These are the weakest IMFs and exist between ALL molecules. They are caused by the random movement of electrons, which creates temporary, instantaneous dipoles. These dipoles can then induce dipoles in neighbouring molecules, leading to a weak attraction. The more electrons a molecule has, the stronger the van der Waals' forces.
  1. Permanent dipole-dipole forces: These are stronger and exist *in addition* to van der Waals' forces, but only between polar molecules. The permanent δ+ on one molecule attracts the permanent δ- on a neighbouring molecule.
  1. Hydrogen bonding: This is the strongest type of IMF. It is a special, strong case of permanent dipole-dipole attraction. It only occurs when a hydrogen atom is covalently bonded to a very electronegative atom with a lone pair of electrons: nitrogen (N), oxygen (O), or fluorine (F).

Key term

Intermolecular Force: A weak attractive force that exists between neighbouring molecules, responsible for the physical properties of simple molecular substances.

Common pitfall

Confusing strong intramolecular covalent bonds with weak intermolecular forces. When a simple molecular substance boils, it is the weak IMFs between molecules that are overcome, not the strong covalent bonds within them.

Worked example 13 marks

The boiling points of fluorine (F₂) and chlorine (Cl₂) are -188 °C and -34 °C respectively. Explain this difference. [3]

  1. 1
    1. Both F₂ and Cl₂ are non-polar molecules.
  2. 2
    1. The only intermolecular forces present are van der Waals' forces.
  3. 3
    1. A chlorine molecule (Cl₂) has more electrons (34) than a fluorine molecule (F₂) (18).
  4. 4
    1. This results in stronger van der Waals' forces between Cl₂ molecules compared to F₂ molecules, which require more energy to overcome, leading to a higher boiling point.

Recap

  • IMFs are weak forces between molecules.
  • Van der Waals' forces exist in all molecules and increase with the number of electrons.
  • Permanent dipole-dipole forces exist between polar molecules.
  • Hydrogen bonding is the strongest IMF, occurring when H is bonded to N, O, or F.

Quick check

  1. Name the strongest type of intermolecular force present in a sample of liquid propanone (CH₃COCH₃). [1]1 mark

5. Special Case: Hydrogen Bonding

Hydrogen bonding is the heavyweight champion of intermolecular forces. It's a particularly strong type of permanent dipole-dipole force that has profound effects on the properties of substances. The strict requirements for hydrogen bonding are: a hydrogen atom must be covalently bonded to a highly electronegative atom (specifically nitrogen, oxygen, or fluorine), and this hydrogen is then attracted to a lone pair of electrons on a neighbouring N, O, or F atom. This creates a strong attraction because:

  1. The N-H, O-H, or F-H bond is highly polar, creating a large δ+ on the hydrogen atom.
  2. The hydrogen atom is very small, so it can get very close to the lone pair on the adjacent molecule.

Water (H₂O) is the classic example. Its surprisingly high boiling point, high surface tension, and the fact that ice is less dense than liquid water are all direct consequences of extensive hydrogen bonding. In ice, the molecules are held in a fixed, open lattice structure by hydrogen bonds, which is less dense than the more disordered arrangement in liquid water.

Key term

Hydrogen Bond: A strong intermolecular attraction between a lone pair of electrons on a nitrogen, oxygen, or fluorine atom and a hydrogen atom that is covalently bonded to another nitrogen, oxygen, or fluorine atom.

Examiner insight

When describing a hydrogen bond, always mention the attraction to a lone pair on the N, O, or F atom. Simply stating it's an attraction between molecules is not specific enough for full marks.

Fun fact

The structure of DNA's double helix is held together by thousands of hydrogen bonds between the base pairs (A-T and G-C). These bonds are strong enough to hold the strands together but weak enough to be 'unzipped' for replication.

Worked example 13 marks

The boiling points of the hydrogen halides are: HCl (-85 °C), HBr (-67 °C), HI (-35 °C). The boiling point of HF is +20 °C. Explain why HF has an anomalously high boiling point. [3]

  1. 1
    1. Fluorine is a very highly electronegative atom. The H-F bond is therefore very polar.
  2. 2
    1. Strong intermolecular forces called hydrogen bonds can form between the δ+ hydrogen atom of one HF molecule and the lone pair on the δ- fluorine atom of a neighbouring HF molecule.
  3. 3
    1. These hydrogen bonds are much stronger than the van der Waals' forces and permanent dipole-dipole forces present in HCl, HBr, and HI. A large amount of energy is required to overcome these strong hydrogen bonds, resulting in a much higher boiling point for HF.

Recap

  • Hydrogen bonds form between molecules containing N-H, O-H, or F-H bonds.
  • The δ+ hydrogen is attracted to a lone pair on a neighbouring N, O, or F atom.
  • Hydrogen bonds are the strongest type of intermolecular force.
  • They are responsible for water's high boiling point and the fact that ice floats.

Quick check

  1. Can hydrogen bonds form between molecules of methane (CH₄)? Explain why or why not. [2]2 marks

6. Predicting Properties from Bonding

By identifying a substance's bonding and structure, we can predict its physical properties. There are four main categories to know:

Giant Ionic Lattice (e.g., NaCl, MgO): Consists of a regular arrangement of positive and negative ions held by strong electrostatic forces. This requires a lot of energy to break, so they have very high melting and boiling points. They are brittle and cannot conduct electricity as solids (ions are fixed), but can when molten or dissolved (ions are mobile).

Giant Metallic Lattice (e.g., Fe, Al): A lattice of positive metal ions surrounded by a 'sea' of delocalised electrons. The strong electrostatic attraction between ions and electrons gives metals high melting points and makes them good conductors of heat and electricity in both solid and liquid states. They are malleable and ductile.

Giant Covalent Structure (e.g., Diamond, Graphite, SiO₂): A huge network of atoms joined by strong covalent bonds. To melt or boil these substances, these strong bonds must be broken, requiring immense amounts of energy. Thus, they have extremely high melting and boiling points and are typically very hard and insoluble. They are non-conductors (except for graphite).

Simple Molecular Structure (e.g., CO₂, H₂O, I₂): Composed of discrete molecules with strong covalent bonds *within* them, but only weak intermolecular forces (IMFs) *between* them. Melting or boiling only requires overcoming these weak IMFs, so they have low melting and boiling points. They are poor conductors of electricity as there are no free-moving charged particles.

Key term

Giant Covalent Structure: A three-dimensional structure of atoms that are all bonded to each other by strong covalent bonds, resulting in a very high melting and boiling point.

Worked example 14 marks

Explain the difference in melting point between magnesium oxide (MgO, m.p. 2852 °C) and sulfur dioxide (SO₂, m.p. -72 °C) in terms of their structure and bonding. [4]

  1. 1
    1. Magnesium oxide (MgO) has a giant ionic lattice structure. It consists of Mg²⁺ and O²⁻ ions held together by strong electrostatic forces of attraction.
  2. 2
    1. A large amount of energy is required to overcome these strong forces and break down the lattice, resulting in a very high melting point.
  3. 3
    1. Sulfur dioxide (SO₂) has a simple molecular structure. Within the molecules are strong covalent bonds, but between the molecules, there are only weak intermolecular forces (van der Waals' and permanent dipole-dipole forces).
  4. 4
    1. Only a small amount of energy is needed to overcome these weak intermolecular forces, resulting in a low melting point.

Recap

  • Giant structures (ionic, metallic, covalent) have high melting points due to strong bonds throughout.
  • Simple molecular structures have low melting points due to weak intermolecular forces.
  • Electrical conductivity requires mobile charged particles (ions or electrons).
  • Metals conduct as solids and liquids; ionic compounds conduct only when molten or aqueous.

Quick check

  1. A substance is a white crystalline solid at room temperature. It dissolves in water to form a solution that conducts electricity. What type of bonding and structure does it have? [2]2 marks

End-of-chapter exercise

Test yourself on the whole chapter. Work through these before moving on.

  1. Define the term 'electronegativity' and explain why fluorine is more electronegative than iodine. [3]3 marks
  2. Draw a 'dot-and-cross' diagram to show the bonding in a molecule of tetrachloromethane, CCl₄. Predict the shape of the molecule and state the bond angle. [3]3 marks
  3. Explain why the boiling point of ammonia (NH₃) is significantly higher than the boiling point of phosphine (PH₃). [4]4 marks
  4. Boron trifluoride (BF₃) has polar bonds but is a non-polar molecule. Explain this observation, stating the shape of the BF₃ molecule. [3]3 marks
  5. Describe the bonding in a typical metal. Use your description to explain why metals are good electrical conductors and have high melting points. [4]4 marks
  6. Ethene, C₂H₄, is a planar molecule with a H-C-H bond angle of approximately 120°. Explain the shape and bond angle in ethene by referring to orbital hybridisation and the types of bonds present. [5]5 marks
  7. Compare and contrast the structure, bonding, and melting points of sodium oxide (Na₂O) and silicon dioxide (SiO₂). [5]5 marks
  8. The boiling points of butane (C₄H₁₀) and propanone (CH₃COCH₃) are -1 °C and 56 °C respectively, despite having similar relative molecular masses. Explain this difference in terms of the intermolecular forces present in each substance. [4]4 marks
  9. Predict the shape of the sulfur hexafluoride (SF₆) molecule. Use the VSEPR theory to justify your answer and state the F-S-F bond angle. [3]3 marks
  10. Astatine (At) is at the bottom of Group 17. Predict whether the boiling point of hydrogen astatide (HAt) would be higher or lower than that of hydrogen iodide (HI, boiling point -35 °C). Justify your prediction by considering all relevant intermolecular forces. [4]4 marks

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