Chemistry (9701) | Electronegativity and bonding

Welcome to Electronegativity and Bonding!

Hello future chemist! This chapter is absolutely fundamental because it explains why atoms join together and how they share (or don't share!) electrons. Understanding these concepts – especially electronegativity – will help you predict the type of bond formed and the physical properties (like boiling point and solubility) of almost every substance you study.
Don't worry if these terms seem complex; we'll break them down using simple language and relatable examples. Let's get started!

---

1. Electronegativity: The Electron Tug-of-War

1.1 Defining Electronegativity

Definition: Electronegativity is defined as the power of an atom to attract electrons to itself (when the atom is part of a compound).
Imagine two atoms bonded together fighting over the shared electrons. The atom with the higher electronegativity is the one that pulls harder in this electron 'tug-of-war'.

Did you know? Fluorine (F) is the most electronegative element in the Periodic Table. It is the ultimate electron bully!

1.2 Factors Influencing Electronegativity

The strength of this pulling power depends on three main factors. Think about how the nucleus (positive charge) attracts the bonding electrons:

• Nuclear Charge (Proton Number): The more protons in the nucleus (higher nuclear charge), the stronger the attraction for the bonding electrons. (Electronegativity increases).
• Atomic Radius: If the bonding electrons are further away from the nucleus (larger atomic radius), the attraction is weaker. (Electronegativity decreases).
• Shielding by Inner Shells: Inner electrons repel the valence (outer) electrons, reducing the effective pull from the nucleus. More inner shells mean more shielding. (Electronegativity decreases).

1.3 Periodic Trends in Electronegativity (LO 3.1.3)

Understanding the factors above helps us explain the trends across the Periodic Table:

1. Across a Period (Left to Right): Electronegativity Increases
   • Nuclear charge increases (more protons).
   • Atomic radius decreases slightly (electrons are pulled closer).
   • Shielding remains relatively constant (same number of inner shells).
   • Result: The nucleus has a greater pull on the bonding electrons.
2. Down a Group (Top to Bottom): Electronegativity Decreases
   • Atomic radius increases (more electron shells are added).
   • Shielding increases (more inner electron shells).
   • Result: The bonding electrons are further away and better shielded, weakening the nuclear pull.

---

2. The Polarity Spectrum: Ionic vs. Covalent Bonds

2.1 Predicting Bond Type using Pauling Values (LO 3.1.4)

The type of bond formed between two atoms depends on the difference in their electronegativity (\(\Delta EN\)). Pauling electronegativity values are used to predict this (and will be provided in exams if needed).

• If \(\Delta EN\) is small (or zero): The electrons are shared relatively equally. This is a Covalent Bond. (Example: H-Cl)
• If \(\Delta EN\) is large: One atom pulls so strongly that the electrons are essentially transferred completely, forming ions. This is an Ionic Bond. (Example: Na-Cl)

In reality, bonding is a spectrum. Almost no bond is 100% ionic or 100% covalent!

2.2 Ionic Bonding (LO 3.2.1)

Ionic bonding occurs when the electronegativity difference is very large, usually between metals (low EN) and non-metals (high EN).

Definition: Ionic bonding is the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions).

The attraction holds the ions tightly together in a rigid, repeating structure called a giant ionic lattice (covered further in Topic 4.2).

Syllabus Examples of Ionic Compounds (LO 3.2.2):
Sodium chloride (NaCl), Magnesium oxide (MgO), and Calcium fluoride (CaF₂).

2.3 Covalent Bonding (LO 3.4.1)

Covalent bonding involves the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. This usually occurs between two non-metals.

When the \(\Delta EN\) is small, the bond is non-polar covalent (e.g., in H₂ or Cl₂). If the \(\Delta EN\) is non-zero, the bond is polar covalent.

---

3. Polar Bonds and Molecular Dipole Moments

3.1 Bond Polarity (LO 3.6.2)

When a covalent bond is formed between two atoms with different electronegativities, the shared electrons are pulled closer to the more electronegative atom.

• This results in unequal sharing.
• The more electronegative atom gains a partial negative charge, denoted \(\delta-\).
• The less electronegative atom gains a partial positive charge, denoted \(\delta+\).

A bond with partial charges is called a polar covalent bond.

Example: In H-Cl, chlorine (EN=3.2) is more electronegative than hydrogen (EN=2.2). The electrons spend more time near Cl, so we write H\(\delta+\)-Cl\(\delta-\).

3.2 Dipole Moments and Molecular Polarity (LO 3.6.2)

The overall polarity of a molecule is measured by its dipole moment. This depends on two things:

1. The polarity of the individual bonds.
2. The shape of the molecule.

The dipole moment is essentially the sum of the polarity 'vectors' (arrows) for all the bonds in the molecule.

• Polar Molecules: If the individual bond dipoles do not cancel out due to an asymmetrical shape, the molecule has an overall dipole moment and is polar. (Example: Water, H₂O, where the V-shape means the bond dipoles point upwards and result in an overall dipole.)
• Non-Polar Molecules: If the bond dipoles cancel out due to a symmetrical shape, the molecule has a zero net dipole moment and is non-polar. (Example: Carbon dioxide, CO₂, is linear. The dipoles pull equally in opposite directions, cancelling each other out.)

---

4. Intermolecular Forces (IMFs) (LO 3.6.3)

4.1 Strong Bonds vs. Intermolecular Forces (LO 3.6.4)

It is vital to distinguish between forces within a molecule (intramolecular/chemical bonds) and forces between molecules (intermolecular forces, IMFs).

• Strong Bonds: Ionic, Covalent, and Metallic bonds are all strong chemical bonds.
• Intermolecular Forces (IMFs): These are weak forces between separate molecules. When you melt or boil a substance like water or methane, you are breaking the IMFs, not the strong covalent bonds inside the molecule.

Important Principle: Strong chemical bonds (Ionic, Covalent, Metallic) are significantly stronger than intermolecular forces.

4.2 Van der Waals' Forces (The Generic Term) (LO 3.6.3a)

The term van der Waals' forces is used as a generic term to describe all intermolecular forces between molecular entities other than those due to bond formation.

There are two main types of van der Waals' forces (excluding hydrogen bonding for now):

A. Instantaneous Dipole-Induced Dipole Forces (id-id)

These are also called London Dispersion Forces (LDFs).

1. Electrons are constantly moving, so at any given instant, the electron cloud in an atom or molecule may be unevenly distributed. This creates a temporary, instantaneous dipole.
2. This temporary dipole can influence a neighbouring atom, causing its electron cloud to shift, creating an induced dipole.
3. The weak attraction between these temporary dipoles is the LDF.

• They are the weakest IMF.
• They are present in all molecules (both polar and non-polar).
• Their strength increases with the number of electrons (and surface area) because larger molecules are easier to polarise. This explains why I₂ (more electrons, larger size) is a solid, while Cl₂ (fewer electrons) is a gas (LO 11.1.3).

B. Permanent Dipole-Permanent Dipole Forces (pd-pd)

These occur only between molecules that are already polar (they possess a permanent dipole moment, Section 3.2).

• The permanent partial positive charge (\(\delta+\)) on one molecule is attracted to the permanent partial negative charge (\(\delta-\)) on a neighbouring molecule.
• Pd-pd forces are stronger than LDFs (for similarly sized molecules).

---

5. Hydrogen Bonding: The Special Case (LO 3.6.1)

5.1 Defining Hydrogen Bonding (LO 3.6.3c)

Hydrogen bonding is a special, very strong type of permanent dipole-permanent dipole force.

It only occurs when a hydrogen atom (H) is directly bonded to one of three highly electronegative atoms: Nitrogen (N), Oxygen (O), or Fluorine (F).

Why is it so strong?

1. The N, O, or F atom is highly electronegative, creating a very strong bond dipole (H\(\delta+\)—X\(\delta-\)).
2. The hydrogen atom is very small and has no other electrons shielding its nucleus.
3. The N, O, or F atom has a lone pair of electrons which forms the attraction site.

The hydrogen bond is the attraction between the H\(\delta+\) atom on one molecule and the lone pair on the N, O, or F atom of a neighbouring molecule.

Syllabus Examples (LO 3.6.1a):
Water (H₂O) and Ammonia (NH₃).

5.2 Hydrogen Bonding and Water's Anomalous Properties (LO 3.6.1b)

Hydrogen bonding is responsible for several unique properties of water, making it essential for life. These properties are often described as "anomalous" because they break the trends expected for hydrides in Group 16.

1. Relatively High Melting and Boiling Points

We expect H₂O to boil at a much lower temperature (based on its small size) but H₂O, NH₃, and HF all have unexpectedly high boiling points.

• Explanation: To boil water, you must provide extra energy to overcome the very strong hydrogen bonds holding the liquid molecules together. This requires far more energy than overcoming the weaker LDFs in larger, non-H-bonded molecules.

2. Relatively High Surface Tension

Water molecules are held together tightly at the surface due to hydrogen bonds pulling the surface molecules inward and sideways.

3. Density of Ice Compared with Liquid Water

Most substances become denser when they solidify, but ice is less dense than liquid water, which is why ice floats.

• Explanation: When water freezes, the hydrogen bonds force the molecules into an ordered, highly structured giant lattice.
• This structure holds the molecules further apart than they are in the constantly shifting liquid state.
• This open structure means that a given mass of ice occupies a larger volume than the same mass of liquid water, resulting in lower density.