CIE Chemistry Topic 3: Chemical Bonding Study Guide

Introduction to Chemical Bonding

• This material is specifically designed for the Cambridge International (CIE) Topic 3: Chemical Bonding syllabus.
• Chemical bonding is fundamental to chemistry as it describes how atoms and ions join together to create molecules.
• The content covers ionic, covalent, and metallic bonding, alongside advanced concepts such as orbital hybridization.
• Knowledge from Topic 1 (Atomic Structure) and Topic 2 (Electrons in Atoms), specifically regarding energy levels and orbitals, is foundational for understanding the bonding mechanisms discussed here.

Ionic Bonding

• Definition: Ionic bonding is the electrostatic attraction between oppositely charged ions or ionic entities.
• Electrostatic Attraction: "Electro" refers to charge, "static" implies the charges are not moving, and "attraction" indicates the binding force between them.
• Formation Process: Atoms form ions to achieve a full valence (outer) shell of electrons.
  • Example: Sodium ($Na$) and Chlorine ($Cl$).
  • Sodium is in Group 1 and has $1$ valence electron.
  • Chlorine is in Group 7 and has $7$ valence electrons.
  • Transition: Sodium loses an electron to become $Na^+$; Chlorine accepts an electron to become $Cl^-$.
  • These oppositely charged particles attract, forming an ionic bond.
• Dot-and-Cross Diagrams: Used to represent the valence electrons in atom-to-ion transitions.
• Common Ions by Group:
  • Group 1: $1+$
  • Group 2: $2+$
  • Group 3: $3+$
  • Group 5: $3-$
  • Group 6: $2-$
  • Group 7: $1-$
  • Noble Gases: Generally do not form ions because they possess a full outer shell.
  • Transition Elements ($d$-block): Can form ions with variable charges.
• Molecular Ions to Memorize:
  • Hydroxide: $OH^-$
  • Nitrate: $NO_3^-$
  • Ammonium: $NH_4^+$
  • Sulfate: $SO_4^{2-}$
  • Carbonate: $CO_3^{2-}$
  • Hydrogen Carbonate: $HCO_3^-$

Determining Chemical Formulas: The Swap and Drop Method

• The Method: A reliable way to determine the formula of an ionic compound by exchanging the magnitude of charges between the cation and anion to become subscripts.
• Example: Calcium Nitrate:
  • Ions: $Ca^{2+}$ and $NO_3^-$.
  • Swap the charges: The $2$ from Calcium moves to Nitrate; the $1$ from Nitrate moves to Calcium.
  • Drop the charges to subscripts: $Ca_1(NO_3)_2$.
  • Requirement: Parentheses must be placed around the molecular ion ($NO_3$) to indicate that two units of nitrate are present. Writing $NO_{32}$ is incorrect.
  • Final Formula: $Ca(NO_3)_2$.
• Example: Calcium Oxide:
  • Ions: $Ca^{2+}$ and $O^{2-}$.
  • Swap and Drop: $Ca_2O_2$.
  • Simplification: Always simplify to the lowest whole-number ratio. Result: $CaO$.

Structure and Properties of Ionic Compounds

• Giant Ionic Lattice: Ionic substances form large, regular repeating patterns of ions in 3D space rather than existing as isolated pairs of atoms.
• Sodium Chloride ($NaCl$):
  • Structure: Cubic shape.
  • Arrangement: Ions arrange themselves to be adjacent to as many oppositely charged particles as possible, providing immense robustness.
• Physical Properties:
  • Solubility: Generally soluble in water. Water is a polar solvent that pulls the lattice structure apart.
  • Electrical Conductivity:
  • Solid: Do not conduct electricity because ions are locked in position.
  • Molten/Aqueous: Conduct electricity because ions are free to move and carry charge.
  • Melting Point: High melting points due to the strength of the electrostatic forces between oppositely charged ions throughout the giant structure.
  • Brittleness: Brittle because a physical blow can shift layers of ions, bringing like-charges ($Cl^-$ next to $Cl^-$) together, causing mutual repulsion and fracturing.
• Evidence for Charged Particles (Electrolysis):
  • Example: Copper(II) Chromate ($CuCrO_4$) on wet filter paper.
  • Observation: Original solution is green.
  • Result: When electricity is applied, the blue $Cu^{2+}$ ions migrate to the negative electrode (cathode), and the yellow chromate ions ($CrO_4^{2-}$) migrate to the positive electrode (anode). This visually demonstrates the presence of charged particles.

Covalent Bonding

• Definition: The sharing of outer (valence) electrons between atoms to obtain a full shell.
• Context: Generally occurs between two non-metals.
• Mechanism: The bond is an electrostatic attraction between the positive nuclei of the atoms and the negative shared electrons.
• Types of Covalent Bonds:
  • Single Bond: One pair of shared electrons (represented by a single line).
  • Double Bond: Two pairs of shared electrons (represented by two lines).
  • Triple Bond: Three pairs of shared electrons (represented by three lines). Triple bonds are stronger and shorter than single or double bonds.
• Dative Covalent (Coordinate) Bonds:
  • Definition: A bond where one atom provides both electrons for the shared pair.
  • Notation: Represented by an arrow pointing from the donor atom to the acceptor.
  • Example: Ammonium Ion ($NH_4^+$). The lone pair on the Nitrogen in Ammonia ($NH_3$) is donated to a Hydrogen ion ($H^+$) which possesses no electrons.
  • Example: Carbon Monoxide ($CO$). Contains a double covalent bond and one dative covalent bond.
  • Example: Aluminum Chloride ($Al_2Cl_6$). Two $AlCl_3$ molecules join via coordinate bonds to form the more stable dimer $Al_2Cl_6$.

Bond Length and Bond Enthalpy

• Forces at Play:
  • Attrition between positive nuclei and shared negative electrons.
  • Repulsion between the two positive nuclei.
  • Repulsion between the electrons.
• Structural Balance: The bond length is the distance where these attractive and repulsive forces are balanced.
• Relationship: The greater the electron density between the atoms (e.g., in double or triple bonds), the stronger the attractive force.
  • Stronger Attraction → Atoms pulled closer → Shorter Bond → Higher Bond Enthalpy (stronger bond).
  • Triple bonds are shorter and have higher enthalpy than double bonds, which are shorter and stronger than single bonds.

Shapes of Molecules (VSEPR Theory)

• Principle: The shape is determined by the repulsion between electron pairs (bond pairs and lone pairs) surrounding a central atom. Electron pairs repel each other to be as far apart as possible in 3D space.
• Repulsion Hierarchy: Two lone pairs repel more than a lone pair and a bond pair, which in turn repel more than two bond pairs.
• Lone Pair Impact: Each lone pair typically reduces the bond angles between remaining bonds by approximately $2.5^{\circ}$.
• Common Shapes and Angles (No Lone Pairs):
  • Linear: $2$ bond pairs, $0$ lone pairs. Angle: $180^{\circ}$. Example: $BeCl_2$.
  • Trigonal Planar: $3$ bond pairs, $0$ lone pairs. Angle: $120^{\circ}$. Example: $BF_3$.
  • Tetrahedral: $4$ bond pairs, $0$ lone pairs. Angle: $109.5^{\circ}$. Example: $CH_4$.
  • Trigonal Bipyramidal: $5$ bond pairs, $0$ lone pairs. Angles: $90^{\circ}$ and $120^{\circ}$. Example: $PCl_5$.
  • Octahedral: $6$ bond pairs, $0$ lone pairs. Angle: $90^{\circ}$. Example: $SF_6$. (Octahedral refers to the $8$ faces of the 3D shape).
• Common Shapes and Angles (With Lone Pairs):
  • Pyramidal: $3$ bond pairs, $1$ lone pair. Angle: $107^{\circ}$ ($109.5 - 2.5$). Example: $NH_3$.
  • Bent / Non-linear: $2$ bond pairs, $2$ lone pairs. Angle: $104.5^{\circ}$ ($107 - 2.5$). Example: $H_2O$.
  • Trigonal Planar (Exception): $3$ bond pairs, $2$ lone pairs. Angle: $120^{\circ}$. Example: $ClF_3$. The lone pairs cancel each other's repulsion out.
  • Square Planar: $4$ bond pairs, $2$ lone pairs. Angle: $90^{\circ}$. Example: $XeF_4$.

Giant Covalent (Macromolecular) Structures

• Graphite:
  • Each carbon atom is bonded to $3$ others in hexagonal rings forming layers.
  • The $4^{\text{th}}$ electron is delocalized between layers.
  • Layers are held by weak forces, allowing them to slide (useful for pencils/lubricants).
  • Conducts electricity due to delocalized electrons.
  • High melting point and low density (relative to diamond).
• Diamond:
  • Each carbon atom is bonded to $4$ others in a rigid tetrahedral arrangement.
  • Extremely hard; used in cutting tools like circular saws and drill bits.
  • Conducts heat well but does not conduct electricity.
  • High melting point and insoluble.
• Silicon Dioxide ($SiO_2$): Also known as sand; has a similar giant covalent structure to diamond and shares similar properties.

Orbital Hybridization

• Orbital Types: $s$ orbitals (spherical) and $p$ orbitals (figure-of-eight; $p_x, p_y, p_z$).
• Hybridization Concept: The mixing of atomic orbitals to form new hybrid orbitals for bonding.
• $sp^3$ Hybridization (Example: methane, $CH_4$):
  • Carbon valence configuration: $2s^2 2p_x^1 2p_y^1 2p_z^0$.
  • To form $4$ bonds, an electron from the $2s$ orbital migrates to the empty $2p_z$ orbital (Excitation).
  • Excitation energy cost: $+404\,kJ\,mol^{-1}$.
  • The $s$ and three $p$ orbitals merge to form four equivalent $sp^3$ hybrid orbitals.
  • Energy released by forming $4$ $CH$ bonds: $-1648\,kJ\,mol^{-1}$.
  • Net energy change: $-1244\,kJ\,mol^{-1}$ (compared to only $-824\,kJ\,mol^{-1}$ if only $2$ bonds were formed). This explains why Carbon forms $4$ bonds rather than $2$.
• Sigma ($\sigma$) and Pi ($\pi$) Bonds:
  • Sigma (\sigma) Bonds: Formed by the end-to-end overlap of orbitals. All single bonds are sigma bonds.
  • Pi (\pi) Bonds: Formed by the sideways overlap of unhybridized $p$ orbitals. Found in double and triple bonds.
• $sp^2$ Hybridization (Example: Alkenes and Benzene):
  • The $s$ orbital merges with only two $p$ orbitals, creating three $sp^2$ hybrid orbitals at $120^{\circ}$ (trigonal planar).
  • One unhybridized $p$ orbital remains perpendicular ($90^{\circ}$) to the plane.
  • Side-on overlap of these unhybridized $p$ orbitals in adjacent atoms forms the pi bond.
  • A double bond consists of one sigma bond and one pi bond.
• $sp$ Hybridization: Mixing of one $s$ and one $p$ orbital.

Electronegativity and Bond Polarity

• Definition: Electronegativity is the ability of an atom to attract electrons toward itself within a covalent bond.
• Pauling Scale: Quantifies electronegativity. Fluorine is the most electronegative element with a value of $4.0$.
• Periodic Trend: Electronegativity increases moving up and to the right of the periodic table (excluding noble gases).
• Polar Bonds: Formed when there is a significant difference in electronegativity between two atoms in a covalent bond.
  • Result: Charge separation indicated by $\delta+$ (partial positive) and $\delta-$ (partial negative).
  • Example: $HCl$. Chlorine is more electronegative, so it is $\delta-$ while Hydrogen is $\delta+$.
• Non-polar Bonds: Occur between atoms with the same or very similar electronegativity (e.g., $Cl-Cl$ or hydrocarbons like methane).
• Polar Molecules: A molecule can have polar bonds but be non-polar overall if it is symmetrical (e.g., Carbon Dioxide, $CO_2$). The dipoles cancel out.

Intermolecular Forces

• Definition: Forces that exist between molecules. These are significantly weaker than chemical bonds.
• Types (in order of increasing strength):
  1. van der Waals Forces (Induced Dipole-Dipole):
  • Weakest force, existing in all molecules with electrons.
  • Caused by temporary, instantaneous dipoles as electrons move, which then induce a dipole in a neighboring molecule.
  1. Permanent Dipole-Dipole Forces:
  • Occur between polar molecules (like $HCl$).
  • The $\delta-$ end of one molecule is attracted to the $\delta+$ end of another.
  • Molecules with these forces also possess van der Waals forces.
  1. Hydrogen Bonding:
  • The strongest intermolecular force.
  • Occurs specifically when Hydrogen is covalently bonded to the three most electronegative elements: Nitrogen ($N$), Oxygen ($O$), or Fluorine ($F$).
  • Represents an interaction between the $\delta+$ Hydrogen and a lone pair on the $N, O,$ or $F$ of a neighboring molecule.
  • Water ($H_2O$) is a classic example and possesses all three types of intermolecular forces.

Metallic Bonding

• Definition: The electrostatic attraction between positive metal ions and a "sea" of delocalized electrons.
• Formation: Metals donate their valence electrons to a shared delocalized system.
• Melting Points: Guided by the number of delocalized electrons.
  • Example: Magnesium ($Mg$) has a higher melting point than Sodium ($Na$) because Magnesium can donate two electrons per atom, leading to stronger attractions.
• Properties:
  • Thermal/Electrical Conductors: Delocalized electrons are free to move and carry kinetic energy or electrical charge.
  • Malleability/Ductility: Metal layers can slide over each other when hammered or pulled; the delocalized electrons shift to maintain the attractive force and keep the structure together.
  • Insolubility: Metals are generally insoluble in water because the metallic bonds are too strong for water to break.

Summary of Substance Types

• Polarity and Solubility: "Like dissolves like." Polar molecules (like ammonia) dissolve well in polar solvents (like water), while non-polar molecules (like cooking oil/hydrocarbons) do not.