Chemistry - Chemical Bonding

Ionic Bonding

Formation

Ionic bonds form when metals (which lose electrons to form cations) react with non-metals (which gain electrons to form anions). The electrostatic attraction between oppositely charged ions constitutes the ionic bond.

Electron Transfer

Example: formation of sodium chloride

$\mathrm{Na} \to \mathrm{Na^+} + e^-$

$\mathrm{Cl} + e^- \to \mathrm{Cl^-}$

$\mathrm{Na^+} + \mathrm{Cl^-} \to \mathrm{NaCl}$

Giant Ionic Lattice

Ionic compounds form a regular three-dimensional lattice. Each ion is surrounded by ions of opposite charge. There are no discrete molecules in an ionic solid.

Properties of Ionic Compounds

Property	Explanation
High melting/boiling points	Strong electrostatic forces throughout the lattice
Conduct when molten or dissolved	Ions are free to move and carry charge
Do not conduct when solid	Ions are fixed in position
Soluble in polar solvents	Polar solvent molecules attract and separate ions
Brittle	Shifting layers brings like charges together

Worked Example 1

Write the formula of magnesium oxide, showing the ionic charges.

Solution

Magnesium ($Z = 12$): loses 2 electrons to form $\mathrm{Mg^{2+}}$.

Oxygen ($Z = 8$): gains 2 electrons to form $\mathrm{O^{2-}}$.

$\mathrm{Mg^{2+}} + \mathrm{O^{2-}} \to \mathrm{MgO}$

The charges balance (one $\mathrm{Mg^{2+}}$ per $\mathrm{O^{2-}}$).

Worked Example 2

Write the formula of aluminium oxide.

Solution

Aluminium loses 3 electrons: $\mathrm{Al^{3+}}$. Oxygen gains 2 electrons: $\mathrm{O^{2-}}$.

To balance charges: LCM of 2 and 3 is 6. Need $2 \times \mathrm{Al^{3+}}$ and $3 \times \mathrm{O^{2-}}$.

Formula: $\mathrm{Al_2O_3}$

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Covalent Bonding

Formation

A covalent bond forms when two non-metal atoms share one or more pairs of electrons so that each atom achieves a noble gas electron configuration.

Single, Double, and Triple Bonds

Bond Type	Shared Pairs	Example
Single	1	$\mathrm{H - H}$
Double	2	$\mathrm{O = O}$
Triple	3	$\mathrm{N \equiv N}$

Electron Dot Diagrams

Electron dot diagrams (Lewis structures) show valence electrons as dots. Shared pairs represent covalent bonds.

Example: Water ($\mathrm{H_2O}$)

Oxygen has 6 valence electrons. Each hydrogen has 1 valence electron. Oxygen shares one electron pair with each hydrogen:

$\mathrm{H} \colon \ddot{\mathrm{O}} \colon \mathrm{H}$

(Each hydrogen shares one pair with oxygen; oxygen has two lone pairs.)

Example: Carbon dioxide ($\mathrm{CO_2}$)

$\mathrm{O} \doubleequal \mathrm{C} \doubleequal \mathrm{O}$

Each oxygen shares two electron pairs with carbon (two double bonds).

Example: Ammonia ($\mathrm{NH_3}$)

Nitrogen has 5 valence electrons and shares one pair with each of three hydrogens. Nitrogen has one lone pair.

Properties of Covalent Compounds (Simple Molecular)

Property	Explanation
Low melting/boiling points	Weak intermolecular forces between molecules
Do not conduct electricity	No free ions or mobile electrons
Soluble in non-polar solvents	Similar intermolecular forces
Many are gases or liquids at rtp	Weak van der Waals forces easily overcome

Worked Example 3

Draw the electron dot diagram for hydrogen chloride ($\mathrm{HCl}$).

Solution

Hydrogen contributes 1 electron, chlorine contributes 7. They share one pair:

$\mathrm{H}\colon\ddot{\mathrm{Cl}}$

Chlorine has three lone pairs.

Dative Covalent (Coordinate) Bonds

A dative covalent bond is a covalent bond in which both electrons in the shared pair come from the same atom.

Example: Ammonium ion ($\mathrm{NH_4^+}$)

The nitrogen in $\mathrm{NH_3}$ donates its lone pair to form a dative covalent bond with $\mathrm{H^+}$:

$\mathrm{NH_3} + \mathrm{H^+} \to \mathrm{NH_4^+}$

All four N-H bonds in $\mathrm{NH_4^+}$ are equivalent once formed.

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Metallic Bonding

Formation

In metallic bonding, metal atoms release their valence electrons to form a "sea" of delocalised electrons. The positive metal ions are held together by electrostatic attraction to the delocalised electrons.

Properties of Metals

Property	Explanation
High melting/boiling points	Strong metallic bonding (more delocalised electrons = stronger bond)
Good electrical conductivity	Delocalised electrons carry charge freely
Good thermal conductivity	Delocalised electrons transfer kinetic energy
Malleable and ductile	Layers of positive ions can slide past each other
Lustrous	Delocalised electrons absorb and re-emit light

Strength of Metallic Bonding

The strength of metallic bonding increases with:

• More delocalised electrons (e.g., $\mathrm{Al}$ is harder than $\mathrm{Na}$)
• Smaller ionic radius (e.g., $\mathrm{Mg}$ is harder than $\mathrm{Ba}$)
• Higher charge on ions (e.g., $\mathrm{Mg^{2+}}$ vs. $\mathrm{Na^+}$)

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Intermolecular Forces

Intermolecular forces are weaker than intramolecular bonds (ionic, covalent, metallic). They act between molecules and determine many physical properties.

Van der Waals Forces (London Dispersion Forces)

Van der Waals forces arise from instantaneous dipole-induced dipole interactions caused by fluctuations in the electron cloud.

• Present between all molecules
• Increase with the number of electrons (and hence molecular size)
• Increase with molecular surface area (straight-chain isomers have stronger forces than branched)
• Weak individually, but significant in large molecules

Dipole-Dipole Forces

Polar molecules have permanent dipoles. The positive end of one molecule attracts the negative end of another. Dipole-dipole forces are stronger than van der Waals forces between molecules of similar size.

Hydrogen Bonding

Hydrogen bonding is a special, strong type of dipole-dipole interaction that occurs when:

1. A hydrogen atom is covalently bonded to a highly electronegative atom (N, O, or F).
2. The hydrogen atom interacts with a lone pair on another N, O, or F atom.

Examples:

• Water ($\mathrm{H_2O}$): extensive hydrogen bonding gives water its unusually high boiling point
• Ammonia ($\mathrm{NH_3}$)
• Hydrogen fluoride ($\mathrm{HF}$)
• DNA base pairing (adenine-thymine, guanine-cytosine)

Worked Example 4

Explain why the boiling point of water ($100^\circ\mathrm{C}$) is much higher than that of hydrogen sulphide ($-60^\circ\mathrm{C}$), despite $\mathrm{H_2S}$ having a larger molar mass.

Solution

Water molecules form strong hydrogen bonds between the hydrogen of one molecule and the lone pairs on the oxygen of another. $\mathrm{H_2S}$ molecules cannot form hydrogen bonds because sulphur is not sufficiently electronegative. $\mathrm{H_2S}$ only has weaker van der Waals and dipole-dipole forces. The strong hydrogen bonding in water requires much more energy to overcome, giving water its much higher boiling point.

Worked Example 5

Explain why butane ($\mathrm{C_4H_{10}}$) has a higher boiling point than 2-methylpropane ($\mathrm{C_4H_{10}}$, an isomer).

Solution

Both have the same molecular formula and number of electrons, so the strength of van der Waals forces depends on surface area. Butane is a straight-chain molecule with a larger surface area for contact between molecules, leading to stronger van der Waals forces. 2-methylpropane is more compact (branched), giving a smaller surface area and weaker van der Waals forces.

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Types of Structures

Giant Ionic Structure

• Example: $\mathrm{NaCl}$, $\mathrm{MgO}$
• Regular lattice of oppositely charged ions
• High melting point, conducts when molten/dissolved

Simple Molecular Structure

• Example: $\mathrm{H_2O}$, $\mathrm{CO_2}$, $\mathrm{CH_4}$, $\mathrm{I_2}$
• Weak intermolecular forces between molecules
• Low melting/boiling point, does not conduct electricity
• $\mathrm{CO_2}$ sublimes because the intermolecular forces are very weak

Giant Covalent (Macromolecular) Structure

Atoms are covalently bonded in a continuous giant lattice.

Substance	Structure	Properties
Diamond	Each C bonded to 4 others (tetrahedral)	Hardest natural substance, very high melting point, insulator
Graphite	Layers of hexagonal C rings; weak van der Waals between layers	Soft and slippery, conducts electricity (delocalised electrons), high melting point
Silicon(IV) oxide ($\mathrm{SiO_2}$)	Each Si bonded to 4 O; each O bonded to 2 Si	Hard, high melting point, insulator

Metallic Structure

• Regular lattice of positive ions in a sea of delocalised electrons
• Variable properties depending on the strength of metallic bonding

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Bond Polarity and Molecular Polarity

A covalent bond is polar if the bonded atoms have different electronegativities. The more electronegative atom carries a partial negative charge ($\delta^-$), and the less electronegative atom carries a partial positive charge ($\delta^+$).

Polar vs. Non-Polar Molecules

A molecule is polar if it has polar bonds and an asymmetric shape (so the bond dipoles do not cancel out).

Molecule	Polar Bonds?	Shape	Polar Molecule?
$\mathrm{HCl}$	Yes	Linear (asymmetric)	Yes
$\mathrm{CO_2}$	Yes	Linear (symmetric)	No (dipoles cancel)
$\mathrm{H_2O}$	Yes	Bent (asymmetric)	Yes
$\mathrm{CH_4}$	Yes	Tetrahedral (symmetric)	No (dipoles cancel)
$\mathrm{NH_3}$	Yes	Trigonal pyramidal (asymmetric)	Yes

Worked Example 6

Explain why $\mathrm{CCl_4}$ is a non-polar molecule despite having polar bonds.

Solution

Each $\mathrm{C - Cl}$ bond is polar because chlorine is more electronegative than carbon. However, $\mathrm{CCl_4}$ has a tetrahedral molecular geometry. The four bond dipoles are arranged symmetrically and cancel each other out completely. The resultant dipole moment is zero, so the molecule is non-polar.

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Common Pitfalls

• Confusing intramolecular forces (covalent bonds within a molecule) with intermolecular forces (forces between molecules). Covalent bonds are much stronger.
• Assuming ionic compounds always have the formula $AB$. The formula depends on the charges of the ions (e.g., $\mathrm{CaCl_2}$, $\mathrm{Al_2O_3}$).
• Thinking that covalent bonds are always weaker than ionic bonds. Giant covalent structures like diamond have extremely strong covalent bonds throughout.
• Forgetting that graphite conducts electricity because of delocalised electrons between layers, even though it is a giant covalent structure.
• Assuming hydrogen bonding can occur between any molecule containing hydrogen. It requires hydrogen bonded to N, O, or F.
• Confusing bond polarity with molecular polarity. A molecule with polar bonds can be non-polar if the shape is symmetric.

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Summary Table

Concept	Key Point
Ionic bonding	Electrostatic attraction between cations and anions
Covalent bonding	Sharing of electron pairs between non-metals
Metallic bonding	Delocalised electrons and positive ions
Van der Waals	Weak forces between all molecules; increase with size
Hydrogen bonding	Strong dipole-dipole; requires H bonded to N, O, or F
Giant ionic	High melting point, conducts when molten
Simple molecular	Low melting point, does not conduct
Giant covalent	Very high melting point; diamond insulates, graphite conducts

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Problem Set

Problem 1: Draw the electron dot diagram for the formation of magnesium chloride ($\mathrm{MgCl_2}$).

If you get this wrong, revise: Ionic Bonding — Electron Transfer

Solution

$\mathrm{Mg} \to \mathrm{Mg^{2+}} + 2e^-$

$2\mathrm{Cl} + 2e^- \to 2\mathrm{Cl^-}$

Each chlorine atom gains one electron to achieve a full outer shell. The magnesium ion transfers two electrons, one to each chlorine atom.

Problem 2: Explain why diamond has a much higher melting point than iodine.

If you get this wrong, revise: Types of Structures — Giant Covalent vs Simple Molecular

Solution

Diamond has a giant covalent structure with strong covalent bonds throughout the entire lattice. A large amount of energy is required to break all these bonds. Iodine has a simple molecular structure with weak van der Waals forces between $\mathrm{I_2}$ molecules, which require relatively little energy to overcome.

Problem 3: Explain why sodium chloride conducts electricity when molten but not when solid.

If you get this wrong, revise: Properties of Ionic Compounds

Solution

In solid $\mathrm{NaCl}$, the ions are fixed in the lattice and cannot move, so they cannot carry charge. When molten, the ions are free to move towards the electrodes, allowing conduction.

Problem 4: Explain the difference in boiling points between $\mathrm{HF}$ ($20^\circ\mathrm{C}$) and $\mathrm{HCl}$ ($-85^\circ\mathrm{C}$).

If you get this wrong, revise: Hydrogen Bonding

Solution

$\mathrm{HF}$ can form strong hydrogen bonds between molecules (H bonded to highly electronegative F), requiring significant energy to overcome. $\mathrm{HCl}$ has polar bonds but chlorine is less electronegative than fluorine and cannot form hydrogen bonds; $\mathrm{HCl}$ only has weaker dipole-dipole interactions and van der Waals forces.

Problem 5: Is $\mathrm{BF_3}$ a polar molecule? Explain your answer.

If you get this wrong, revise: Bond Polarity and Molecular Polarity

Solution

No, $\mathrm{BF_3}$ is non-polar. Although each $\mathrm{B - F}$ bond is polar (F is more electronegative than B), the molecule has a trigonal planar geometry. The three bond dipoles are arranged symmetrically at $120^\circ$ and cancel each other out completely.

Problem 6: Explain why graphite is used as a lubricant while diamond is used as an abrasive.

If you get this wrong, revise: Giant Covalent Structures

Solution

Graphite has a layered structure with weak van der Waals forces between layers. The layers can slide over each other easily, making graphite a good lubricant. Diamond has a rigid tetrahedral network of strong covalent bonds in all three dimensions, making it the hardest known natural substance and an excellent abrasive.

Problem 7: Predict the shape and bond angle of the $\mathrm{NH_4^+}$ ion.

If you get this wrong, revise: Covalent Bonding — Dative Bonds and VSEPR

Solution

Nitrogen has 5 valence electrons. Three form bonds with hydrogen, and the lone pair forms a dative bond with $\mathrm{H^+}$, giving four bonding pairs and zero lone pairs around nitrogen.

Total electron pairs = 4 (all bonding). Molecular shape: tetrahedral. Bond angle: $109.5^\circ$.

Problem 8: Explain why magnesium oxide has a higher melting point than sodium chloride.

If you get this wrong, revise: Ionic Bonding — Properties and Ionic Radii

Solution

Both compounds have giant ionic lattices. The strength of the ionic bond depends on the charge and size of the ions. $\mathrm{Mg^{2+}}$ and $\mathrm{O^{2-}}$ have higher charges ($+2$ and $-2$) than $\mathrm{Na^+}$ and $\mathrm{Cl^-}$ ($+1$ and $-1$). Higher ionic charges produce stronger electrostatic attraction according to Coulomb's law ($F \propto q_1 \times q_2$). Additionally, $\mathrm{Mg^{2+}}$ and $\mathrm{O^{2-}}$ are smaller ions, allowing them to get closer together, further increasing the electrostatic attraction.

Problem 9: Explain why $\mathrm{H_2O}$ is a liquid at room temperature while $\mathrm{H_2S}$ is a gas, even though both are in Group 16 hydrides.

If you get this wrong, revise: Hydrogen Bonding

Solution

Water can form extensive hydrogen bonds between molecules because hydrogen is bonded to oxygen, which is highly electronegative (EN $= 3.5$). Each water molecule can form up to four hydrogen bonds, creating a network that requires significant energy to break.

$\mathrm{H_2S}$ cannot form hydrogen bonds because sulphur (EN $= 2.1$) is not electronegative enough. $\mathrm{H_2S}$ molecules only have weaker dipole-dipole interactions and van der Waals forces between them. These are easily overcome at room temperature, making $\mathrm{H_2S}$ a gas.

Problem 10: Write the formula of calcium fluoride, showing the ionic charges, and explain why ionic compounds are generally brittle.

If you get this wrong, revise: Ionic Bonding — Formation and Properties

Solution

Calcium ($Z = 20$, Group 2): $\mathrm{Ca^{2+}}$. Fluorine ($Z = 9$, Group 17): $\mathrm{F^-}$.

To balance charges: $1 \times (+2) = 2 \times (-1)$. Formula: $\mathrm{CaF_2}$.

Ionic compounds are brittle because if a force is applied, layers of ions shift so that ions of the same charge come adjacent to each other. The resulting electrostatic repulsion between like charges causes the crystal to crack or shatter.