Bonding Types and Ionic Bonding

Think of ionic bonding like a magnet attraction between oppositely charged particles. When metals lose electrons, they become positive cations, whilst non-metals gain electrons to become negative anions. These oppositely charged ions attract each other strongly, forming a regular 3D structure called a lattice.

You'll need to memorise key compound ions like sulfate (SO₄²⁻), hydroxide (OH⁻), nitrate (NO₃⁻), carbonate (CO₃²⁻), and ammonium (NH₄⁺). These pop up constantly in exams, so knowing them saves precious time.

The periodic table is your best friend for predicting charges. Group 1 metals always form +1 ions, Group 2 form +2 ions, whilst Group 7 non-metals form -1 ions. Once you know the charges, writing formulas becomes straightforward - just balance the positive and negative charges.

Quick Tip: Remember that in ionic compounds, the total positive charge must equal the total negative charge!

Covalent and Dative Covalent Bonds

Covalent bonds form when atoms share electrons - it's like two friends sharing a pizza equally. You can represent single covalent bonds with a simple line (-), whilst double and triple bonds use multiple lines $= or ≡$.

Dative covalent bonds (also called coordinate bonds) are slightly different - here, one atom provides both electrons for sharing. Think of it as one friend buying the whole pizza but still sharing it equally. You draw these with an arrow (→) pointing from the electron donor to the acceptor.

Once a dative covalent bond forms, it behaves exactly like a regular covalent bond - the only difference is where the electrons originally came from. This concept appears frequently in complex ion questions, so understanding the arrow notation is crucial.

The key difference to remember: covalent bonds involve equal electron sharing, whilst dative bonds involve one-sided electron donation that still results in equal sharing.

Remember: The arrow in dative bonds shows the direction of electron donation, not electron movement!

Metallic Bonding

Picture metallic bonding as a "sea of electrons" surrounding positive metal ions. The electrons are delocalised, meaning they're free to move around the entire structure rather than being stuck between specific atoms. This creates strong electrostatic attraction in all directions.

Three factors determine metallic bond strength: smaller metal ions create stronger attraction, higher charges on ions increase the force, and more delocalised electrons strengthen the bonding. This explains why aluminium $3+ charge$ is stronger than sodium $1+ charge$.

You'll encounter four main crystal structures: ionic (like sodium chloride), metallic (like magnesium), macromolecular or giant covalent (like diamond), and molecular (like ice). Each type has distinct properties based on its bonding.

Remember the difference between intermolecular forces (between separate molecules) and intramolecular bonds (within one molecule). This distinction is vital for understanding why substances behave differently when heated or dissolved.

Key Point: Metallic bonding explains why metals conduct electricity - those delocalised electrons can carry electrical current!

Crystal Structures and Properties

Simple molecular crystals like iodine and ice have strong bonds within molecules but weak forces between them. That's why ice melts easily - you're breaking the weak intermolecular forces, not the strong covalent bonds within water molecules.

Giant covalent structures like diamond and graphite contain millions of atoms joined by strong covalent bonds. They need enormous energy to break apart, giving them extremely high melting points. Graphite conducts electricity because it has delocalised electrons between its layers.

Ionic crystals like salt form regular 3D lattices of alternating positive and negative ions. They don't conduct electricity as solids because the ions can't move, but they do when melted or dissolved in water as the ions become mobile.

Metallic crystals have some of the strongest bonding due to electrostatic forces acting in all directions. They're insoluble in water but conduct electricity brilliantly thanks to those mobile delocalised electrons.

Exam Tip: If asked about conductivity, always mention whether charged particles can move freely!

Molecular Shapes

Imagine electrons as negative charges that hate being near each other - this is the basis of VSEPR theory (Valence Shell Electron Pair Repulsion). Electron pairs arrange themselves as far apart as possible to minimise repulsion, which determines molecular shape.

Lone pairs $non-bonding electrons$ repel more strongly than bonding pairs. They push bonding pairs closer together, reducing bond angles by about 2-2.5 degrees. This explains why ammonia (NH₃) has a smaller bond angle than methane (CH₄).

The number of electron pairs around the central atom determines the basic shape: 2 pairs give linear, 3 pairs give trigonal planar, 4 pairs give tetrahedral, and so on. However, lone pairs affect the final molecular geometry.

You need to predict shapes and bond angles for molecules with up to six electron pairs. Start by counting total electron pairs, then consider how many are lone pairs to determine the final shape.

Memory Aid: Lone pairs are "bigger bullies" - they push bonding pairs around more than bonding pairs push each other!

Bond Polarity

Electronegativity measures how strongly an atom attracts electrons in a covalent bond. When atoms with different electronegativities bond, the electrons spend more time near the more electronegative atom, creating a polar covalent bond.

Three factors affect electronegativity: nuclear charge $more protons = stronger pull$, atomic radius (smaller atoms pull harder), and electron shielding $fewer electron shells = less interference$. This explains why fluorine is the most electronegative element.

Electronegativity increases across periods (more nuclear charge) and decreases down groups (larger atoms, more shielding). You can use partial charges $δ+ and δ-$ to show which end of a bond is slightly positive or negative.

Some molecules contain polar bonds but aren't polar overall - this happens when the molecule is symmetrical, so the polar bonds cancel each other out. Carbon dioxide (CO₂) is a perfect example of this.

Quick Check: A molecule is only polar if it has polar bonds AND an asymmetrical shape!