Bonding, Structure, and Gases

Most ionic, metallic, and covalent compounds form crystalline lattices.
Ions, atoms, or molecules are arranged in a regular, repeating pattern within these lattices.

An ionic bond is the electrostatic force of attraction between a positively charged metal ion (cation) and a negatively charged non-metal ion (anion).
Metals become positively charged by transferring electrons to non-metals, which then become negatively charged.
Ionic compounds arrange themselves into giant ionic lattices (also called giant ionic structures).
The specific type of lattice structure depends on the relative sizes of the positive and negative ions, which are arranged in an alternating pattern.
Examples of ionic lattices with a cubic structure include Magnesium Oxide ( $MgO$ ) and Sodium Chloride ( $NaCl$ ).
In general ionic lattices, negative ions (anions) are typically larger than the positive ions (cations).
The ionic lattices for $NaCl$ and $MgO$ are similar due to the 1:1 ratio of cations to anions.

Covalent bonds involve the sharing of electrons between non-metal atoms.
Covalent compounds can form simple molecular lattices or giant molecular lattices.
- Simple molecular lattices: Examples include iodine ( $I<em>2$ ), buckminsterfullerene ( $C</em>{60}$ ), and ice ( $H_2O$ ).
- Giant molecular lattices: Examples include silicon(IV) oxide ( $SiO_2$ ), graphite, and diamond.
Different simple molecular lattices (e.g., ice, buckminsterfullerene, iodine) exhibit different structures and intermolecular forces.
Giant molecular lattices generally have higher melting and boiling points due to the significant energy needed to overcome intramolecular and/or intermolecular forces.

Metals form giant metallic lattices where metal ions are surrounded by a 'sea' of delocalized electrons.
Metal ions are often arranged in hexagonal layers or cubic arrangements.
In copper ( $Cu$ ), the cations are arranged in regular layers.

Different types of bonding and structure influence physical properties like melting and boiling points, electrical conductivity, and solubility.

Ionic compounds are strong due to strong electrostatic forces that hold the ions together.
They tend to be brittle because ionic crystals can easily split apart under stress.
Ionic compounds have high melting and boiling points because strong electrostatic forces between ions act in all directions, maintaining a strong lattice.
Melting and boiling points increase with the charge density of the ions due to greater electrostatic attraction. For example, $Mg^{2+}O^{2-}$ has a higher melting point than $Na^+Cl^-$ .
Ionic compounds are often soluble in water because they can form ion-dipole bonds.
Ionic compounds conduct electricity only when molten or in solution, as ions can move freely. In the solid state, ions are fixed in position.

Metallic compounds are malleable; metal layers can slide when a force is applied, and metallic bonds reform.
The attractive forces between metal ions and delocalized electrons act in all directions.
Metallic compounds are strong and hard due to strong attractive forces between metal ions and delocalized electrons.
Metals have high melting and boiling points because of these strong attractive forces.
A greater number of delocalized electrons and smaller cation size result in greater attractive force and higher melting/boiling points.
Pure metals are insoluble in water.
Metals conduct electricity in solid and liquid states because mobile electrons can move freely.

Simple covalent lattices have low melting and boiling points because of weak intermolecular forces.
- Little energy is needed to break the lattice.
Most compounds are insoluble in water unless they are polar (like $HCl$ ) or can form hydrogen bonds (like $NH_3$ ).
They do not conduct electricity in solid or liquid states due to the absence of charged particles.
Some simple covalent compounds conduct electricity in solution (e.g., $HCl$ forms $H^+$ and $Cl^-$ ions).

Giant covalent lattices have high melting and boiling points because of a large number of covalent bonds linking the entire structure.
A lot of energy is required to break the lattice.
The compounds can vary in hardness.
- Graphite is soft due to weak forces between carbon layers.
- Diamond and silicon(IV) oxide are hard due to their 3D network of strong covalent bonds.
Most compounds are insoluble in water.
Most compounds do not conduct electricity, but some exceptions exist.
- Graphite has delocalized electrons between carbon layers that can move when a voltage is applied.
- Diamond and silicon(IV) oxide do not conduct electricity because all outer electrons are involved in covalent bonds.

Characteristic	Giant Ionic	Giant Metallic	Simple Covalent	Giant Covalent
Melting/Boiling Points	High	Moderately high to high	Low	Very high
Electrical Conductivity	Only when molten/solution	When solid or liquid	Do not conduct electricity	Do not conduct electricity (except graphite)
Solubility	Soluble	Insoluble (some may react)	Usually insoluble unless they are polar	Insoluble
Hardness	Hard, brittle	Hard, malleable	Soft	Very hard (diamond/SiO2) or soft (graphite)
Physical State	Solid	Solid	Solid, liquid, gas	Solid
Forces	Electrostatic attraction	Delocalised sea of electrons attracting positive ions	Weak intermolecular forces between molecules	Electrons in covalent bonds between atoms
Particles	Ions	Positive ions in a sea of electrons	Small molecules	Atoms
Example	$NaCl$	Copper	$Br_2$	Graphite, silicon(IV) oxide

Problem: Given the physical properties of substances X, Y, and Z, identify their structure.
Solution:
- X (Melting Point: 839°C, Conductivity: Good when molten, Solubility: Soluble) - likely a giant ionic structure.
- Y (Melting Point: 95°C, Conductivity: Very poor, Solubility: Almost insoluble) - likely a simple molecular structure.
- Z (Melting Point: 1389°C, Conductivity: Good, Solubility: Insoluble) - likely a giant metallic structure.

Decreasing volume (at constant temperature) increases pressure due to more frequent collisions.
Volume is inversely proportional to pressure at constant temperature.

Increasing temperature (at constant volume) increases pressure as molecules gain kinetic energy and collide more frequently.
Temperature is directly proportional to pressure at constant volume.

Gas molecules are constantly moving randomly.
Molecules have negligible volume.
No intermolecular forces exist.
Collisions are elastic (no kinetic energy loss).
Temperature is related to average kinetic energy.
Gases following these rules are called ideal gases; real gases approximate this.

Volume depends on pressure and temperature.
Heating a gas (at constant pressure) increases volume as particles gain kinetic energy and collide more frequently.
Volume is directly proportional to temperature.

Real gases deviate at high pressures and low temperatures because:
- Molecules are closer together.
- Intermolecular forces become significant.
- Molecular volume is no longer negligible.
Assumptions of no intermolecular forces and negligible volume break down.
Attractive forces reduce impact force, lowering measured pressure.