Section B

Intermolecular forces

Covalent bonds join atoms together inside a molecule and are the intramolecular forces.

Inter are the forces that bond molecules together.

Intramolecular forces

Covalent bonds fix the atom sequence; non covalent contacts stabilise some shapes. 

Conformation

Electronegativity 

Electronegativity refers to how strong the bonds are shared in an atom. 

The symbols

𝛿+ and 𝛿- mean partial positive and partial negative. 

This separation of charge is called a bond dipole

Bond dipole

A separation of partial positive and partial negative charge within a bond. 

Polarity

the uneven distribution of electric charge across atoms in a bond or molecule. The uneven shared electrons are known as a polar bond

A polar molecule

A polar molecule is a molecule that has an uneven distribution of electrical charge, 

Dipole

A dipole is negative end of one molecule attracts the positive end of another molecule

Causes of polarity

A polar bond does NOT automatically mean a polar molecule 

A bent molecule will have dipoles that do not point in opposite directions; therefore, they do not cancel.

This makes molecules like H2O polar. But CO2 non polar.

Types of intermolecular forces

Most molecules have MORE THAN ONE intermolecular force 

Dispersion forces 

Also known as London dispersion forces.

Movement of electrons

• At one instant, the electrons in an atom or molecule can be slightly more concentrated on one side.

• That temporary dipole affects a neighbouring particle

• Creates a temporary dipole there too.

It’s the movement of electrons, not because of an unequal distribution of electrons

Bigger molecules usually have stronger dispersion

More electrons mean a larger electron cloud, that can be distorted more easily and they are more polarisable.

Therefore, it has higher intermolecular forces.

Higher boiling point.

Dipole-dipole forces 

Permanently polar molecules are partially positive or negative.

Polar molecules can orient themselves to bond together.

For example; HCl.

Dipole-dipole attractions can co-exist with dispersion attractions. For example, HCl.

Hydrogen bonding

Not covalent; attraction between partial charges and lone electrons. 

Hydrogen (the donor) bonded to N, O and F (the acceptor) 

Represented by a dotted line compared to the covalent, straight line. 

Significance of hydrogen bonding 

Boiling point 

• Water molecules form lots of hydrogen bonds with one another. 

• Overcoming many of those intermolecular hydrogen bonds lead to high boiling point. 

Frozen and density 

• In ice, water molecules arrange themselves into a relatively open hydrogen-bonded structure. It is less dense than water, so ice floats. 

Ion–dipole forces 

Ions interacting with dipoles of non-charged particles creates ion-dipole interactions / hydration of ions.

How intermolecular forces affect boiling point 

If a molecule like dimethyl ether, is missing an OH group relative to ethanol, then it’s hydrogen bonds are weaker and it will boil faster. 

States of matter

Molecular Solids	Intermolecular solids
Ionic Solids	Ionic bonding
Metallic Solids	Metal ion cores and delocalised electrons.
Covalent network solids	Have covalent bonds extending throughout the structure.

Determinants of states of matter

Intermolecular attractions:

• Attractions favour particles staying close together. Hydrogen bonds, dispersion etc.

Temperature 

• Higher temperature means greater average particle kinetic energy, helping thermal motion disrupt those attractions. 

Pressure 

• Higher pressure generally favours the denser phase. 

Phase changes

Diagonal parts mean that temperature is increasing. 

Flat parts mean that phase change is occurring 

Along the diagram: endothermic reactions take place

The vice versa applies

Endo & exothermic reaction 

A reaction that requires energy is exothermic. 

Such as freezing. Freezing conditions can release heat as the water freezes 

A reaction that releases energy is endothermic. 

Such as evaporation.  It absorbs heat from surroundings. 

Evaporation

Evaporation, molecules at the surface gain enough kinetic energy to overcome forces of attraction holding it to neighbouring molecules. 

Boiling

Vapour pressure= external pressure = boiling

Boiling happens within a liquid, where water molecules gain kinetic energy on average, causing them to escape into the gas phase. The liquid's vapour pressure increases, equalling the external pressure causing vapour bubbles to form and grow inside the liquid, rising towards the surface. 

Triple point:

Triple Point: the pressure where three phases co exist. 

Critical Point: End of the liquid–gas line.  

• Liquid density decreases, mainly because temperature increases. 

• Gas density increases, as the vapour becomes denser at higher pressure and temperature. 

• Eventually, their densities become equal, and the liquid–gas distinction disappears. 

Phase diagram with pressure and temperature

Sublimation 

An endothermic phase change in which a solid changes directly into a gas without passing through the liquid state. 

This occurs because the surrounding pressure is below the triple-point pressure, allowing a solid to turn into a gas. 

Freeze drying

In an atmosphere, below the triple pressure point, ice is sublimed and a porous solid remains.

It works because below the triple pressure point, liquid water cannot exist as a stable phase; pressure is reduced to the point that the boiling temperature is so low that water molecules escape into gas phase.

Dissolution

Like dissolves like: similar intermolecular forces usually mix well; entropy also favours mixing.

Polarity dissolution

Polar substances often dissolve well in polar solvents. 

Nonpolar substances often dissolve well in nonpolar solvents. 

Oil non polarity

Oil is non polar and therefore does not dissolve well. Emulsifiers contain polar and non polar regions that bridge together the water and the oil, allowing for dispersion.  

Entropy

Entropy: mixing can give particles more possible arrangements.

Ion dipole solubility

The polar water molecules are attracted to the ions.

Na+ ions are attracted to the negative oxygen molecules.

H- ions are attracted to the positive hydrogen ions.

These electrostatic forces dissolve the molecule in water.

Hydrogen bonding solubility

Sugar contains many OH groups that form with hydrogen bonds with water molecules. The partially positive hydrogen in sugar's –OH group can form a hydrogen bond with the lone pairs of electrons in oxygen.

Acetone has oxygen groups that bond to hydrogen. Water has partially positive hydrogen atoms (δ+), which are attracted to the partially negative oxygen (δ−) in acetone.

Model of Gases

Particles move randomly and continuously, travelling in straight lines between collisions.

Their own volume is negligible compared with the total gas volume.

There are no forces between particles between collisions.

Collisions are elastic: the total kinetic energy of the colliding particles is conserved

Average kinetic energy is proportional to temperature in kelvin.

When it fails

Particle volume and attractions are the two assumptions that fail in real gases.

The model becomes less accurate when:

Negligible particle volume

- Pressure is high: particles are crowded, so their own volume matters.

Temperature is low

- Temperature is low, particularly near condensation: attractions between particles become important (Fails at low T, near condensation).

Pressure 

Collisions with walls