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INTERMOLECULAR FORCES
In a molecular substance (most solid and liquid), there must be forces between the particles or molecules that attract them to each other; otherwise, the molecules would move apart. These forces are called intermolecular forces. Intermolecular force: a weak force of attraction between molecules, ions, or atoms of noble gases.

INTRAMOLECULAR FORCES
In contrast to intermolecular forces, intramolecular forces are classified as ionic, metallic, or covalent bonds. Interatomic forces (also known as intra-molecular forces):  a bond that occurs between atoms within molecules.

NB: Interatomic forces are stronger than intermolecular forces.

It requires much less energy to evaporate a particular liquid than to break the bonds within its molecules to form separate atoms, ions, or new molecules or compounds.

For example, It takes 464 kJ/mol to break the H-O bonds within a water molecule and only 19 kJ/mol to weaken the forces between water molecules.

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1. induced dipole forces

2. dipole-dipole forces

3. ion-dipole forces

1. INDUCED DIPOLE FORCES

Induced dipole forces result when an ionic, polar or a non-polar substance induces a dipole in an atom or a molecule with no dipole (non-polar molecule). The three induced dipole forces are namely,  London or dispersion (also called induced dipole-induced dipole) forces, ion-induced dipole forces, and dipole-induced dipole forces.

1.1 London dispersion forces: temporary attractive forces that result when the electrons in two adjacent atoms or molecules occupy positions that make them form temporary or momentary dipoles. London forces are the attractive forces that cause non-polar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently

Because of the constant motion of the electrons, an atom or molecule can develop a temporary (instantaneous) dipole when its electrons are distributed asymmetrically about the nucleus, thus causing a distortion in the electron cloud density of the neighbouring atom or molecule (because electrons repel one another) which leads to an electrostatic attraction between the two atoms or molecules. Dispersion forces are present between any two molecules (even polar molecules) when they are almost touching.

NB: The London dispersion force is the weakest intermolecular force.

The strength of London forces is influenced by the size of the temporary dipoles, which is influenced by:
- the number of electrons
- the interacting surface of the molecule (e.g. long unbranched chain vs branched-chain).

Strength of London forces and molecular size

Larger and heavier atoms and molecules exhibit stronger dispersion forces than smaller and lighter ones.

In a larger atom or molecule, the valence electrons are, on average, farther from the nuclei than in a smaller atom or molecule. They are less tightly held and can move easily form temporary dipoles.

The ease with which the electron distribution around an atom or molecule can be distorted is called the polarizability.

Therefore, London dispersion forces tend to be stronger between molecules that are easily polarized and weaker between molecules that are not easily polarized.

Strength of London forces and molecular shape

The shapes of molecules also affect the magnitudes of dispersion forces between them.

At room temperature, neopentane (C₅H₁₂) is a gas whereas n-pentane (C₅H₁₂) is a liquid.

The somewhat cylindrical shape of n-pentane molecules allows them to come in contact with each other more effectively than the somewhat spherical neopentane molecules.

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1.2 I on-induced dipole force: a weak attraction that results when the approach of an ion induces a dipole in an atom or non-polar molecule by disturbing the arrangement of electrons in the non-polar species.

1.3 Dipole-induced dipole force:  a weak attraction that results when a polar molecule induces a dipole in an atom or non-polar molecule by disturbing the arrangement of electrons in the non-polar species.

2. DIPOLE-DIPOLE FORCES

Dipole-dipole force is an attractive force between the positive end of one polar molecule and the negative end of another polar molecule. They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).

Quiz: Both ICl and Br₂ have the same number of atoms and approximately the same molecular weight, but ICl is a solid whereas Br₂ is a liquid at 0oC. Why?

Answer: Intermolecular dipole-dipole attractions between ICl molecules are sufficient to cause them to form a solid at 0oC, whereas the intermolecular attractions between non-polar Br₂ molecules are not.

2.1 Hydrogen bonding
Hydrogen bonding exists between molecules in which hydrogen is bonded to:

a small atom;

of high electronegativity

with at least one lone pair of electrons

Hydrogen bonding is a special type of dipole-dipole attraction and forms between a hydrogen atom in one polar molecule and a small, highly electronegative atom (N, O, and F) in a neigbouring polar molecule. Hydrogen bonding is generally stronger than ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.

3. ION-DIPOLE FORCES
An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. They are most commonly found in solutions, especially important for solutions of ionic compounds in polar liquids. A positive ion (cation) attracts the partially negative end of a neutral polar molecule while a negative ion (anion) attracts the partially positive end of a neutral polar molecule.

NB: Ion-dipole attractions become stronger as either the charge on the ion increases or as the magnitude of the dipole of the polar molecule increases.

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Recap:
What type of forces exist in the mixtures of the following pairs:

1. HBr and H₂S

2. I₂ and NH₃

3. H₂O and HF

4. H₂O and CH₄

5. HCl and H₂O

6. H₂O and O₂

Dissolution of ionic compounds
When an ionic solid dissolves in a polar substance, ion-dipole forces are set up, preventing the particles of the ionic solid to from re-bonding again.

Dissolve: become or cause to become incorporated into a liquid so as to form a solution.
Dissolution: the act or process of dissolving.
Solute: a solid that is dissolved into a liquid
Solvent: the liquid in which a solute is dissolved to form a solution.

There are two types of dissolution processes, namely, ionisation and dissociation.

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Ionisation: the process in which molecular compounds dissolve in water to produce ions.

Most molecular compounds do not undergo ionization, except Acids. All acids produce hydrogen ions in solution. Some examples are as follows:

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The reactions as written in the left hand column of the table above are actually simplified versions of what really occurs. Evidence suggests that the hydrogen ion, H⁺, actually bonds to a water molecule (H₂O) to form the hydronium ion, H₃O⁺, as seen in the right hand column.

For our class, it won't make a difference if you write an acid ionization reactions as producing hydrogen ions (H⁺) or hydronium ions (H3O⁺). For our purposes they will represent the same thing. You should be comfortable using either; one will mean the same as the other.

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Dissociation of ionic compounds occurs when water molecules “pull apart” the ionic crystal. This occurs due to strong attractions between the polar ends of the water molecule and the positive and negative ions within the crystal. Water molecules then surround the positive cations and negative anions; this is called hydration.

Examples of dissociation equations:

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NB: It is worth noticing that the last three reactions shown above all produce hydroxide ions (OH−). These three compounds - NaOH, KOH, and Mg(OH)2 are all Arrhenius bases since they produce hydroxide ions in solution.

There are two important things to notice about writing dissociation equations:

1. Generally DO NOT include H₂O as a reactant. We know something has been dissolved in water when we see the (aq) notation.

2. Ion charges MUST BE included.

DISCUSSION
Beaker A: contains 500 ml of water
Beaker B: contains 500 ml of Sodium chloride (0.2 M)
Beaker C: contains 500 ml  of Sodium chloride (0.5 M)

If all beakers were to be heated continuously, which mixture would boil first? Explain your answer.

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Question and answer:
Four beakers, each containing 10 ml of liquid water, copper sulphate solution (1.0 M), liquid carbon dioxide, and hydrocloric acid (1.0 M)are to be boiled.

Which beaker will require more energy to boil?

Below are the reletave strengths of intermolecular forces in terms of energy required to overcome them:

1. Ionic bond: 400-4000 kJ/mol

2. Hydrogen bonding: 10-40 kJ/mol

3. Dipole-dipole forces: 2-25 kJ/mol

4. Van der Walls: 0.05-40 kJ/mol

States of matter
All substances exist in one of the three states; solid, liquid, or gas. The sate in which a substance exists depends on the following factors:

the kinetic energy of the particles; and

the intermolecular forces between the particles.

These different states of matter have different properties, which are illustrated in the table below.

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As another way that we to describe the properties of matter is the state (also called phase), physical properties of a substance depend partly on temperature and air pressure. At a particular temperature, a substance will exist as either a solid, liquid or gas.

For example, at the air pressure found at sea level, water exists as a liquid at temperatures between 0oC and 100oC. Above 100oC, water exists as a gas (water vapor). Below 0oC, water exists as a solid (ice). Different substances have a different range of temperatures at which they exist in each state. These differences explain why some substances are always solids at normal Earth temperatures, whereas others are always gases or liquids.

Changes in state

Adding energy to matter gives its particles the ability to resist some of the forces holding them together. Continuous addition of energy will leads to particles gaining energy, vibrating more strongly, and eventually having enough energy to overcome the attractive forces holding them together, thus the matter expanding and change of phase occurring. For example, heating a solid to its melting point gives its molecules enough energy to move, thus the solid melting and becoming liquid. Similarly, heating a liquid  to its boiling point gives its molecules enough energy to pull apart from one another so they no longer have contact. The liquid vaporizes and becomes water vapor.

Boiling point: the temperature where the liquid-gas phase change occurs.
Melting point: the temperature where the solid-liquid phase change occurs.

The temperature at which phase change occurs depends on the identity of the substance and the atmospheric pressure. Each substance has its own boiling and melting points that depend on the properties of the substance, including the strength of its intermolecular forces. The different processes of phase change are shown in the figure below.

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The effect of intermolecular forces on boiling point

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The effect of intermolecular forces on boiling point

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1. DENSITY
The density of a substance is the ratio of its mass and the volume its particles occupy. For example, the density of liquid water at 4oC is 1 g/cm³. This means that molecules in one gram of liquid water occupy a volume of 1 cm³.

Effect of intermolecular forces on density
Since the way in which particles in a substance are packed is dependent on intermolecular forces, the space the particles occupy depends on the strength of the intermolecular forces. Therefore, the density of a substance is also dependent on the strength of the intermolecular forces in the substance.

NB: since particles of different substances pack differently, volumes of the two substances will be different even if their masses are the same.

Effect of temperature on density
Since an increase in temperature results in the expansion of a substance, an increase in its temperature causes a decrease in its density.
2. SURFACE TENSION
Surface tension is the energy, or work, required to increase the surface area of a liquid due to intermolecular forces.

Since these intermolecular forces vary depending on the nature of the liquid (e.g. water vs. gasoline) or solutes in the liquid (e.g. surfactants like detergent), each solution exhibits differing surface tension properties.

In a sample of water, there are two types of molecules. Those that are on the outside (exterior), and those that are on the inside (interior). The interior molecules are attracted to all the molecules around them, while the exterior molecules are attracted to only the other surface molecules and to those below the surface. This makes it so that the energy state of the molecules on the interior is much lower than that of the molecules on the exterior. Because of this, the molecules try to maintain a minimum surface area, thus allowing more molecules to have a lower energy state. This is what creates what is referred to as surface tension. An illustration of this can be seen in figure below:

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The molecules at the surface of water experience a net attraction to other molecules in the liquid, which holds the surface of the bulk sample together. In contrast, those in the interior experience uniform attractive forces.

Surface tension of a substance can cause things to float which are more dense than the substance. The effect of surface tension is shown in the following figures:

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There are several other important concepts that are related to surface tension. The first of these is the idea of coercive and adhesive forces.

Coercive forces: the forces exerted between molecules holding them together. If cohesive forces are strong, a liquid tends to form droplets on a surface.

Adhesive forces: the forces between liquid molecules and a surface. If adhesive forces are strong, a liquid tends to spread across a surface.

So if the cohesive forces are stronger than the adhesive forces, the body of water will maintain its shape, but if the opposite is true then the liquid will be spread out, maximizing its surface area. Any substance that you can add to a liquid that allows a liquid to increase its surface area is called a wetting agent. A visual effect of cohesive and adhesive forces is evidenced in the concept of meniscus as shown in figure below:

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Figure (a) illustrates the shape of the meniscus and the relative height of a mercury column when a glass capillary is put into liquid mercury. The meniscus is convex and the surface of the liquid inside the tube is lower than the level of the liquid outside the tube. Mercury has greater cohesive forces than adhesive forces, so the level of the liquid will actually be lower in the capillary tubes than compared to the rest of the mercury.

Figure (b) illustrates that, because water adheres strongly to the polar surface of glass, it has a concave meniscus, whereas mercury, which does not adhere to the glass, has a convex meniscus.

Capillary action
The idea of a meniscus (concave or convex) explains why water or other liquids have when they are in narrow tubes. This is caused by the attraction between the tube and the liquid. With water, this causes it to climb up the sides of a tube. This attraction is amplified as the diameter of the tubes decreases; this is called capillary action.

This can be seen if  a tube with a very small diameter (a capillary tube) is taken and lowered it into a body of water. The liquid will climb up into the tube, even though there is no outside force. You may have also seen this when you put a straw into a drink and notice that the liquid level inside the straw is higher than it is in your drink.

NB: Capillary action requires that the adhesive forces (between the liquid and the capillary surface) be higher than the cohesive forces (between the liquid and itself), otherwise there will be no capillary action or the opposite can even happen.

3. VISCOSITY
Viscosity (η) is the resistance of a liquid to flow. Some liquids, such as gasoline, ethanol, and water, flow very readily and hence have a low viscosity. Others, such as motor oil, molasses, and maple syrup, flow very slowly and have a high viscosity.

There is also a correlation between viscosity and molecular shape. Liquids consisting of long, flexible molecules tend to have a higher viscosity than those composed of more spherical or shorter-chain molecules. The longer the molecules, the easier it is for them to become “tangled” with one another, making it more difficult for them to move past one another. London dispersion forces also increase with chain length. Due to a combination of these two effects, long-chain hydrocarbons (such as motor oils) are highly viscous.

NB:  Viscosity decreases with increase in temperature in liquids and solids, but increases with increase in temperature for gases.

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Practical demonstration of capillary action in plants
The aim of the experiment to model capillary action in plants

The Procedure

a Fill the trough with water and place a capillary tube  in it. The water will rise inside the tube, to a height h above the water surface in the cup (diagram a).

b Observe the concave curvature of water inside the capillary tube. The curvature forms an angle, say θ, on the inside wall of the capillary tube.

c Now, lower the capillary tube such that the height of the capillary tube above the water surface in the beaker is less than h (diagram b).

d Observe the concave curvature starts reducing and the angle at which the water surface meets the inside wall of the capillary tube is greater than θ.

e When the tube is fully lowered to the surface level of water in the cup you will notice that the water surface flattens (diagram c).

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Types  solids
1.  Molecular solids
2. Network solids
3. Ionic solids
4. Metallic solids

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How are they formed?

Examples of molecular solids:

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Why does ice float in water? - George Zaidan and Charles Morton

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Generally, carbon steel is the most important commercial steel alloy. Increasing carbon content increases hardness and strength and improves hardenability. But carbon also increases brittleness and reduces weldability.

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Ionic Solids are solids composed of oppositely charged ions. They consist of positively charged cations and negatively charged anions held together by electrostatic forces. Ionic solids can be composed of simple ions as seen in NaCl (sodium chloride) or can be composed of polyatomic ions as seen in ammonium nitrate (NH₄NO₃) with NH₄⁺ and NO₃⁻ ions.

In an ionic compound, the cations and anions are arranged in space to form an extended three-dimensional array that maximizes the number of attractive electrostatic interactions and minimizes the number of repulsive electrostatic interactions (as shown in the figure below).

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As shown in Equation 1 below, the electrostatic energy of the interaction between two charged particles is proportional to the product of the charges on the particles and inversely proportional to the distance between them: where Q₁ and Q₂ are the electrical charges on particles 1 and 2, and r is the distance between them.

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When Q₁ and Q₂ are both positive, corresponding to the charges on cations, the cations repel each other and the electrostatic energy is positive. When Q₁ and Q₂ are both negative, corresponding to the charges on anions, the anions repel each other and the electrostatic energy is again positive. The electrostatic energy is negative only when the charges have opposite signs; that is, positively charged species are attracted to negatively charged species and vice versa.

Lattice energy
The lattice energy of a crystalline ionic solid is the energy required to separate one mole of its solid into its component ions in the gas phase. Lattice energy is directly proportional to the product of the ionic charges and inversely proportional to the sum of the radii of the ions. For example, NaF and CaO both crystallize in the face-centered cubic (fcc) sodium chloride structure, and the sizes of their component ions are about the same: Na⁺ (102 pm) versus Ca²⁺ (100 pm), and F− (133 pm) versus O2− (140 pm). However, because of the higher charge on the ions in CaO, the lattice energy of CaO is almost four times greater than that of NaF (3401 kJ/mol versus 923 kJ/mol).

The forces that hold Ca and O together in CaO are much stronger than those that hold Na and F together in NaF, so the heat of fusion (melting) of CaO is almost twice that of NaF (59 kJ/mol versus 33.4 kJ/mol), and the melting point of CaO is 2927°C versus 996°C for NaF.

Ionic solids typically do not go from a solid state to gas state at ordinary temperatures. They can although be melted by applying thermal energy enough to interrupt the crystalline lattice. Therefore, the higher the lattice energy is of an ionic compound, the higher the melting point is.

NB:  The lower the lattice energy, the higher the quantity an ionic solid that can be dissolved in any quantity of solvent.