Bonding

Ionic bonds form when one atom gives an electron to another, creating solid crystals with high melting points that can conduct electricity in water; covalent bonds form when atoms share electrons, resulting in substances that can be gases, liquids, or solids with lower melting points and no electrical conductivity; metallic bonds occur when metal atoms share a pool of free electrons, making metals shiny, malleable, and good conductors of heat and electricity.

At the end of this section, you will be able to:

• Explain ionic bonding and its formation and state the properties of ionic compounds.

• Explain covalent bonding and its formation and state the properties of covalent compounds.

Key Ideas

• Chemical bonds are the forces that hold atoms together in a molecule or compound.

• Lewis dot symbol is a representation of an atom’s valence electrons using dots around the symbol of the element.

• Ionic bond is a type of chemical bond formed between two atoms when one atom donates an electron to another, resulting in the formation of positively charged cations and negatively charged anions.

• A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms.

• Metallic bonds are the forces that hold metal atoms together in a solid.

• Electron affinity is the energy change that occurs when an electron is added to a neutral atom in the gas phase to form a negative ion.

• Large lattice energy is the amount of energy released when ions in the gas phase come together to form a crystalline solid.

• Polarising power is the ability of a cation to distort the electron cloud of an anion.

Chemical bonds are defined as whenever two or more atoms are held strongly together. They do so in a particular way, which tends to give the constituent atoms specific whole number ratios. Gilbert Lewis, an American scientist, explained that atoms combine to achieve a more stable electron configuration. Maximum stability is attained when an atom loses, gains or shares electrons to achieve a similar electron configuration to that of a noble gas (inert gas).

In chemical bond formation, only the outermost shells of the constituent atoms come into contact. So, only the valence electrons are involved in bond formation.

The Lewis dot symbol, which consists of the symbol of an element and a dot for one valence electron in an atom, helps to explain how atoms interact to form chemical bonds. Elements in the same group on the periodic table have the same number of valence electrons, hence similar Lewis’ dot symbols, except for their atomic symbols. The constituent atoms in a molecule can donate, accept or share a certain number of electrons to form a specific type of chemical bond. The type of bond created depends on the electronegativity difference between the atoms involved.

Fig. 7.1: A chemical bond.

Image source: https://sciencenotes.org/how-to-draw-a-lewis-structure/ Atoms connect through chemical bonds. Electronegativity strongly influences how atoms interact with each other and how they bond. The electronegativity difference between two bonded atoms determines the nature of the chemical bond that forms between them. If the electronegativity difference is large (equal to or greater than approx. 1.7), the bond that forms between the atoms will be ionic, and if it is small (less than approx. 1.7), a covalent bond will generally form.

Ionic Bond

This is an electrostatic force of attraction formed between a positive ion and a negative ion. Usually, the less electronegative element completely transfers its valence electron(s) to the more electronegative element. The bond is formed between a positive ion (cation) and a negative ion (anion). This exchange results in a more stable, noble gas electron configuration for both atoms involved. An ionic bond is based on attractive electrostatic forces between two ions of opposite charge.

Ionic bonds can also be formed between species that are already ions rather than neutral atoms forming ions e.g. precipitation reactions.

Cations and Anions

When an atom, typically a metal, loses an electron or electrons and becomes a positive ion, it is called a cation. When an atom, typically a non-metal, gains an electron or electrons and becomes a negative ion, it is called an anion.

One example of an ionic bond is the formation of sodium fluoride, NaF, from sodium and fluorine atoms. In this reaction, the sodium atom loses its single valence electron to the fluorine atom, which has just enough space to accept it.

Na Na+ + e− and F + e− F− The ions produced are oppositely charged and are attracted to one another due to electrostatic forces. Na+ + F− NaF Lithium cation Fluorine anion Fig 7.2: Formation of ionic compound Lithium Fluoride (LiF).

Image source: https://mmerevise.co.uk/gcse-chemistry-revision/ionic-bonding/ Factors that affect ionic bond formation

1. Low ionisation energy of the metallic element which forms the cation.

2. Large electron affinity of the non-metallic element which forms the anion.

3. Large lattice energy i.e., smaller size and higher charge of the ions.

The electronegativity difference of the elements affects the bonding of atoms.

Elements with high electronegativity tend to form ionic bonds with elements of low electronegativity.

Properties of Ionic Compounds

1. Ionic compounds have high melting and boiling points.

2. Some dissolve in polar solvents like water.

3. They conduct electricity in the aqueous form or molten state (as the ions are free to move).

4. In their solid form, they serve as good insulators (as the ions cannot move).

5. They are hard and brittle in nature.

Covalent Bonds

This is a chemical bond that is formed when two atoms mutually share a pair of electrons. Covalent bonds are usually found between non-metal atoms. By doing so, the atoms attain a stable duplet or octet electronic configuration. In covalent bonding, overlapping of the atomic orbitals having one electron from each of the two atoms takes place, resulting in the sharing of the pair of electrons. Generally, the orbitals of the electrons in the valence shell of the atoms are used for electron sharing. For example, in a hydrogen molecule (H₂), a covalent bond is formed by the overlap of the two s-orbitals each containing one electron from each of the two H atoms of the molecule. Each electron in a shared pair of electrons is attracted to the nuclei of both atoms. Each H atom attains a 1s²configuration.

H ∙ + ∙ H H: H Fig 7.3: A covalent bond.

Image source: https://www.dummies.com/article/academics-the-arts/science/chemistry/covalent-bonds-a- hydrogen-example-194230/ Covalent bonding between atoms that have electrons beyond the first shell involves only the unpaired valence electrons. In the covalent bond formation by two fluorine atoms in F₂, there are seven (7) valence electrons, which signifies that only one (1) unpaired electron exists on a fluorine atom Only two (2) valence electrons participate in the formation of the covalent bond in F₂, leaving six (6) valence electrons in each atom not involved in forming the covalent bond. The pairs of electrons that are not used in bonding are called lone pairs. Therefore, F₂has six lone pairs of electrons (3 lone pairs on each F atom).

Fig 7.4. Image source: Raymond Chang and Jason Overby, 2009 Types of Covalent Bonds Pure covalent bond When the atoms that form a covalent bond are identical, as in H₂, Cl₂and other diatomic elements, then the electrons in the bond are shared equally. We refer to this as a pure covalent bond.

Electrons shared in pure covalent bonds have an equal probability of being near each nucleus.

Dative bond or coordinate covalent bonding This is a type of covalent bond in which both electrons in the bond are donated by one atom. That is, the pair of electrons that are shared between the two atoms comes from only one of the bonding atoms. This type of bonding is different from a typical covalent bond, where each atom contributes one electron to form a shared pair.

Formation of dative bonds (coordinate covalent bonding)

a. Formation of Ammonium ion (NH₄ +). The reaction between an ammonia molecule (NH₃) and a proton (H+) forms an ammonium ion (NH₄ +). NH₃+ H+ NH₄ +. In ammonia, three hydrogen atoms combine directly with the nitrogen atom by normal covalent bonding. The nitrogen atom has a lone pair of electrons in its outermost shell. It acts as a lone pair donor, i.e. can combine with the hydrogen ion. It shares the lone pair of electrons with a proton from an acid to produce the ammonium ion, NH₄ +. The proton carries over its positive charge to give the ammonium ion, NH₄ +.

Fig 7.5: Formation of Ammonium ion.

Image source: https://byjus.com/chemistry/ammonium-ion/

b. Formation of a hydroxonium ion (H₃O+). H₂O + H+ H₃O+. In a water molecule, two hydrogen atoms share two pairs of electrons with an oxygen atom by normal covalent bonding. The oxygen atom in a water molecule has two lone pairs of electrons in its outermost shell. It shares this with a proton (hydrogen ion, H+) from an acid to produce the hydroxonium ion H₃O+. The positive charge on the hydrogen ion is carried over to give the positively charged hydroxonium ion, H₃O+.

Fig. 7.6: Formation of hydroxonium ion.

Image source: https://www.sarthaks.com/274945/explain-the-formation-of-h3o-and-nh4-ion Polar covalent bond This is a type of chemical bond formed between two atoms, where the electrons are shared unequally. In this type of bond, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom. The atom that attracts the electrons more strongly acquires a partial negative charge whilst the other acquires a partial positive charge. For

example, the electrons in the H–Cl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom.

Whether a bond is non-polar or polar covalent is determined by a property of the bonding atoms called electronegativity. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarised the electron distribution and the larger the partial charges of the atoms.

Fig. 7.7: Formation of HCl.

Image source: https://socratic.org/questions/how-does-a-covalent-bond-become-polar Factors That Affect the Formation Of Covalent Bonds

1. Number of valence electrons

2. High ionisation energy

3. Comparable electron affinity

4. Comparable electronegativities

5. Atomic size

6. High nuclear charge and small internuclear distance.

Properties Of Covalent Compounds

Compounds that contain covalent bonds exhibit different physical properties than ionic compounds. Covalent compounds;

1. Exist in three states of matter namely solid, liquid and gas.

2. Do not conduct electricity in a solid, molten or aqueous state.

Polarisation During the formation of an ionic compound or ionic molecules, two oppositely charged ions (cations and anions) must come closer to each other. During this process, the cation attracts the electron charge cloud of the outermost shell of the anion toward itself. Therefore, the symmetrical shape of the anion gets distorted, deformed or polarised.

The phenomenon of the distortion of the symmetrical shape of the electron cloud of an anion by the nearby cation is called polarisation of anion.

There is also the chance of the polarisation of a cation by an anion. However, due to the smaller size of the cation, its electron cloud is strongly held to the nucleus, and the shape of the cloud is not distorted to an appreciable extent. Hence, the polarisation of a cation by an anion is not generally considered.

Polarising Power

The ability of a cation to polarise (distort) an anion is called its polarising power or polarisability.

Fig. 7.8: Polarising power of a cation Factors affecting polarising power The factors affecting the magnitude of a cation’s polarising power are listed below:

a. Magnitude of positive charge on the cation The greater the charge on the cation, the more strongly it attracts the outermost shell electron cloud of an anion toward itself and polarises the given anion easily. Therefore, the polarising power of a cation is directly proportional to the magnitude of the positive charge on it. E.g. Na +< Mg ₂₊< Al ³⁺. If the same element has a different positive charge, the higher positive charge has greater power of polarisation. E.g., Sn ⁴⁺> Sn ²⁺.

b. Size of cation The smaller the size of the cation, the more strongly it attracts the outermost shell electron cloud of an anion towards itself, hence, the greater its polarising ability. In other words, with the decreasing size of the cation, the polarising power of the cation increases. Thus, the polarising power of the cation is inversely proportional to the size of the cation. Example:

Li+> Na+> K+> Rb+

Fig. 7.9: Factors affecting polarising power.

Polarisability of anions The tendency of an anion to get polarised by a cation is called its polarisability.

The factors affecting the polarisability of an anion are magnitude of the negative charges on the anion and the size of the anion.

a. Magnitude of negative charges on anion The higher the negative charge on the anion, the more easily its outermost electron cloud is attracted by cations, hence the Polarisability of an anion is directly proportional to the magnitude of the negative charges on it. E.g., C ⁴⁻> N ³⁻> O ²⁻> F −

b. Size of the anion The larger the size of the anion, the more easily its outermost shell electron cloud is attracted by the cation towards itself, hence, the greater the polarisability of an anion. Thus, the polarisability of an anion is directly proportional to the size of an anion.

This is a type of chemical bond that occurs in metals. It is an electrostatic force of attraction between the fixed positive metal ions and the delocalised electrons around the cations. The lattice structure of the atoms is held together by a sea of delocalised valence electrons. The delocalised electrons can also be referred to as a sea of electrons or mobile electrons.

Metallic Bonds

Metallic bonds are formed by the process of metal atoms transferring their outermost electrons to an electron sea, which is the collection of shared electrons that surround the positively charged atomic cores of the metal atoms. When heated, the metal atoms release their outermost electrons, which then move freely throughout the lattice. Since metal atoms have very low electronegativity, they tend to lose their valence electrons, readily forming positive ions in the process.

These positive ions are then surrounded by a cloud of delocalised electrons, which form the metallic bond that holds the ions together in a regular lattice structure.

Fig. 7.10: A metallic bond The strength of the metallic bond depends on factors such as the number of valence electrons in the metal atom, the size of the atoms and the proximity of the atoms in the lattice. Metallic bond strength increases as the number of valence electrons increases. Thus, aluminium is harder than magnesium, which is in turn harder than sodium because 3, 2, and 1 valence electrons are attracted by the fixed positive lattice points respectively. As the atomic size of these atoms reduces with increased nuclear attraction, their melting points also increase accordingly.

Properties of Metals

Malleable Think of a lump of clay. You can flatten it with your hands or a rolling pin into a thin sheet. Metals are kind of like that. Malleable means a metal can be hammered or pressed into different shapes, especially thin sheets, without breaking.

This is why we can make things like aluminium foil, car bodies, and metal sheets for roofs. If metals were brittle (like glass), they would just shatter when you tried to flatten them.

Gold is extremely malleable. You can hammer it into incredibly thin sheets called gold leaf, which are used for decoration.

Ductile Ductile means a metal can be stretched into thin wires. This is how we make electrical wires, which are essential for carrying electricity to our homes and devices. Copper is very ductile, which is why it’s used so much for electrical wiring.

Good conductors of heat and electricity

1. Heat: If you put a metal spoon in a hot cup of tea, the spoon quickly gets hot too. This is because metals allow heat to flow through them easily. “Good conductor of heat” means heat travels through the metal quickly. This is why we use metal pots and pans for cooking. They heat up quickly and evenly.

2. Electricity: Electricity is the flow of electrons. Metals have free electrons that can move easily, allowing electricity to flow through them. “Good conductor of electricity” means electricity travels through the metal easily. This is why we use metal wires to carry electricity.

Silver and copper are excellent conductors of both heat and electricity.

Activity 7.1: Malleability Demonstration (Aluminium Foil Shaping)

Materials needed:

• Aluminium foil

• A small hammer (or a heavy spoon)

• A flat surface (like a cutting board) Procedure:

1. Take a piece of aluminium foil.

2. Try to shape it with your hands. Notice how easily it bends and changes shape.

3. Place the foil on a flat surface.

4. Gently tap the foil with the hammer (or spoon). Observe how the foil flattens and spreads out without breaking easily.

Activity 7.2: Ductility Demonstration (Copper Wire Stretching)

Materials:

• Thin copper wire (easily found in some discarded electronics or crafting supplies)

• Two pairs of pliers (optional, for better grip) Safety Note: Be careful not to cut yourself with the wire.

Procedure:

1. Hold each end of the copper wire firmly.

2. Gently and slowly pull on the wire, trying to stretch it.

3. Observe how the wire becomes thinner and longer before it eventually breaks.

4. If you have pliers, grip each end of the wire with the pliers, and pull.

Activity 7.3: Heat Conductivity Demonstration (Spoon in Hot Water) Materials:

• A metal spoon (preferably a stainless steel or silver spoon)

• A plastic or wooden spoon (for comparison)

• A mug or heat-safe container

• Hot water (be careful!)

Procedure:

1. Pour hot water into the mug.

2. Place the metal spoon and the plastic/wooden spoon into the hot water, with the handles sticking out.

3. Wait for a minute or two.

4. Carefully touch the handles of both spoons.

5. Notice the temperature difference.

Safety notes:

• Teacher supervision is recommended, especially when working with hot water or tools.

• Always be careful when handling any type of wire, especially when handling any electricity. Never use household current to test conductivity.

Review Quetion 7.1

1. In which of these compounds is the ionic bond strongest?

NaCl, MgCl₂,AlCl₃

2. Explain why sodium chloride does not conduct electricity when solid.

3. Why do ionic compounds have high boiling and melting points?

4. Use Lewis dot structures to explain the nature of chemical bonding found in oxygen gas (O₂) and identify the number of bonds formed in the molecule

There are different types of intermolecular forces between molecules. These forces arise from the molecular structures of the molecules, and they influence the physical properties of the compounds. In this section, we will delve into explaining how these forces affect the properties of the molecular compounds.

At the end of this section, you will be able to:

• Describe the different types of intermolecular forces and explain how they arise from the structural features of molecules

• Explain how intermolecular forces affect the physical properties of compounds

Key Ideas:

• Dipoles - these are pairs of equal and oppositely charged species.

• Hydrogen bond - a weak bond between two molecules resulting from an electrostatic attraction between a proton in one molecule and an electronegative atom in the other.

• Van der Waal Forces - weak electrostatic forces that attract neutral molecules to one another.

• Intermolecular forces - are the attractive and repulsive forces that arise between molecules of a substance.

• Intra-molecular forces - Intramolecular forces are the attractive forces that hold atoms together within a molecule.

• Induced dipole–induced dipole (London Dispersion Forces) - this is a weak attraction that occurs when a polar molecule causes a dipole to form in an atom or non-polar molecule by disrupting the electron configuration in the non-polar species.

Intermolecular forces (IMF) are the forces that attract or push away molecules in a substance. These forces control how molecules interact with each other. They also help explain many of the physical and chemical properties of different materials.

Intermolecular forces are the attractions that pull molecules together. This is different from intramolecular forces, which hold the atoms together inside a molecule, like covalent bonds. Covalent bonds are much stronger than intermolecular forces because they keep the molecule itself together.

Intermolecular forces happen because of electric charges, dipoles (molecules with positive and negative ends), and hydrogen bonds. The most common types of intermolecular forces are:

1. Dipole-dipole forces

2. Hydrogen Bonds

3. Induced Dipole–Induced Dipole Forces

Covalent bonds are stronger than intermolecular forces. For example, covalent bonds are usually 50-200 kJ/mol in strength, while intermolecular forces are weaker, usually 1-12 kJ/mol. Intramolecular forces keep molecules stable (the atoms making up the molecule remain the same), while intermolecular forces affect properties like:

a. Melting and boiling points

b. Solubility

c. Viscosity To boil a substance, enough energy is needed to completely break the intermolecular forces holding the molecules together, so they can turn into gas. Substances with higher boiling points have stronger intermolecular forces. Those with the highest boiling points have intermolecular forces that are also intramolecular forces e.g.

diamond is made up of a network of covalent bonds hence its high boiling point.

Van der Waal Forces Covalent molecules are usually held together by Van der Waals forces, which are forces of attraction between molecules. These forces of attraction come about because of opposite charges that are either permanently present or temporarily present, depending on the atoms in the compound.

There are three types of Van der Waals forces based on how the charges are spread in the molecules:

1. Dipole-Dipole Interactions are forces that occur between molecules that have positive and negative ends (dipoles).

2. Hydrogen Bonding is a strong type of dipole interaction that happens when hydrogen is bonded (within the molecule) to highly electronegative atoms like oxygen or nitrogen or fluorine.

3. Induced Dipole–Induced Dipole (London Dispersion) Forces: These are weak forces that happen when temporary charges appear in molecules as they get close to each other.

Dipole - Dipole interaction Dipole-dipole interaction, also called dipolar interaction, is a type of force that happens between two polar molecules. Polar molecules have a permanent separation of positive and negative charges. This happens because the atoms in the molecule pull on the shared electrons with different strengths.

Since some atoms in molecules pull harder on the shared electrons, one atom ends up with the shared electrons nearer to it and gets a partial negative charge, while the other atom has fewer electrons nearer to it and gets a partial positive charge.

In dipole-dipole interactions, the positive end of one molecule is attracted to the negative end of another, creating a force that pulls the molecules together.

Fig. 8.1: Dipole-dipole interactions Dipole-dipole interactions are weaker than covalent bonds but stronger than London dispersion forces (another type of force that will be explained later).

These interactions are important because they help explain the physical and chemical properties of polar molecules.

Dipole-dipole interactions also help determine how easily a polar molecule can dissolve in a polar solvent (like water). The attraction between the molecules of the substance and that of the solvent makes the substance more soluble. These interactions can also affect the shape and stability of molecules when they are in solid or liquid forms.

Hydrogen Bond

A hydrogen bond is a special kind of dipole-dipole interaction. It happens when a hydrogen atom is bonded to a small but very electronegative atom like oxygen, nitrogen, or fluorine. The electronegative atom pulls more on the electrons, creating a partial negative charge on itself, while the hydrogen atom gets a partial positive charge.

The partially positively charged hydrogen is then attracted to a partially negatively charged electronegative atom in a nearby molecule. This attraction forms a hydrogen bond between the two molecules.

Fig 8.2: Hydrogen bond Hydrogen bonds are stronger than London dispersion forces (also called induced dipole–induced dipole forces) and dipole-dipole forces, but they are weaker than covalent bonds. Hydrogen bonds are important because they affect the properties of many substances, like water, ice, and biomolecules such as DNA and proteins.

For example, the special properties of water—like its unusually high boiling point, surface tension, and density—are caused by hydrogen bonds between molecules of water. In DNA, hydrogen bonds between base pairs help keep the double helix structure together.

Some common compounds that have their hydrogen atoms bonded to electronegative atoms (like oxygen, nitrogen, and fluorine) include:

Water (H₂O), Ammonia (NH₃), Hydrogen fluoride (HF), Methanol (CH₃OH), Ethanol (C₂H₅OH), Acetic acid (CH₃COOH), Phenol (C₆H₅OH), Sugars (like glucose and fructose), Proteins (like collagen and keratin) and RNA.

These compounds have special properties that make them important in many fields, such as medicine, chemistry, and biology.

Induce dipole–induce dipole (London dispersion) forces London dispersion forces are types of intermolecular bonds that occur between non-polar covalent molecules. These forces occur because of small changes in electron density, which cause an uneven distribution of charge.

This uneven charge distribution creates temporal dipoles in a molecule, meaning one part of the molecule becomes slightly positive and another part slightly negative. This temporary created dipole can cause a nearby molecule to also form a dipole, and then the two molecules are attracted to each other.

Fig. 8.3: Induce dipole-induce dipole forces Examples of non-polar molecules which have weakly induced dipole-induced dipole intermolecular forces are diatomic molecules (O₂, H₂, l₂, Br₂); noble gases (He, Ar, Ne): CO₂, CH₄, CCl₄, polythene and rubber.

Activity 8.1

Pair up with your colleague and answer the following questions:

1. In terms of intermolecular forces, explain why the solubility of ethanol in water is greater than the solubility of ether in water.

2. Consider the following compounds: CH₄, CO₂, NH₄Cl. Select the compound capable of undergoing dipole-dipole interactions and explain your answer.

Intermolecular forces (IMFs) help us understand the physical properties of substances, such as boiling points, melting points, and viscosity. They play a key role in determining how molecules behave in different situations.

1. Boiling and Melting Points

The strength of IMFs affects how much energy is needed to change a substance from a solid to a liquid (melting) or from a liquid to a gas (boiling).

Substances with stronger IMFs need more energy to melt or boil, so they have higher melting and boiling points. For example, water has strong hydrogen bonds and a high boiling point.

Substances with weaker IMFs need less energy, so they have lower melting and boiling points. For example, methane has weak forces and a low boiling point.

2. Solubility Solubility is the ability of one substance to dissolve in another. It depends on the type of intermolecular forces between the solute (the substance being dissolved) and the solvent (the substance doing the dissolving).

The rule “like dissolves like” means substances with similar types of IMFs dissolve in each other. For example, polar substances dissolve well in polar solvents (like water), while non-polar substances dissolve in non-polar solvents (like oil).

Hydrogen bonding helps substances like alcohol dissolve in water.

It is worth noting that this is not the only thing which could affect solubility;

temperature and pressure also affect solubility. For example, gases dissolve better at higher pressures.

3. Surface Tension

Surface tension is the property of a liquid’s surface that causes it to behave like a stretched elastic membrane. This phenomenon arises from the cohesive forces between molecules at the surface, which create a tendency for the surface to minimize its area.

Liquids with stronger IMFs have higher surface tension. Liquids with weaker forces have lower surface tension.

It is worth noting that this is not the only thing which could affect surface tension; Increasing temperature lowers surface tension because molecules move faster (and with greater energy) meaning they can more readily overcome the forces at the surface.

4. Enthalpy of Vaporisation:

The enthalpy of vaporisation is the energy needed to change a liquid into a gas. Substances with stronger IMFs need more energy to vaporise, so they have a higher enthalpy of vaporisation.

Substances with weaker IMFs, like noble gases, have a lower enthalpy of vaporisation because they need less energy to become a gas.

5. Viscosity Viscosity is how thick or sticky a liquid is and how easily it flows. Liquids with stronger IMFs have higher viscosity because their molecules are held together more, making it harder to flow. For example, honey flows slowly because of strong hydrogen bonds.

Liquids with weaker IMFs, have lower viscosity and flow easily.

6. Volatility Volatility refers to how easily a substance evaporates. Substances with stronger IMFs have lower volatility because their molecules are held tightly together, making it harder to evaporate.

Substances with weaker IMFs have high volatility and evaporate easily.

These properties help us understand how substances behave in different situations, from boiling water to the flow of syrup!

Activity 8.2: Intramolecular and intermolecular forces of attraction.

1. Using the internet, watch videos of intramolecular and intermolecular forces by clicking the links below.

• https://www.youtube.com/watch?v=jucya5UsZz0

• intermolecular forces-Khan Academy

2. Using this knowledge, define Intramolecular and intermolecular forces of attraction. In pairs, discuss your findings and answers. Make sure to include and refer to a labelled diagram in your discussion, as well as a comparison of all the forces you have found.

Activity 8.3: Hydrogen bond formation Materials needed: pencil, drawing paper (or marker board and a marker).

1. Using the following molecules (HF, NH₃, CH₄,and H₂O), draw the structural formula of two of the molecules (e.g., ammonia, NH₃, and methane, CH₄).

2. Identify the electronegative atoms and the hydrogen atoms bonded to them.

3. Use different colours for the atoms for identification.

4. Identify how the shared electrons will be influenced by the different atoms on either side of the covalent bond.

5. Show how the molecules approach each other, forming hydrogen bonds.

6. Discuss and compare your results with your friends

Activity 8.4

Study the molecular structure of the following:

Water (H₂O); Ammonia (NH₃); Methanol (CH₃OH); Ethanol (C₂H₅OH);

Acetic acid (CH₃COOH

1. Draw the structural formula of a molecule capable of forming hydrogen bonds.

2. Identify the electronegative atom and hydrogen atoms bonded to it.

Extended Activity

1. Compare and contrast the physical properties of ethanol (C₂H₅OH) and methane (CH₄) in terms of their intermolecular forces. Explain why ethanol has a higher boiling point than methane.

2. Design an experiment to investigate the relationship between the strength of intermolecular forces and the surface tension of different liquids.

Include a hypothesis, materials list, procedure and expected results.

3. Analyse the impact of intermolecular forces on the solubility of various substances in water. Consider how factors such as molecular structure, polarity and hydrogen bonding influence solubility. Provide examples to support your analysis.

Review Questions 8.1

1. Which type of intermolecular force is responsible for the high boiling point of water?

2. Arrange each of the following sets in order of increasing boiling point,

a. HCl, H₂O, SiH₄

b. F₂, Cl₂, Br₂

c. CH₄, C₂H₆, C₃H₈

d. O₂, NO, N₂

3. Deduce the type of intermolecular force present in each of the following compounds:

a. H₂S

b. NH₃

c. CH₄

d. SF₆ e . HF