Chemical Bonding and Intermolecular Forces

Chemical Bonding Overview

• Elements react in order to have a full outer shell and become stable.

• Intramolecular Bond: The bond between atoms in a molecule. It is a force that holds the atoms or ions in a compound together.

• Elements bond by either losing, gaining, or sharing electrons.

• Valency: The electrons used in chemical bonding are the unpaired electrons in the last shell.

• There are three main types of intramolecular bonds:   - (a) Covalent bonding   - (b) Ionic bonding   - (c) Metallic bonding

Covalent Bonding

• Definition: A covalent bond is the sharing of at least one pair of electrons by two non-metal atoms.

• The electrons shared are the unpaired electrons in the outer shell.

• Examples of Covalent Molecules:   - Hydrogen Molecule ($H_2$):     - Formed by the bonding of two hydrogen atoms ($H + H$).     - Lewis diagram represents the sharing of a single pair of electrons ($H:H$).     - By sharing, each atom attains a full outer shell.     - Single Bond: When only one pair of electrons is shared. Represented as $H-H$.   - Oxygen Molecule ($O_2$):     - There are 2 pairs of electrons shared.     - By sharing, each atom attains a full outer shell.     - Double Bond: When two pairs of electrons are shared. Represented as $O=O$.   - Nitrogen Molecule ($N_2$):     - There are 3 pairs of electrons shared.     - Each nitrogen atom attains a full outer shell.     - Triple Bond: When three pairs of electrons are shared. Represented as $N \equiv N$.   - Ammonia ($NH_3$):     - Formed by 1 nitrogen atom and 3 hydrogen atoms.     - Represented structurally as hydrogen atoms bonded to a central Nitrogen: $H-N(H)-H$.   - Hydrogen Fluoride ($HF$):     - Formed between 1 fluorine atom and 1 hydrogen atom.     - Represented as $H-F$ (The transcript also notes representation as $H-Cl$ likely as a comparison or alternate example).   - Water ($H_2O$):     - Formed between 2 hydrogen atoms and 1 oxygen atom.     - Represented as $H-O-H$.   - Ethene ($C_2H_4$):     - Each carbon atom shares all four of its electrons.     - It shares two electrons with two hydrogen atoms and two with the other carbon atom.     - This forms a carbon-carbon double bond ($C=C$).

Properties of Covalent Compounds

• They are made up of molecules.

• The molecules are held together by weak forces.

• Melting and Boiling Points: They have low melting and boiling points because the intermolecular forces (weak forces) are easy to break.

• Volatility: They are very volatile.

• Solubility:   - They are generally insoluble in water.   - They are soluble in covalent solvents.

• Conductivity: They do not conduct electricity.

Molecular Elements and Diatomic Elements

• Molecule: A group of atoms held together by covalent bonds.

• Molecular Element: Elements that exist as molecules (e.g., Hydrogen, where the formula $H_2$ indicates two atoms per molecule).

• Diatomic Elements: Elements made up of molecules containing exactly two atoms. Examples include:   - Iodine: $I_2$   - Fluorine: $F_2$   - Nitrogen: $N_2$   - Bromine: $Br_2$   - Oxygen: $O_2$

• Note: Sulphur exists as $S_8$ (polyatomic).

Questions & Discussion: Exercise 1HW

• 1. Draw Lewis diagrams for:   - (a) $HF$   - (b) $Cl_2$   - (c) $CH_4$   - (d) $C_2H_6$   - (e) $CO_2$

• 2. What is a covalent bond? (Sharing of at least one pair of electrons by two non-metal atoms).

• 3. How do atoms in covalent compounds gain a stable outer shell? (By sharing unpaired outer-shell electrons).

• 4. Give 3 properties of covalent compounds. (Low melting point, non-conductive, insoluble in water).

Ionic Bonding

• Definition: An ionic bond is a transfer of electrons from a metal to a non-metal and subsequent electrostatic attraction.

• It is a force of attraction between oppositely charged ions in a compound.

• Formation:   - Metals: Lose electrons and form positive ions (cations) because they then have more protons than electrons.   - Non-metals: Gain electrons and form negative ions (anions) because they then have more electrons than protons.

• Examples:   - Sodium Chloride ($NaCl$):     - Sodium (electronic configuration $2,8,1$) loses 1 electron to become $[2,8]^+$.     - Chlorine ($2,8,7$) gains 1 electron to become $[2,8,8]^-$.     - The resulting $Na^+$ and $Cl^-$ ions attract each other due to opposite charges.   - Magnesium Oxide ($MgO$):     - Magnesium atom ($2,8,2$) loses 2 outer electrons to an Oxygen atom ($2,6$).     - Forms Magnesium ion ($Mg^{2+}$) and Oxide ion ($O^{2-}$).   - Lithium and Fluorine: Formation of Lithium Fluoride (Note: Transcript text mentions "forming aluminium chloride" in error on page 7; diagram shows $LiF$).

Properties of Ionic Compounds

• They are made up of ions.

• Ions are held together by strong electrostatic forces.

• Melting and Boiling Points: They have high melting and boiling points because more energy is required to break the strong electrostatic forces.

• Solubility: They are usually soluble in water.

• Conductivity:   - They conduct electricity when molten or dissolved in water (ions are free to move).   - They do not conduct electricity when solid (ions are fixed in position).

• Volatility: They are less volatile because ions are held by strong forces.

Questions & Discussion: Exercise on ionic bond : 2

• 1. Explain why a calcium ion has a charge of $2+$? (It loses 2 valence electrons and has 2 more protons than electrons).

• 2. Why is the charge on an aluminium ion $3+$? (It loses 3 valence electrons).

• 3. Show bond in high compounds: (a) $MgCl_2$, (b) $Al_2O_3$, (c) $Na_2O$.

• 4. What is the ionic bond? (Transfer of electrons and electrostatic attraction between ions).

• 5. State 3 properties of ionic compounds. (High melting point, conductive when molten, ionic lattice state).

Metallic Bonding

• Definition: The metallic bond is the force of attraction between the free-moving (delocalised) electrons and positive metal ions.

• Atoms of metals are kept together due to the attraction between the positive kernels and the "sea" of delocalised electrons.

• Properties of Metals:   - High melting points.   - Good conductors of electricity.   - Good conductors of heat.   - High density.   - Malleable (can be hammered into shape).   - Ductile (can be drawn into wires).   - Soluble in ionic solvents.   - Shiny (lustrous) in nature.

Structures of Matter

There are five primary structures of matter:

1. Simple Atomic Structure

• Exists in noble gas elements ($He, Ne, Ar, Kr, Xe, Rn$).

• The atoms are held together by very weak Van der Waals forces, specifically London forces (induced dipole).

• They have very low melting and boiling points.

• Trend: Melting and boiling points increase down the group because the increase in the number of electrons increases the strength of the London forces.

• Melting and Boiling Point Data:   - Helium: MP $-272\,^\circ\text{C}$, BP $-269\,^\circ\text{C}$   - Neon: MP $-249\,^\circ\text{C}$, BP $-246\,^\circ\text{C}$   - Argon: MP $-189\,^\circ\text{C}$, BP $-186\,^\circ\text{C}$   - Krypton: MP $-157\,^\circ\text{C}$, BP $-153\,^\circ\text{C}$   - Radon: MP $-71\,^\circ\text{C}$, BP $-62\,^\circ\text{C}$

2. Simple Molecular Structure

• Consists of molecules where atoms are joined by strong covalent bonds.

• Molecules themselves are held together by weak Van der Waals forces.

• Often liquids or gases at room temperature.

• Melting/boiling points are higher than noble gases because Van der Waals forces in molecules are stronger than in atoms.

• Types of Van der Waals Forces:   - (a) London Forces: If the molecule is non-polar.   - (b) Permanent Dipole: If the molecule is polar.

• Example: Halogens:   - Fluorine ($F_2$): Yellow, MP $-220\,^\circ\text{C}$, BP $-188\,^\circ\text{C}$, State: GAS.   - Chlorine ($Cl_2$): Green, MP $-101\,^\circ\text{C}$, BP $-35\,^\circ\text{C}$, State: GAS.   - Bromine ($Br_2$): Orange-brown, MP $-7\,^\circ\text{C}$, BP $59\,^\circ\text{C}$, State: LIQUID.   - Iodine ($I_2$): Purple, MP $+114\,^\circ\text{C}$, BP $184\,^\circ\text{C}$, State: SOLID.

• Trend: Colours darken, states change from gas to solid, and MP/BP increase down the group due to stronger Van der Waals forces (more electrons).

3. Giant Covalent Structure

• Atoms are linked in a continuous network of covalent bonds.

• Examples include Diamond, Graphite, and Silicon (IV) Oxide ($SiO_2$).

• Substances have very high melting points because strong covalent bonds must be broken.   - Diamond: $3550\,^\circ\text{C}$   - Graphite: $2880\,^\circ\text{C}$   - Silicon (IV) Oxide: $1710\,^\circ\text{C}$

• Allotropes of Carbon:   - Diamond: Each Carbon is bonded to 4 others. Rigid network. Hardest known substance. High refractive index (jewellery). Used in drills.   - Graphite: Each Carbon is bonded to 3 others in hexagonal layers. The 4th electron is delocalised, making it a good conductor. Layers slide over each other (lubricant). Used for electrodes and carbon brushes.

• Silicon (IV) Oxide: Each silicon is bonded to 4 oxygen atoms in a rigid network similar to diamond.

4. Giant Ionic Structure

• Exists in all ionic compounds.

• Ions are held by strong electrostatic forces in a lattice.

• High melting and boiling points.

• Conduct electricity only when molten or aqueous.

5. Giant Metallic Structure

• Exists in all metals.

• High melting and boiling points.

• Conducts electricity in both solid and molten states due to delocalised electrons.

Questions & Discussion: Exercise 5

Substance Property Table:

• (a) Giant metallic structure? (D and G, because they conduct in both solid and liquid states and have high melting points).

• (b) Giant ionic structure? (B and E, because they conduct only when liquid and have high melting points and are soluble).

• (c) Simple molecular structure? (A and C, because they have very low melting points and poor conductivity).

• (d) Giant covalent structure? (F, due to the extremely high melting point and poor conductivity).

• (e) Bonding type? B is ionic; A is covalent.

Electronegativity and Polarity

• Electronegativity: A measure of the ability of an atom in a molecule to attract shared electrons.   - Fluorine: $4.0$   - Oxygen: $3.5$   - Nitrogen: $3.0$

• Polarity: Uneven distribution of the electron cloud within a molecule, caused by differences in electronegativity.

• Polar Molecules: Occur when atoms share electrons unequally, forming a dipole (slight charges $\delta+$ and $\delta-$).   - If the electronegativity difference is between $0.5$ and $2.0$, it forms a polar covalent bond.   - Examples: $H_2O$, $HF$, $NH_3$, $HCl$, Chloromethane ($CH_3Cl$).   - In Chloromethane, Chlorine pulls electrons towards itself, gaining a partial negative charge.   - Polar molecules have higher melting/boiling points than non-polar molecules because Van der Waals forces (permanent dipoles) are stronger.

• Non-Polar Molecules: Electrons are shared equally or distributed symmetrically.   - Electronegativity difference is less than $0.5$.   - Examples: Noble gases ($He, Ne$), Diatomic elements ($H_2, N_2, O_2, Cl_2$).   - Symmetrical Molecules with polar bonds: $CO_2$ and $CCl_4$. In $CO_2$, polar bonds are present, but the oxygen atoms pull with equal force in opposite directions, resulting in zero net polarity.

Intermolecular Forces

An intermolecular force is a weak force of attraction between molecules, ions, or atoms. Broken during melting and boiling (except in giant molecular structures).

1. London Forces (Temporary or Induced Dipole)

• Weakest intermolecular force.

• Results from temporary fluctuations in electron density which induce dipoles in adjacent atoms.

• Strength increases with an increase in the number of electrons.

• Group 7 Halogen Trend (Boiling Points):   - Fluorine ($18$ electrons): $-188\,^\circ\text{C}$   - Chlorine ($34$ electrons): $-35\,^\circ\text{C}$   - Bromine ($70$ electrons): $+59\,^\circ\text{C}$   - Iodine ($106$ electrons): $+184\,^\circ\text{C}$

• Group 8 Noble Gas Trend (Increasing Atomic Number increases BP):   - $He (2) \rightarrow BP\, 4.4\,K$   - $Ne (10) \rightarrow BP\, 27.3\,K$   - $Ar (18) \rightarrow BP\, 87.4\,K$   - $Kr (36) \rightarrow BP\, 121.5\,K$   - $Xe (54) \rightarrow BP\, 166.6\,K$   - $Rn (86) \rightarrow BP\, 211.5\,K$

• Hydrocarbons (Alkanes):   - Boiling point increases with size: Methane ($-162\,^\circ\text{C}$) to Pentacontane ($625\,^\circ\text{C}$).   - Branching Effect: More branched isomers have lower boiling points due to less surface area for contact between molecules, weakening London forces.     - Pentane (linear): $36.1\,^\circ\text{C}$     - 2-methylbutane: $27.8\,^\circ\text{C}$     - 2,2-dimethylpropane: $9.5\,^\circ\text{C}$

2. Dipole-Dipole Forces (Permanent Dipole)

• Attraction between the $\delta+$ end of one polar molecule and the $\delta-$ end of another.

• Stronger than London forces but much weaker than ionic/covalent bonds.

• Significant effect only when molecules are close together.

• Example: $HCl$ molecules.

3. Hydrogen Bond

• A special, strong type of dipole-dipole attraction.

• Conditions:   - Hydrogen must be covalently bonded to a very electronegative atom ($N$, $O$, or $F$).   - The $N$, $O$, or $F$ must have at least one lone pair of electrons.

• Water ($H_2O$):   - Oxygen has 2 lone pairs and 2 hydrogens; forms 2 H-bonds per molecule.   - This results in water having a much higher boiling point than other group 6 hydrides ($H_2S, H_2Se, H_2Te$).

• Ammonia ($NH_3$) and Hydrogen Fluoride ($HF$):   - Ammonia forms 1 H-bond (restricted by one lone pair).   - $HF$ forms 1 H-bond (restricted by one hydrogen atom).   - $HF$ has a higher boiling point than $NH_3$ because it is more polar ($F$ is more electronegative than $N$).

4. Ion-Dipole Forces

• Attraction between an ion and a neutral polar molecule.

• Important for the solubility of ionic compounds in polar liquids.

• Example: $Na^+$ surrounded by water molecules ($\delta-$ oxygen end attracts to the cation).

Comparison of Intramolecular and Intermolecular Forces

• Intramolecular forces are strong bonds (Covalent, Ionic, Metallic) within a molecule or lattice.

• Intermolecular forces are weak attractions between separate molecules (London, Dipole-dipole, Hydrogen bonding).

• Standard strength hierarchy: London < Dipole-dipole < Hydrogen bond << Ionic/Covalent/Metallic bond.