The covalent model - IB DP Chemistry- Study Notes - New Syllabus 2025
The COVALENT model – IB DP Chemistry- Study Notes
IITian Academy excellent Introduction to the Particulate Nature of Matter – Study Notes and effective strategies will help you prepare for your IB DP Chemistry 2025 exam.
- IB DP Chemistry 2025 SL- IB Style Practice Questions with Answer-Topic Wise-Paper 1
- IB DP Chemistry 2025 SL- IB Style Practice Questions with Answer-Topic Wise-Paper 2
- IB DP Chemistry 2025 HL- IB Style Practice Questions with Answer-Topic Wise-Paper 1
- IB DP Chemistry 2025 HL- IB Style Practice Questions with Answer-Topic Wise-Paper 2
S2.2.1 – Covalent Compounds and Molecules
Electron Sharing: nonpolar
- 2 non-metals react together, both want enough electrons to create a full shell of 8.
- Known as the octet rule
- The shared pair of electrons is concentrated in the region between the 2 positively charged nuclei.
- The electrostatic attraction between the 2 nuclei and the electrons constitutes the covalent bond.
Covalent Character:
Two ways to predict:
- Position on the Periodic Table
- Electronegativity differences
Position on the Periodic Table
- Covalent compounds tend to form between 2 nonmetals.
- The closer together two elements are, the more covalent.
Electronegativity
- Electronegativity values are given in the IB Data Booklet
- Differences less than 1.8 are covalent.
Naming Covalent Compounds:
- End the name with -ide
- Sodium Chloride, Sulfur Dioxide, Carbon Tetrachloride
- Prefixes depending on the amount of the atoms
- Mono (this one isn’t really mandatory)
- Di
- Tri
- Tetra
- Penta
- Hexa
- Hepta
- Octa
- Nona
- Deca
Lewis Formula or Dot Structures:
- Used to represent the valence electrons of atoms in covalent molecules
- Dots are written between symbols to represent bonding electrons
Lewis Formula Rules:
- Add up the total number of valence electrons in the molecule.
- Draw the skeletal structure.
- Use a line between each element to symbolize an electron pair.
- Distribute the remaining electrons around the elements in pairs to form octets. (Hydrogen can only ever have 2 electrons.) Make sure ALL electrons are shown in the diagram.
- Adjust or charge if it is a polyatomic ion
- Ions must have square brackets around them with the charge notated in the top right hand corner.
S2.2.2 – Bond Order
Bond Length and Bond Strength:
- Single bond → 1 electron pair shared, 2 electrons total
- Double bond → 2 electron pairs shared, 4 electrons total
- Triple bond → 3 electron pairs shared, 6 electrons total
1. Bond Strength:
- Triple bond > Double bond > Single bond
- Bond strength is a measure of how much energy is required to break the bond.
2. Bond Length:
- Shared electrons increase, electrostatic attraction increases.
- Stronger attraction, less repulsion, less bond length
S2.2.3-4 – Coordination Bond
Coordination Bonds:
- Occur when one atom donates both electrons that are shared between two atoms.
- These bonds are also called Dative Bonds
VSEPR(T): Molecular Geometry – (this is important)

The term ‘Actual Shape’ refers to the Molecular Domain Geometry (MDG)
The term ‘Basic Shape’ refers to the Electron Domain Geometry (EDG)
Determining the Shape:
- Movement of delocalized electrons in double/triple bonds
- Molecular Domain Geometry shape (MDG): Bent
- Electron Domain Geometry shape (EDG): Trigonal Planar
- The MDG and EDG are NOT always the same, however they CAN be
- Ex → H2O:
- MDG is bent
- EDG is tetrahedron/triangular pyramid
Expanded Octet (Exceptions to Octet Rule):
- PCl5- Trigonal bipyramidal
- SO2- Bent (MDG) OR Trigonal Pyramidal (EDG)
- SiF6- Octahedral
Multiple Bonds:
- Multiple bonds (double and triple bonds) count as one domain.
- A double bond represents one domain, but it contains two electron pairs.
- A triple bond is one domain composed of three electron pairs.
- Multiple bonds contain more than one pair of electrons; they exert a greater repulsion than single bonds.
- The increased repulsion causes the bond angles in the molecule to deviate from predicted values
Sample Question:
- Suggest why the $O_{3}$ (Ozone) bond angle (117o) is smaller than the bond angle (119°) of $SO_{2}$ (Sulfur Dioxide)?
- The extra repulsion in Sulfur Dioxide from the lone pairs and double bonds causes the angle to go to 119o. The structure of Ozone contains resonance and one lone pair of electrons, causing it to go to 117°
- $SO_{2}$ has bond pair-bond pair repulsion, whereas $O_{3}$ has lone pair-bond pair repulsion. Lone pair-bond pair repulsion is stronger than bond pair- bond pair repulsion which is why $O_{3}$ has an angle of 117° and $SO_{2}$ has an angle of 119°
S2.2.5-6 – Bond Polarity and Molecular Geometry
Periodic Trends (review):
1. Electronegativity → Ability to attract electrons
- Increase left to right
- Decrease top to bottom
- Exception: noble gases (Group 18), lanthanides, and actinides (the last two are the small row of elements at the very bottom of the periodic table)
2. Ionization Energy → Energy needed to remove an electron
- Increase left to right
- Decrease top to bottom
3. Ionic Radius → Distance between a nucleus to which it has influence on its electron cloud
Ionic radius reverse to ions while atomic radius refers to neutral atoms
- Decrease left to right
- Increase top to bottom
- Exceptions: Noble Gases
Bond Polarity:
- Results from the difference in the electronegativities of the bonded atoms.
- If the difference in electronegativity between two bonded atoms (x) is, greater than 1.5 and less than 1.8 makes it polar covalent
- 5 < x > 1.8 = Polar Covalent bond
- Separation of charge between two non-identical bonded atoms is called a dipole moment.
- The value of the dipole moment is often reported in the unit debye, D.
Molecular Polarity:
- The polarity of a molecule depends upon:
- The polar bonds it contains.
- The shape of the molecule.
- If the bonds are equally polar and arranged symmetrically, then they cancel each other out and are non-polar.
- If the molecule contains bonds of different polarities or the bonds are not arranged symmetrically then the molecule will be polar.
- You can usually tell by the shape and lone pairs of electrons if the molecule is polar or not.
S2.2.7 – Covalent Network Structure
*Covalent Network Structure:
- Covalent bonds give rise to two different types of structure:
- Molecular Network
- Consist of discrete groups of covalently bonded atoms called molecules
- Covalent Network (or Giant Structure)
- Contain atoms that are held together by covalent bonds in a continuous three dimensional lattice. Ex. (Silicon, Silicon Dioxide and carbon allotropes)
- The structures are very different in terms of their properties.
*Allotropes of Carbon:
- Allotropes are different forms of an element in the same physical state.
- Different bonding within the structures gives rise to different physical properties.
- Carbon has four allotropes.
1.Graphite
- $sp^{2}$ hybridized covalently bonded to 3 others forming hexagons in parallel layers with bond angles of 120 degrees. TRIGONAL PLANAR
- The layers are held together by van der Waals forces so they can slide over each other.
- Density is 2.26 g/cm3
- Contains one non-bonded, delocalized electron per atom so graphite conducts electricity due to the movement of these electrons.
- Not a good heat conductor
- Very high melting point, most stable allotrope
- Gray solid
- Used as a lubricant and in pencils
- Can conduct electricity
2. Diamond
- Each C atom is sp3 hybridized covalently bonded to 4 others tetrahedrally in a regular repeating pattern with bond angles of 109.5 degrees. TETRAHEDRAL
- It is the hardest known natural substance.
- Density is 3.51 g/cm3
- All electrons are bonded so it does not conduct electricity. Thus, it has no free electrons.
- Weakest conductor of electricity
- Does conduct heat better than metals.
- Very high melting point, brittle
- Lustrous crystal
- Used in jewelry and tools
3.Fullerene, C60
- sp2 hybridized covalently bonded in a sphere of 60 carbon atoms consisting of 12 pentagons and 20 hexagons.
- The structure is a closed spherical cage in which each carbon atom is bonded to 3 others.
- Density is 1.726 g/cm3
- It easily accepts electrons to form negative ions so it is a semiconductor at normal temp and pressure due to some electron mobility.
- Very low heat conductivity
- Low melting point
- Yellow crystalline solid – soluble in benzene
- Related forms are used to make nanotubes for the electronics industry.
4.Graphene
- sp2 hybridized covalently bonded to three other carbons forming hexagons with bond angles of 120. HEXAGONAL
- The structure is a two dimensional single layer described as a honeycomb or chicken wire structure.
- Density is 1.5 g/cm3
- Contains one non-bonded, delocalized electron per atom so graphene conducts electricity due to the movement of these electrons.
- Strongest conductors of electricity
- Excellent heat conductor – better than diamonds
- Very high melting point, thinnest material to exist, stronger than steel
- Almost completely transparent
S2.2.8 – Intermolecular Forces
Intermolecular Forces:
- These are forces between molecules.
- These are forces within the molecule such as covalent, ionic and metallic bonding.
- The strength determines the volatility of a substance.
- The stronger the forces, the higher the melting and boiling points.
- Intermolecular forces are also the forces of attraction that exist between molecules
- The strength of these forces determine:
- The state of matter: solid, liquid, or gas
- The melting and boiling points of the compound.
- The solubilities of one substance in another.
- The strength of these forces determine:
Types of Intermolecular Forces:
- Hydrogen bonding
- Van der Waals forces
- Dipole to dipole interactions
- Dispersion forces
- Dipole induced interaction
London Dispersion Forces:
- Very weak forces of attraction between molecules
- They result from:
- Momentary dipoles occurring due to uneven electron distributions in neighboring molecules as they approach one another
- The weak residual attraction of the nuclei in one molecule for the electrons in a neighboring molecule.
Effect of London Dispersion Forces:
- The attractive forces that cause non-polar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently.
- Phase changes occur when molecules are sufficiently close and dispersion forces are sufficiently strong to hold molecules together.
Dispersion Forces:
- More electrons present in molecule → Stronger dispersion forces
- The only type of intermolecular force that operates between nonpolar molecules
- Weakest of the intermolecular forces
- Typically only 0.1% to 1% as strong as covalent bonds between atoms in a molecule
- Example of Dipersion forces:
- H2, CO2, Cl2
Dipole-Induced Dipole Forces:
- Type of related intermolecular force occurring between a polar molecule and a nearby non-polar molecule.
- Presence of a permanent dipole in a POLAR molecule will cause the formation of a temporary dipole in another neighboring NON-POLAR molecule.
- For example, this type of intermolecular force attracts nonpolar oxygen molecules, O2, to polar water molecules.
- Dipole–induced dipole forces are weak, which explains why the aqueous solubility of oxygen is relatively low.
Dipole-Dipole Forces:
- Occur between molecules that have permanent net dipoles. (polar molecules).
- The partial positive charge on one molecule is electrostatically attracted to the partial negative charge on a neighbouring molecule.
- Example:
- SCl2, PCl3, CH3Cl
Hydrogen Bonding:
- Strongest of all intermolecular forces: 1/10 of the strength of a covalent bond
- Substances with significant hydrogen bondings have a higher than normal melting/boiling point.
- Occurs between polar covalent molecules that possess a hydrogen atom that is bonded to an extremely electronegative element; specifically – N, O, and F.
- The weak attractions that result from hydrogen bonding cause molecules to stick together.
- As a result molecules with significant hydrogen bonding have higher melting/boiling points than they would otherwise have.
S2.2.9 – Properties of Covalent Compounds
Relative Strength of Intermolecular Forces:
- London (dispersion) forces < dipole–dipole forces < hydrogen bonding
Volatility:
- Determined by the strength of the intermolecular forces between the molecules.
- Covalent bonds with weak intermolecular forces are more volatile, meaning they evaporate faster.
- Molecules will easily break apart and become a gas.
- Covalent bonds with intermolecular forces less volatile will evaporate slowly.
- Substances with covalent network structures are solids at room temp. and pressure.
- Vaporizing them will take a lot of energy due to the strong covalent bonds holding it
- This makes them non-volatile with high melting/boiling points.
- To vaporize molecular structures, intermolecular forces between molecules must be overcome.
- Because intermolecular forces tend to be weak, the energy req. to overcome is low.
- Meaning, molecular substances are generally volatile.
- Volatility of a molecule is inversely proportional to its size for only LONDON DISPERSION FORCES.
- Small molecules → higher volatility →Bigger molecules → lower volatility
- Example: O2 (Oxygen) is more volatile than I2 (Iodine)
Electrical Conductivity:
- Covalent compounds are generally not electrically conductive.
- Because they lack free electrons to move around.
- Free electrons are essential for electrical conduction.
- Exception: Graphite → covalent compound but it is electrically conductive
- Why? Because the carbon atoms are arranged in a way that causes the electrons to move between them easily.
- Similarly, Silicon is a semiconductor.
- Its intermediate conductivity (small energy gap between valence and conduction bands) places it between conductors and insulators.
- Which is why Silicon is used in solar panels.
S2.2.10 – Chromatography
Relative Strength of Intermolecular Forces:
- Chromatography is a separation technique that enables the separation of mixtures
- 2 types:
- Paper chromatography
- Thin-layer chromatography
Key components of chromatography:
Mobile phase: A substance that the molecules can move through(always a liquid or a gas) In the case of paper chromatography, it is the solvent that you begin with.
Stationary phase: The substance or material that the molecules can’t move in, (usually a solid but can be a really thick liquid) in this case it is the paper
Principles of separation
- The stationary phase constraints components in a mixture, causing them to move more slowly than the mobile phase.
- The movement of the components in the mobile phase is controlled by the significance of their interactions with the mobile and/or stationary phases.
- Differences in factors like solubility in the mobile phase and affinities for the stationary phase cause certain components to move faster, aiding in the separation of mixture components. These affinities are due to intermolecular forces.
The stationary phase is the paper and the mobile phase is the solvent
The stationary phase is a rectangular plate made of glass or metal coated with silica or alumina. The mobile phase is a non-polar organic solvent.
Advantages of both-
Paper chromatography:
- Cheap
- Little preparation
- More efficient for polar and water-soluble compounds
- Easy to store and handle
Thin layer chromatography:
- Faster
- Detects smaller amounts
- Better separation of less soluble compounds
- Corrosive material can be used
- A wide range of stationary phase is available
S2.2.11 – [HL] Resonance
Resonance:
- Used to describe structures with multiple ways of depicting the same molecule.
- If a double bond can be put in multiple positions, it’s expected to draw the resonance structures of it.
- Electrons are delocalized in areas that have double bonds and are spread equally in each bonding position.
- Bond strength and length are in between single and double bonds
- Resonance structures allow the depiction of all possible double bond positions
- True resonance structure is an intermediate form → Resonance Hybrid
- Double arrows are placed between ALL resonance structures
Delocalization:
- Occurs when electrons are shared by more than two atoms in a molecule or ion, as opposed to being localized between a pair of atoms.
- Resonance structures can be described as a single structure using the concept of delocalization.
- Localized → only two atoms sharing a pair of electrons
S2.2.12 – [HL] Benzene and Resonance
S2.2.13 – [HL] Expanded Octets
Expanded Octet:
- Some atoms are an exception to the octet rule. (Have more than 8 valence electrons)
- To do this, the atom needs easily available d orbital electrons such as Phosphorus and Sulfur which is in period 3.
- Drawing lewis structures for atoms with expanded octets is the same as normal. However, in this case the central atom has more than 8 electrons.
S2.2.14 – [HL] Formal Charge
Formal Charge:
- The possibility of having multiple valid lewis structures for one compound.
- Most often seen with molecules that have expanded octets.
- Example: Sulfur Dioxide (SO2)
Which structure is more stable?
- The concept of formal charge treats covalent bonds as if they were purely covalent with equal electron distribution and ignores electronegativity.
- Count how many electrons belong to each atom in the Lewis structure
- Compare this with the number of valence electrons in the non-bonded atom.
- The difference is the formal charge.
FC = V-(1/2B + L) → Low formal charge → More stable structure
V – number of valence electrons
B – number of bonding electrons
L – number of lone-paired (non-bonding) electrons
S2.2.15 – [HL] Sigma bonds (Ω) and Pi bonds (π)
Summary:
- Triple bond:1 sigma bond and 2 pi bonds
- Double bond:1 sigma bond and 1 pi bond
- Single bond:1 sigma bond
Sigma bonds (Ω):
- Form from overlapping along the bond axis.
- All single covalent bonds are sigma bonds
- Example:
- s and s (H2)
- s and p (HCl)
- p and p (Cl2)
- Hybrid orbitals and s
- Hybrid orbitals and hybrid orbitals
Pi bonds (π):
- Weaker than sigma bonds because their electron density is farther away from the positive nucleus.
- All double covalent bonds contain one pi and one sigma bond.
- All triple covalent bonds contain two pi and one sigma bond.
- The following form pi bonds:
- p and p sideways
- The double bonds break more readily and are more reactive than those with only sigma bonds.
S2.2.16 – [HL] Hybridization
Hybridization:
- The concept of mixing atomic orbitals to form new hybrid orbitals for bonding.
- 4 bonding domains – sp³
- 3 bonding domains – sp²
- 2 bonding domains – sp
- The rest is extra information but up to is basically the gist of this
Sp³ Hybridization:
- CH4, is a classic example of sp³
- This type of hybridization occurs when the three “p” orbitals and one “s” orbital hybridize to form four identical sigma bonds.
- The shape is tetrahedral and has bond angles of 109.5 degrees.
Sp² Hybridization:
- When carbon forms a double bond as in ethene, it undergoes sp2
- This type of hybridization occurs when the two “p” orbitals and one “s” orbital hybridize to form three hybrid orbitals and leave one unhybridized “p” orbital.
- The shape is trigonal planar with bond angles of 120 degrees
- The unhybridized “p” orbitals overlap sideways forming a pi bond.
Sp Hybridization:
- When carbon forms a triple bond as in ethyne, it undergoes sp Hybridization.
- This type of hybridization occurs when the one “p” orbitals and one “s” orbital hybridize to form two hybrid orbitals and leave two unhybridized “p” orbitals.
- The shape is linear with bond angles of 180 degrees.
- The unhybridized “p” orbitals overlap sideways forming two pi bonds.