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Topic 4. Chemical bonding and structure


Intramolecular bonds are strong chemical bonds.

  • ionic

    • ionic structure​

    • ionic lattice structure: NaCl, LiI, MgO, Fe2O3, etc.

  • covalent

    • discrete molecules: CO2, H2O, H2, O2, P4, HClO, CH3COOH, etc.​

      • only ones that have intermolecular forces​

    • giant covalent network: diamond, graphite, etc.

  • metallic: metal cations surrounded by a sea of delocalized electrons

Molecules are chemical species that are covalently bonded and have the structure of discrete molecules.

  • they have intermolecular forces

  • they can exist independently (unlike lattice structures & giant covalent networks which are repeating units)

Compounds are 2 or more different atoms bonded together.

Melting/boiling point rankings

  1. giant covalent molecules: there are a lot of covalent bonds to break

  2. ionic lattice structures / metallic structures

  3. discrete covalent molecules: only have to break intermolecular forces → very volatile

Conditions for conducting electricity

  • either have free-moving / mobile ELECTRONS

  • or have free moving / mobile IONS

    • rarely in solid form. look for (aq) or (l)​

    • ALL molten ionic compounds and ionic compounds dissolved in water can conduct electricity

    • giant covalent networks can't EXCEPT for graphite & graphene, which have seas of delocalized electrons

4.1 Ionic bonding and structure

Ionic bonding is the electrostatic attraction between oppositely charged ions (metal cation after losing electrons and a nonmetal anion after gaining electrons).

  • ions with a + charge are called cations

  • ions with a – charge are called anions

Forms an ionic lattice structure or crystal lattice (regular repeating patterns)

  • high melting/boiling point because its hard to break the bonds

    • ​low volatility (the tendency of a substance to vaporize)

  • (when molten) can conduct electricity

  • brittle

  • non-directional bond (force is equally distributed)

The melting of ionic compounds → breaking ionic bonds requires high levels of energy (strong electrostatic attraction) → high MP/BP

Ionic bond formation

  • depends on the difference in electronegativity between the two elements

  • the table roughly outlines the general trend, but there can be exceptions

* reminder, although HCl is H+ and Cl–, it is still covalent (nonmetal-nonmetal).

  • written as H+ and Cl– because HCl completely dissociates

Coordination number: the number of atoms a central atom is surrounded by.

Solubility of ionic compounds


  • All nitrates

  • All sodium, potassium, and ammonium compounds

  • All sulfates EXCEPT barium, lead, and calcium

  • All chlorides, bromides, and iodides EXCEPT silver and lead


  • All carbonates EXCEPT Group 1 and ammonium carbonates

    • Na2CO3, K2CO3, K2CO3, Li2CO3, etc. are soluble​

  • All hydroxides EXCEPT Group 1 and ammonium carbonates

    • NaOH, LiOH, KOH, RbOH, etc. are soluble​

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From Ms. Fu's powerpoint

4.2 Covalent bonding

Covalent bonding is the electrostatic attraction between a shared pair of electrons between two non-metals

  • these electrons are called bonding electron pairs

  • electrons not involved in bonding are called non-bonding electron pairs

Lewis diagrams & octet rule:

  • whichever atom is short of more electrons to satisfy the octet rule → the central atom

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Bond strengths:

Single < Double < Tiple

Bond lengths:

Single > Double > Triple

The strong force pulls the atoms closer and makes the bond shorter.

Dative / coordinate / dative covalent bond: a covalent bond in which both electrons come from the same atom.

Bond polarity

difference in electronegativity → bond polarity, + asymmetry of domain arrangement → molecule polarity
* bond polarity ≠ molecule polarity

when the electron density is unevenly distributed → polar molecule

atoms of different elements → polar bond

  • regardless of electronegativity difference.

    • for example, PH3: both P and H have a data booklet EN value of 2.2, but their bond is still polar

      • this is because the data booklet provides a rounded value. in reality, they are slightly different.​

atoms of the same element → non-polar bond

  • for example, Cl2

polar bonds have dipole moments: an uneven distribution of electrons

  • it is a vector quantity, meaning there is both a direction and magnitude

if there is a SYMMETRIC domain arrangement, like CH4, then the dipole movement is CANCELED.

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How to answer paper 2 questions:

Q. Is CH4 a polar molecule?

A. 1) The C-H bond is polar due to the electronegativity difference between C and H

     2) The dipole moment can be canceled out due to the symmetrical arrangement of electron domains around the central carbon atom and the 

          central carbon atom is bonded to the same element.

     3) Thus, CH4 is a non-polar molecule.

Q. Is PH3 a polar molecule?

A. 1) The P-H bond is polar due to the electronegativity difference between P and H

     2) The dipole moment generated by the polar bond is even and can not be canceled out due to the asymmetrical arrangement of electron domains

          around the central carbon atom's electron domains and the presence of a lone pair.

     3) Thus, PH3 is a polar molecule.

Symmetrical arrangements:

  • linear

  • trigonal planar

  • tetrahedral

BUT! Some are symmetrical yet polar, such as dichloromethane CH2Cl2

Asymmetrical arrangements:

  • bent / V-shaped (104.5º)

  • bent (118º)

  • trigonal pyramidal (109.5º)


Melting / boiling points

polar molecules have higher MP/BP due to the difference in electron density distribution

  • creates extra attraction between oppositely charged molecules (dipoles)

  • more energy is required to break IMFs

4.3 Covalent structures

Giant covalent networks

allotropes are made of the same element but have different structures

  • ex. diamond, graphite, fullerene

1. diamond - tetrahedral

  • extremely high MP/BP (have to break all the covalent bonds)

  • each C is covalently bonded to 4 other Cs

    • no remaining electrons → is not an electrical conductor

  • HL: bonded via sp³ hybridization 

2. graphite

  • high MP/BP (have to break all the covalent bonds)

  • each C is covalently bonded to 3 other Cs

    • 1 remaining electron is free to roam between layers → can conduct electricity

  • slippery - can be used as a lubricant

  • insoluble

  • HL: bonded via sp² hybridization 

3. graphene - pure carbon

4. silica (SiO2)

fullerene / C-60 / buckyballs

  • has 60 carbon atoms (not a G.C.N. though, it is a discrete molecule)

  • super strong

  • can conduct electricity (but limited since the free moving electrons only move on the outside)

  • conducts heat (low resistance)

  • free radical scavenger

  • HL: bonded via sp² hybridization 

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Tricky lewis diagrams to remember:

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Resonance occurs when a Lewis structure allows for the same arrangement of atoms but different (yet equally valid) arrangement of electrons

The result is a resonance hybrid. A common example is ozone O3.

  • electrons are delocalized and spread over two bonding orbitals

  • represented through dotted lines

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From Ms. Fu's powerpoint


Valence Shell Electron Pair Repulsion theory (VSEPR):

  1. all the electron domains (electron clouds, electron pairs) around central atoms repel each other

  2. molecules are most stable when the repulsive forces are at a minimum

  3. when the electron domains are arranged as far apart as possible, the repulsive forces are minimized

An electron domain is the number of bonds / electron pairs around the central atom.

  • ex. NH4+ has 4, SO2 has 3

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From Ms. Fu's powerpoint

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4.4 Intermolecular forces

Intermolecular forces (IMFs) are attractions between molecules that have temporary dipoles, permanent dipoles, or hydrogen bonds.

Volatility describes how likely a substance is to evaporate

  • basically asking how weak the IMFs are

  • if only LDF is present → high volatility

There is a phase change when IMFs are overcome.

Three main types of IMFs:

  1. London dispersion forces: interactions between discrete molecules that form temporary dipoles (induced dipoles)

  2. Dipole-dipole forces: permanent dipoles of polar molecules (polar bonds)

  3. Hydrogen bonds: (a type of dipole-dipole) forms direct bonds with H and either N, O, or F

London dispersion forces

LDF is present in all discrete molecules

  • seems to be more significant in non-polar molecules since non-polar ones have no other IMFs...

london dispersion force = dispersion force = temporary dipole-dipole = dipole induced dipole = instantaneous dipole-dipole

Increasing LDF... *can only compare LDF within the same group

  • larger radius → contains more electrons

  • heavier molecular mass → more polarization → more electrons

  • less branching → more surface area interaction

    • linear: stronger LDF​

    • spherical: weaker LDF and lower MP/BP

Dipole-dipole forces

Is when partially charged ends of molecules attract or repel each other

  • also called permanent dipoles

  • only present in POLAR molecules

  • a larger EN difference → stronger attraction

Hydrogen bonds

A hydrogen bond is an attractive force that forms between an unshared electron pair and a hydrogen atom covalently bonded to N, O, or F

  • N, O, F are very electronegative → the molecule is likely very polar

  • a special type of dipole-dipole attraction

  • the H that forms the hydrogen bond must be bonded to N, O, or F within a molecule

  • H bonds form BETWEEN molecules (intermolecular)

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Q. Why does carboxylic acid have a stronger attractive force than alcohol?

A. the red bond has a greater electrostatic attraction; carboxylic acid forms 2 hydrogen

     bonds and thus has a higher MP/BP; alcohol only forms 1 hydrogen bond

Aldehydes (left) and ketones (right) DO NOT form hydrogen bonds.

  • lower melting point

  • but aldehyde BP > ketone BP

    • aldehydes can be polarized more easily​

Alkanes, alkenes, alkynes, carboxylic acids, alcohol, etc. all can form H.B.s

Q. Why is ice less dense than water?

A. ice has a crystalline structure with hydrogen bonds; this orientation pushes atoms further apart

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4.5 Metallic bonding

Metallic bonds are the electrostatic attraction between a lattice of metal cations and a sea of delocalized electrons

The strength of the bond depends on:

  • charge of the ion (larger charge magnitude = stronger)

  • size of the ion (smaller radius = stronger)


Delocalized electrons make metals excellent conductors

  • metals lose electrons to satisfy the octet → electrons can flow freely through the metal

The lattice structure is made of many thin layers (no definite bonds) → malleable

Very high MP/BP

  • have to break strong metallic bonds

INsoluble in water

Alloys are metals made from mixtures of two or more metals and/0r other elements

  • brass = copper + zinc

  • steel = iron + copper + others

  • solder = lead + tin

  • bronze = copper + tin

  • pewter = tin + (copper, lead, antimony, bismuth)

  • amalgam = mercury + (tin, gold, silver, sodium)

Alloys make metals stronger:

  • The atoms of different sizes distort the layers of atoms in the pure metal → a greater force is required for the layers to slide over each other → therefore, the alloy is harder and stronger than the original, pure metal.


Giant covalent networks are made of atoms held together by numerous strong covalent bonds.

Metals are made of atoms; ionic compounds are made of ions; the two have very similar MPs/BPs

Discrete molecules are made of individual molecules → very weak IMFs

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