Chemistry Notes
Chemical Bonding
Atomic Combinations
1. A chemical bond is a net electrostatic force between atoms, ions or molecules that enables
the formation of chemical compounds.
2. The bond may result from the electrostatic force of attraction between oppositely charged
ions (ionic bonds) or through the sharing of electrons (covalent bonds).
3. Atoms react to achieve a stable noble electron structure and a lower potential energy.
4. Intermolecular force – a weak force of attraction between molecules, ions, or atoms of
noble gases.
5. Intramolecular force (interatomic bond) – a strong force of attraction between atoms within
a molecule.
Ionic bond
1. Is the transfer of electrons and subsequent electrostatic attraction.
2. Metals react by losing electrons and non-metals react by gaining electrons.
3. Positive ions (cations) and negative ions (anions) attract each other with strong
electrostatic/Coulombic forces.
4. In an ionic bond
a. Smallest particle is an ion
b. Ionic bonds form a crystal lattice in the solid phase
5. Ionic substances don’t conduct electricity in a solid state.
Metallic bonding
1. Is the attraction between a positive kernel and a sea of delocalised electrons.
2. Metallic bonds:
a. Bonds between atoms of metals.
b. Atoms are closely packed so the outermost orbitals overlap.
c. The valence electrons become delocalised and can move freely in empty orbitals.
d. Positive core forms a compact crystal lattice.
e. Smallest particle is a positive ion.
Covalent bonding
1. Is the sharing of at least one pair of electrons by two non-metal atoms.
2. Properties of covalent bonds
a. They arise between two non-metal atoms when their outermost energy levels
overlap.
b. Each atom contributes an electron to be shared.
c. Covalent bonds result in the formation of molecules.
3. A covalent bond forms when the valence orbitals of atoms overlap and share an electron
pair between the nuclei.
4. A single bond is one electron pair of shared electrons. Also known as bonding/shared pair.
5. A lone pair is an unshared electron pair.
6. Sometimes atoms share more than one pair electrons to make a double/triple bond.
1
,The Lewis Diagram
1. For each atom, determine the number of valence electrons in the atoms and represent them
using dots/crosses.
Valence electrons
1. Electrons in the highest energy level of the atom.
Valency electrons
1. The number of electrons an atom will accept, donate, or share in order to achieve a stable
electronic configuration.
Dative covalent bonds
1. Is a lone pair of electrons that are shared with an empty orbital of another atom or an ion.
2. Is formed by the overlapping of a filled orbital containing a lone pair with an empty orbital of
another atom or ion.
Non-polar and polar covalent bonds
1. Non-polar covalent bond (pure covalent bond) – is the equal sharing of electrons.
2. Polar covalent bonds – is the unequal sharing of electrons leading to a dipole forming.
3. Polarity of molecules affects properties such as solubility, melting and boiling points.
4. Electronegativity – is the measure of the tendency of an atom to attract a bonding pair of
electrons.
5. Formula for calculating electronegativity differences
a. 𝛥𝐸𝑁 = 𝐺𝑟𝑒𝑎𝑡𝑒𝑟 𝑣𝑎𝑙𝑢𝑒 − 𝑙𝑜𝑤𝑒𝑟 𝑣𝑎𝑙𝑢𝑒
6. To whether a bond is polar/non-polar/ionic
Electronegativity difference Bond polarity
0 Non-polar (pure) covalent bond
0 < ∆EN ≤1,7 Polar covalent bond
∆EN >1,7 Ionic bond
7. A non-polar (pure) covalent bond occurs between two identical non-metal atoms or
between atoms with the same electronegativity.
8. A polar covalent bond occurs if two different non-metal atoms bond then the shared
electron pair will be attracted more strongly to the atom with the higher electronegativity.
9. A polar molecule has one end with a slightly positive charge and the other with a slightly
negative charge.
10. A polar molecule:
a. Has dipoles
b. Partially positive and negative ends
c. Has a net dipole not equal to zero 9the dipole moment of each bond doesn’t cancel
each other out).
d. Has an asymmetrical charge distribution
11. A polar bond has a dipole moment ( ) which shows the direction in which the
bonding pair is attracted.
12. A dipole moment is a vector quantity, which means it has magnitude and direction.
13. A non-polar molecule is where the charge is equally spread across the molecule.
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, a. Has a net dipole equal to zero.
b. Bond dipoles cancel each other out (i.e. the vector addition of the dipoles equal
zero) and the overall molecule has a zero dipole moment.
c. Has a symmetrical charge distribution.
Dipole
1. Is a molecule that has two poles as regions with opposite charges.
2. Is represented by a dipole arrow pointing towards the more negative end.
Shapes of molecules and VSEPR theory
1. Valence Shell Electron Pair Repulsion (VSEPR) theory is a model in chemistry which is used to
predict the shapes of simple molecules.
2. Electron pairs (lone and bond pairs) around the central atom repel each other so that the
bond angle has a maximum value, and the potential energy is at its lowest.
3. Bonding and lone pairs arrange themselves so that they are as far apart as possible.
4. Molecular shapes are determined by the repulsions between electron pairs present in the
valence shell of the central atom.
5. The assumption of the VSEPR is that the molecule adapts the geometrical shape that
minimizes the repulsive force among a given number of electron pairs.
6. The shape of the molecule depends on the number of bonding electron groups (or atoms
bonded to the central atom) and the number of lone pairs on the central atom.
7. A represents the central atom, X represents the terminal atoms.
8. Five ideal shapes are found when there are no lone pairs. ONLY bond pairs.
9. The five ideal shapes are:
a. Linear – (AX2) TWO
b. Trigonal planar – (AX3) THREE
c. Tetrahedral – (AX4) FOUR
d. Trigonal bipyramidal – (AX5) FIVE
e. Octahedral – (AX6) SIX
10. Non ideal molecular shapes (the central atom has bond pairs and lone pairs)
a. Molecules with bond and lone pairs around their central atom can’t have ideal
molecular shapes.
Number of bonding electron Number of lone pairs Geometrical Shape General formula
pairs
2 2 bent or angular A X2E2
3 1 trigonal pyramidal A X3E
Definitions
1. Intramolecular Bonds
a. A covalent bond – is a sharing of at least one pair of electrons by two non-metal
atoms.
b. Non-polar covalent bond/pure covalent – is an equal sharing of electrons.
c. Polar covalent bond – is the unequal sharing of electrons leading to a dipole forming
(as result of electronegativity difference).
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Chemical Bonding
Atomic Combinations
1. A chemical bond is a net electrostatic force between atoms, ions or molecules that enables
the formation of chemical compounds.
2. The bond may result from the electrostatic force of attraction between oppositely charged
ions (ionic bonds) or through the sharing of electrons (covalent bonds).
3. Atoms react to achieve a stable noble electron structure and a lower potential energy.
4. Intermolecular force – a weak force of attraction between molecules, ions, or atoms of
noble gases.
5. Intramolecular force (interatomic bond) – a strong force of attraction between atoms within
a molecule.
Ionic bond
1. Is the transfer of electrons and subsequent electrostatic attraction.
2. Metals react by losing electrons and non-metals react by gaining electrons.
3. Positive ions (cations) and negative ions (anions) attract each other with strong
electrostatic/Coulombic forces.
4. In an ionic bond
a. Smallest particle is an ion
b. Ionic bonds form a crystal lattice in the solid phase
5. Ionic substances don’t conduct electricity in a solid state.
Metallic bonding
1. Is the attraction between a positive kernel and a sea of delocalised electrons.
2. Metallic bonds:
a. Bonds between atoms of metals.
b. Atoms are closely packed so the outermost orbitals overlap.
c. The valence electrons become delocalised and can move freely in empty orbitals.
d. Positive core forms a compact crystal lattice.
e. Smallest particle is a positive ion.
Covalent bonding
1. Is the sharing of at least one pair of electrons by two non-metal atoms.
2. Properties of covalent bonds
a. They arise between two non-metal atoms when their outermost energy levels
overlap.
b. Each atom contributes an electron to be shared.
c. Covalent bonds result in the formation of molecules.
3. A covalent bond forms when the valence orbitals of atoms overlap and share an electron
pair between the nuclei.
4. A single bond is one electron pair of shared electrons. Also known as bonding/shared pair.
5. A lone pair is an unshared electron pair.
6. Sometimes atoms share more than one pair electrons to make a double/triple bond.
1
,The Lewis Diagram
1. For each atom, determine the number of valence electrons in the atoms and represent them
using dots/crosses.
Valence electrons
1. Electrons in the highest energy level of the atom.
Valency electrons
1. The number of electrons an atom will accept, donate, or share in order to achieve a stable
electronic configuration.
Dative covalent bonds
1. Is a lone pair of electrons that are shared with an empty orbital of another atom or an ion.
2. Is formed by the overlapping of a filled orbital containing a lone pair with an empty orbital of
another atom or ion.
Non-polar and polar covalent bonds
1. Non-polar covalent bond (pure covalent bond) – is the equal sharing of electrons.
2. Polar covalent bonds – is the unequal sharing of electrons leading to a dipole forming.
3. Polarity of molecules affects properties such as solubility, melting and boiling points.
4. Electronegativity – is the measure of the tendency of an atom to attract a bonding pair of
electrons.
5. Formula for calculating electronegativity differences
a. 𝛥𝐸𝑁 = 𝐺𝑟𝑒𝑎𝑡𝑒𝑟 𝑣𝑎𝑙𝑢𝑒 − 𝑙𝑜𝑤𝑒𝑟 𝑣𝑎𝑙𝑢𝑒
6. To whether a bond is polar/non-polar/ionic
Electronegativity difference Bond polarity
0 Non-polar (pure) covalent bond
0 < ∆EN ≤1,7 Polar covalent bond
∆EN >1,7 Ionic bond
7. A non-polar (pure) covalent bond occurs between two identical non-metal atoms or
between atoms with the same electronegativity.
8. A polar covalent bond occurs if two different non-metal atoms bond then the shared
electron pair will be attracted more strongly to the atom with the higher electronegativity.
9. A polar molecule has one end with a slightly positive charge and the other with a slightly
negative charge.
10. A polar molecule:
a. Has dipoles
b. Partially positive and negative ends
c. Has a net dipole not equal to zero 9the dipole moment of each bond doesn’t cancel
each other out).
d. Has an asymmetrical charge distribution
11. A polar bond has a dipole moment ( ) which shows the direction in which the
bonding pair is attracted.
12. A dipole moment is a vector quantity, which means it has magnitude and direction.
13. A non-polar molecule is where the charge is equally spread across the molecule.
2
, a. Has a net dipole equal to zero.
b. Bond dipoles cancel each other out (i.e. the vector addition of the dipoles equal
zero) and the overall molecule has a zero dipole moment.
c. Has a symmetrical charge distribution.
Dipole
1. Is a molecule that has two poles as regions with opposite charges.
2. Is represented by a dipole arrow pointing towards the more negative end.
Shapes of molecules and VSEPR theory
1. Valence Shell Electron Pair Repulsion (VSEPR) theory is a model in chemistry which is used to
predict the shapes of simple molecules.
2. Electron pairs (lone and bond pairs) around the central atom repel each other so that the
bond angle has a maximum value, and the potential energy is at its lowest.
3. Bonding and lone pairs arrange themselves so that they are as far apart as possible.
4. Molecular shapes are determined by the repulsions between electron pairs present in the
valence shell of the central atom.
5. The assumption of the VSEPR is that the molecule adapts the geometrical shape that
minimizes the repulsive force among a given number of electron pairs.
6. The shape of the molecule depends on the number of bonding electron groups (or atoms
bonded to the central atom) and the number of lone pairs on the central atom.
7. A represents the central atom, X represents the terminal atoms.
8. Five ideal shapes are found when there are no lone pairs. ONLY bond pairs.
9. The five ideal shapes are:
a. Linear – (AX2) TWO
b. Trigonal planar – (AX3) THREE
c. Tetrahedral – (AX4) FOUR
d. Trigonal bipyramidal – (AX5) FIVE
e. Octahedral – (AX6) SIX
10. Non ideal molecular shapes (the central atom has bond pairs and lone pairs)
a. Molecules with bond and lone pairs around their central atom can’t have ideal
molecular shapes.
Number of bonding electron Number of lone pairs Geometrical Shape General formula
pairs
2 2 bent or angular A X2E2
3 1 trigonal pyramidal A X3E
Definitions
1. Intramolecular Bonds
a. A covalent bond – is a sharing of at least one pair of electrons by two non-metal
atoms.
b. Non-polar covalent bond/pure covalent – is an equal sharing of electrons.
c. Polar covalent bond – is the unequal sharing of electrons leading to a dipole forming
(as result of electronegativity difference).
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