Lewis Formulas | General Chemistry 1
Covalent bonds result from electron sharing between two atoms. Electrons are shared to allow the atoms to attain noble-gas configurations and reach an octet in the valence shell.
Cl: Z = 17 ⇒ 1s22s22p63s23p5 ⇒ 7 valence electrons
If 2 chlorine atoms each share an electron:
⇒ they have 8 electrons in their outer shell
⇒ they reach the electron configuration of argon
⇒ formation of Cl2
A covalent bond is formed by 2 shared electrons.
Bond order: number of chemical bonds between a pair of atoms.
Cl2 has one chemical bond between the 2 chorines
⇒ bond order = 1
Bond length: distance between 2 covalently bonded nuclei
triple bond < double bond < single bond
Bond strength: energy required to break a bond
triple bond > double bond > single bond
(when comparing similar bonds ⇒ bond length/atomic radius becomes a factor)
Lewis structures show us how the electrons are arranged around atoms in a molecule:
dot ⇒ one non-bonding electron
pair of dots ⇒ lone electron pair (lone pair)
line ⇒ two shared electrons (bond)
We always try to satisfy the octet rule (or duet rule for hydrogen) when writing Lewis structures.
Lewis structure of NH3:
N: 5 valence electrons
⇒ needs 3 shared electrons ⇒ 3 covalent bonds
H: 1 valence electron
⇒ needs 1 shared electron ⇒ 1 covalent bond
Lewis structure of CO2:
C: 4 valence electrons
⇒ needs 4 shared electrons ⇒ 4 covalent bonds
O: 6 valence electrons
⇒ needs 2 shared electrons ⇒ 2 covalent bonds
How to write Lewis structure:
1) Add together the valence electrons for all of the atoms in the compound.
Add or subtract electrons as necessary if you have an ion.
2) Determine which atom is your central atom (usually the first one in the formula unless the first atom is H).
3) Represent a covalent bond by placing a line between the atoms that are assumed to be bounded to each other.
4) Arrange the remaining valence electrons as lone pair to satisfy the octet rule.
Lewis structure of CCl4:
1) Number of electrons = valence electrons + charge
= 1 x 4 (1 carbon) + 4 x 7 (4 chlorines) + 0 (charge) = 32
(⇒ 16 pairs of electrons ⇒ 16 covalent bonds or lone pairs).
2) Cl wants to form only one covalent bond ⇒ C is probably the central atom.
8 electrons are used to create the covalent bonds.
C follows the octet rule.
There are 24 electrons left ⇒ 12 pairs of electrons ⇒ 12 lone pairs (C already follows the octet rule) ⇒ 3 lone pairs on each chlorine
Cl follow the octet rule.
Formal Charge = (nb of valence electrons in free atom) - (nb of valence electrons in bound atom)
Valence electrons in bound atom = nb of covalent bonds + 2 x nb of lone-pair on the atom
What are the formal charges of CH3NO2?
N: nb of valence electrons (VE) in free atom = 5 (Z =5)
nb of valence electrons in bound atom = 4 (4 covalent bonds)
Formal charge = 5-4 = +1
O: nb of VE in free atom = 6 (Z =6)
nb of VE in bound atom = 6 (2 covalent bonds + 2 lone pairs)
Formal charge = 6-6 = 0
O: nb of VE in free atom = 6 (Z =6)
nb of VE in bound atom = 7 (1 covalent bond + 3 lone pairs)
Formal charge = 6-7 = -1
Resonance hybrid is a superposition of Lewis structures
⇒ attempt to represent a real structure that is a combination of several extreme possibilities.
Resonance forms differ only in the placement of their π or non-bonding electrons and obey rules of valency.
Different resonance forms of a substance are not equivalent: one structure usually contributes to the hybrid structure more than others.
Selecting the best resonance structure:
- lower formal charges (positive or negative) are preferable to higher ones.
- formal charges on adjacent atoms are not desirable.
- a more negative formal charger should reside on a more electronegative atom.
Free radical: species with one or more unpaired electrons.
They are formed by homolytic cleavage.
CH3 has one unpaired electron on the carbon
NO has one unpaired electron on the nitrogen
⇒ these are free radicals
Expanded Valence Shells
Exceptions to the octet rule:
- odd number of electrons (free radical)
- incomplete octet (B and Be are prime candidates for this)
- more than 8 electrons on the central atom ⇒ expanded valence shell.
The last one is the largest class of exceptions and can happen only for atoms with a valence shell n 3.
P has 10 electrons in its outer shell.
Electronegativity: ability of one atom in a molecule to attract electrons to itself
⇒ electrons in a bond are pulled toward the atom with the highest electronegativity
Electronegativity increases across a period and towards the top of a group
⇒ same trend as electron affinity.
Bond energies (BE): energy required to break bonds between two atoms.
Enthalpy of reaction ΔH0: difference in enthalpy between products and reactants (in kJ.mol-1)
= sum BE of reactants - sum BE of products
ΔH0 < 0 ⇒ the reaction is exothermic (heat out)
ΔH0 > 0 ⇒ the reaction is endothermic (heat in)
ΔH0 of the following reaction:
CH4 (g) + 3 Cl2 (g) → CHCl3 (g) + 3 HCl (g)
Data: BE (C-H) = 413 kJ.mol-1; BE (Cl-Cl) = 243 kJ.mol-1; BE (C-Cl) = 339 kJ.mol-1; BE (H-Cl) = 427 kJ.mol-1
ΔH0 = 4 BE (C-H) + 3 BE (Cl-Cl) – BE (C-H) – 3 BE (C-Cl) – 3 BE (H-Cl)
ΔH0 = -330 kJ.mol-1
ΔH0 < 0 ⇒ the reaction is exothermic
Polar Bonds & Dipole Moments
Polar covalent bonds: covalent bonds in which there are unequal sharing of electrons between two atoms.
Difference in electronegativity is a gauge of bond polarity:
the greater the electronegativity difference, the more polar the bond.
If the difference of electronegativity is too high ⇒ transfer of electrons ⇒ ionic bonds.
H2 ⇒ same atom ⇒ no difference of electronegativity ⇒ nonpolar bond
HCl ⇒ Electronegativity: H = 2.1, Cl = 3.0 ⇒ difference of electronegativity ⇒ polar bond
NaCl ⇒ Electronegativity: H = 0.9, Cl = 3.0 ⇒ very high difference of electronegativity ⇒ ionic bond
Dipole moment: occurs whenever there is a separation of positive and negative charges.
It is represented as an arrow pointing along the bond from δ+ to δ-.