Chemical Bonding and Lewis Structure Flashcards
What are the two main ways of creating a chemical bond:
Electron sharing - covalent bonds
Conplete electron transfer - ionic bonds
What is metallic bonding
Postive metal ions held together by a ‘sea’ of declocalised electrons
What are ionic bonds
Ionic bonds are formed between atoms with low ionisation energies and those with high electron affinities
Why is covalent bonding characteristic of elements in the middle of the Periodic table
Complete electron transfer is not practical
When the two atoms forming a covalent bond are the same, this is called
A homonuclear bond
The electrons will be equally shared
If the atoms forming a covalent bond are different, this results in a
Heteronuclear bond
Sharing will be unequal
In a heteronuclear bond, sharing of electron density is unequal
What determines which atom get more of the electrons
It depends on the electronegativity of each of the atoms
What is electronegativity
The ability of an atom in a molecule to attract electron density to itself
How can we usually determine the nature of the bond between two atoms
By considering the electronegativity of each of the atoms
If there is no electronegativity difference, the bond formed is likely to be
‘pure’ covalent bond (X-X)
for example C-H
If there is a small to moderate electronegativity difference, the bond formed is likely to be…
Polar covalent bonds (Xδ+ - Yδ-)
e.g. Cδ+ - Oδ-
If there is a very large electronegativity difference, the bond formed is likely to be…
An ionic bond (X⁺ - Y¯)
e.g. (K⁺ - Br¯)
What do lewis structures allow us to do
determine the connectivity and formal charges of the atom in a structure (BUT not the shape)
What are the basic rules of Lewis Structure
- Every atom tried to achieve an octet of electrons (or two for hydrogen)
- Each pair of shared electrons (symbolised by line between atoms) gives one bond
- Often no more than 4 bonds to an atom
What are the 6 key step when drawing the lewis structure of a compound
- Add up the total number of valence electrons to determine how many pairs of electrons can be used in the structure
- Decide on the central atom (usually has lowest electronegativity) and surround it with the other atoms to which it is bonding
- Join atoms together using single bonds
- Keep adding the remaining pairs of electrons to form multiple bond, IF appropriate for the atom being considered
- Add any remaining electrons as lone pairs
- Check the number of electrons in the immediate surroundings of the atoms and then assign formal charges if necessary
Example: how many valence electrons and pairs of electrons will CH₄ have
C = 4 valence electrons
H = 1 valence electron
So total 4 + 4(1) = 8 valence electrons
4 pairs
There is no general rule to distinguish between central and terminal atoms, however a few things to keep in mind:
- Lowest electronegativity atoms are often central
- Chains of the same type of atom are rare (unless in a hydrocarbon)
How would you assign a formal charge to an atom
Imagine a circle around the atom
If the number of electrons inside the circle:
= valence electrons there no charge
< valence electrons then positive charge
> valence electrons then negative charge
When drawing lewis structure, what are some exceptions worth knowing about
- Some elements (towards the LHS of the Periodic table) cannot achieve a full octet
- Elements in second and subsequent rows of the Periodic table can accommodate >8 electrons (can expand their octet)
- It is often not possible to draw a single, unambiguous Lewis structure for a given molecule or ion
Some elements (towards the LHS of the periodic table) cannot achieve a full octet
Using BF₃ as an example, explain why
Boron only has 3 valence electrons and hence would need 5 more to fill the octet
However only forms 3 bonds, because it is too small to form any more bonds
Elements in the second and subsequent rows of the periodic table can accommodate >8 electrons (can expand their octet)
How is this possible
Such elements possess d-orbitals that are close enough in energy to the p-orbital to allow use in bonding, so now >4 bonds can be formed
They are not just restricted to just one s and three p orbitals
It is often not possible to draw a single, unambiguous Lewis structure for a given molecule or ion
Using the example of a carbonate ion, why is this the case
In the carbonate ions structure, there are two single bonds and one double bond
There is 3 different oxgyens, so the double bond could exist on any of the 3 oxygens
This is resonance, where electrons are delocaised in the resonance hybrid which is lower in energy than any single component form
What is resonance
Resonance involves the blending of two or more Lewis structures, where the atoms have the same relative positions, but with different electronic arrangements
What evidence is there for resonance structures
If just a single structure e.g. for the carbonate ion where to exist, there would be one shorter double bond and two longer single bonds
Experiments show that all C-O bonds in the carbonate are the same length
This is the Lewis Structure of Dinitrogen Tetroxide
It has a N-N bond, but this is very long (weak) and the molecule readily splits into two NO₂ fragments
Why?
Repulsion between the positive charges on the nitrogen atom leads to the N-N bond being very weak, hence breaks
If the nitrogens also both have positive charges it means they are electron deficient, so the covalent bond is lacking in electron density
It is found that the O-O distance is H₂O₂ is much longer than in F₂O₂. Why is this the case
Because Fluorine is so electronegative, in the F₂O₂ case, it is pulling the electron density of the molecule out towards the terminal fluorines
The lone pair on the oxygen can become a double bound, and the bond pair between F-O being accepted by fluorine to form an F¯ ion
When this is averaged out, Oxygens bond length is somewhere between a single and a double bond
The ion (NCO)¯ is stable; the ion (CNO)¯ is known, but is very unstable. Why is this?
(NCO)¯ can form two resonance forms, with the negative charge more predominantly on oxygen rather than Nitrogen, from their relative electronegativities
(CNO)¯ has 3 formal charges - one positive and two negative. The positive charge is on the electronegative nitrogen atom
What is a Coordinate (dative) bond
Formed by one atom (Donor) ‘giving’ both electrons to another (Acceptor)
Shown by an arrow
