Chapter 5 - Electronegativity, bond polarity, and Lewis structures

The ability of an atom to attract bonding electrons from another atom to itself

- Smaller the EN

- Increases

- Decreases

The degree of polarity in a chemical bond between two atoms based on the EN difference between them

Polar = bond that contains a partially charged + end and a partially charged - end between atoms with different EN

nonpolar = bond that does not contain partial charges between atoms with equal EN

polar = 0.4 - 2.0 EN difference

nonpolar = 0 - 0.4 EN difference

ionic = 2.0 + EN difference

A measure of bond polarity where μ = q x r

- q = size of equal but opposite charges

- r = distance between them

- % ionic character: dipole moment if the electron was completely transferred (measured dipole moment · 100%)

1) write the chemical symbol of the elements used

2) draw brackets around the anion and draw dots to resemble its valence electrons, include a dot for the transferred electron from the cation

3) write the assigned charges next to the elements in the upper right corner, with the anions charge outside the brackets

1) Count the total number of valence electrons for the molecule

2) The first element listed in the central element (UNLESS it's hydrogen or a tri-atom)

3) Draw the central atom and disperse the others around it, connect each with a single bond ( -2 electrons for every bond)

4) Satisfy octets for outer atoms first and disperse remaining electrons as lone pairs on the central atom

5) If there's any unpaired electrons on the central atom and an outer atom, 'push' electrons to the bonding region to create double/triple bonds to make as many octets as possible

6) determine the formal charges of each atom to determine the most stable Lewis structure

- If a molecule has a charge symbol on it, add/subtract 1 electron from the total v.e for every +/- and draw brackets around the entire structure with the charge outside in the upper right corner

The charge of an atom if all the bonding electrons were shared equally

- used to help determine the most stable Lewis structure

# of valence electrons - (# of nonbonding electrons + 1/2 bonding electrons)

1) Determine the number of valence electrons

2) determine the number of nonbonding electrons and the number of bonding electrons (count 2 electrons from every bond)

3) use the formula

4) if the result is negative/positive, the atom has a -/+ charge, if the result is 0, the atom has no formal charge

- The sum of all formal charges on the individual atoms must equal the charge of the neutral molecule or ion

- Small or 0 formal charges on individual atoms are more stable/better for the structure than large formal charges on individual atoms

- If calculating formal charges results in a negative charge, the most stable structure will the the one where the formal charge is on the most electronegative atom

- structures that have an odd # of valence electrons and results in 1 lone electron and an incomplete octet

- also called free radicals or radicals

- are highly reactive because they want to lose or bond their lone electron pair

- structures that have less than 8 bonding electrons which result in incomplete octets, but creating additional bonds would create a formal charge and is thus, not the most stable structure

- most common elements are B and Al

- structures where one of the atoms has the ability to bond more than 8 electrons around it

- only exist in 3rd row elements and beyond

- only exists when an expanded octet results in a lower formal charges and more stable structure, otherwise additional bonds are used

When a molecule or compound has more than 1 observable Lewis structure

The length of a bond between atoms measured in picometers (pm)

- more bonds between two atoms → shorter bond length → stronger bond strength

The number of electron pairs between 2 atoms in a bond

- Higher bond order → stronger bond strength

single bond = BO of 1

double bond = BO of 2

triple bond = BO of 3

(# of electron pairs in the equivalent bonds of interest) ÷ (# of total resonance structures for molecule) = BO of overall molecule

The energy needed to break 1 mole of the given bond

- shorter bonds (length) and multiple bonds (strength) → more energy to break