Chem 121A Chapter 8: Lewis Structures Dr Fetto

Ionic bonds

Formed thru electron transfer

Covalent bonds

Formed through electron sharing

Metallic bonds

Electrons are free to move throughout the metal

Octet rule

Noble gases have 8 valence electrons and satisfy the octet rule.

Atoms like to have 8 valence electrons

Ne: (He)2s²2p^6. This has full s and p orbitals!

Elements will gain lose, or share electrons in order to satisfy the octet rule.

Transition metals do not typically follow the octet rule, which is why their charges vary.

Lewis symbols

A way to visualize valence electrons on an element.

How to?

• Electrons are represented as dots

• Distribute dots evenly around chemical symbols

• Max 2 electrons per side

Ionic bonding

Metal + nonmetal. They bond through electron transfer, and this can be visualized with Lewis symbols. Explains swap and drop ;3

Energetic of ionic bonding

Alkali metals and halogens react violently to form salts. Salts are formed by creating a lattice. We measure the stability of an ionic compound by looking at lattice energy

Lattice energy

The energy required to separate one mol of a solid ionic compound into its gaseous ions.

NaCl (s) → Na+(g) + Cl-(g) Δ H= 788 kJ/mol

High energy needed to break it apart means a stable solid is formed. High energy equals being more stable. This is why salts have such high melting points.

Trends in lattice energy

Ion size

• as ion size increases down a column, lattice energy will decrease

• Ions can’t get as “snug” into the lattice.

Ion charge

• Ion charge will alter the potential energy between atoms

• Larger ion chargers (+ or -) will have higher lattice energy. Stick together more aggressively the higher the charge

Calculate by using charge first, size second.

Covalent bonds

Nonmetal + nonmetal. Bond through electron sharing, pooling valence electrons and evenly distributed until the octet is satisfied for all atoms.

Lewis structures: how to

1. Sum all valence electrons into a total pool

2. Determine the central atom

The least electronegative atom goes in the center. Fr lowest, F highest. Increases left to right, increases down to up.

3. Draw single bonds between central and outer atoms.

Single boned are 2 shared electrons. Do not draw bonds between outer atoms. Subtract used electrons from total.

4. Use the remaining electrons to satisfy octet of the outer atoms

5. Double check the octet of the central atom

If all atoms are satisfied and no electrons are left, the structure is complete

6. When needed:

• If too many electrons: place remaining electrons on the central atom as lone pairs

• If too few electrons: form multiple bonds to satisfy octet

Single bond (2 electros shared)

Double bond (4 electrons shared)

Triple bond (6 electrons shared) *

*rare

Exceptions to octet

Less than octet

• H: 2 electrons (1 bond.) can’t be a central atom.

• He/be: 4 electrons (2 bonds)

• B: 6 electrons (3 bonds)

More than octet

• Some elements can expand their octet, hypervalent

• Central atoms in row 3 or below can be hypervalent

Odd electron species

• Some compounds will have an odd number of electrons

• Unpaired electrons called free radicals

Structures with charge

When a covalent compound has a charge, the charge will add or remove electrons from the total pool.

Formal charge

Yip can double check your structure by determining the formal charge of each atom in the compound.

Formal charge= valence electrons - (1/2 bonding electrons) - (nonbonding electrons)

Dominant structures

Use formal charge to determine dominant/best overall structure.

The best structure will have:

1. As many formal charges of 0 as possible

2. The ∑ of all formal charges should equal the overall charge

3. Negative formal charges should be on the most electronegative atoms when possible

Resonance structures

Equivalent structures with multiple bonds at different locations are called resonance structures. You must draw all possible options for resonance structures

Lewis structure: hints

Only make double bonds if absolutely necessary. Avoid double bonds to halogens, it is rarely corrrect.

Double check formal charges. Formal charges do not trump octet! Don’t make more bonds than possible.

Follow the steps! Many structures lack anomalies. Don’t make ur life more difficult

Properties of covalent bonds

Bond strength

• determined by bond enthalpy: amount of energy required to break the bond3240

Higher bond enthalpy means stronger bonds multiple bonds are stronger than single bonds - have higher bond enthalpy

Bond lengths shorter bonds are stronger than longer bonds. When looking at multiple bonds, multiple bonds are shorter and therefore stronger

Bond enthalpy

As can determine Δ H is through looking at bond energies.

Bond enthalpy is the enthalpy change for the breaking of a particular bond in one molecule of a gaseous substance.

Enthalpies of reactions (ΔHrxn)

We can use bond enthalpies to determine Δ H using a slightly different reaction where:

Δ H rxn= ∑(bonds broken) - ∑(bonds formed)