Chapter 8

Chapter 8 lecture notes for CHM - 151 (General Chemistry)

Chemical Bonds

Chemical Bonds

• Metallic

• Ionic

• Covalent.

Metallic Bonding

Free electrons hold metals atoms together.

Ionic Bonding

Electrostatic attraction between ions. Sometimes describe as the “transfer of electrons".

Describes: Na(s) + ½ Cl2 → NaCl(s); Ag + Cl^- → AgCl(s)

Covalent Bonding

Sharing of electrons. Attractions of electrons and nuclei (p+). Repulsions between electrons and repulsions between nuclei (p+).

Properties of Ionic Substance

• Brittle

• High melting points

• Crystalline

• Cleave along smooth line.

Lewis Structures (Meaning)

Lewis symbols contain a number of dots to represent the number of valence electrons.

• Can be drawn to represent covalent bonding (none for ionic bonding).

• All valence electrons are shown; a line is a single bond and represents the sharing of two electrons (also called a “bonding pair”)

Octet Rule

Atoms share, lose, or gain electrons in order to have 8 valence electrons to become like noble gases.

Lattice Energy

Energy releases to make ionic compounds. Related to the strength of an ionic bond.

Larger ion size equals greater separation which equals…

smaller lattice energy, smaller attractive force.

Bigger charge equals greater attraction which equals…

greater attractive force and lattice energy.

Electronegativity

The ability of an atom to pull an electron towards itself.

Electronegativity ________ as you go left to right of a periodic table. Electronegativity ________ as you go down a periodic table.

increase; decrease.

Polarity of Bonds

• Nonpolar covalent: 0 - 0.4

• Polar Covalent: 0.5 - 1.9

• Ionic: 2.0+

Polar

Electrons are unequally shared.

Nonpolar

Electrons are equally shared.

Dipole Moment

Shows which element pulls under (more electronegative).

Duet Rule

Lewis Structure (Steps)

• Step 1: Draw skeleton structure “how are the atoms connected”. Usually the least electronegative atom are the center atom (not H).

• Step 2: Count valence electrons.

• Step 3: Add electrons so the total number equals step 2. Add electrons to the outside atoms first.

• Step 4: Check total number of electrons and check octet/duet rule.

Double Bonds (=)

4 electrons shared, 2 pairs.

Triple Bonds (≡)

6 electrons shared, 3 pairs.

• Have the highest bond enthalpy, hardest to break, strongest bonds, and shortest bonds.

Formal Charges (Meaning)

Pretends all electrons are equally shared.

• All follow the octet rule and have the right amount of electrons. Formal charges will help pick which is best.

• Note: Formal charges are in circles.

Formal Charges (Formula)

Formal Charge = (valence e-) - ½(bonding e-) - (all nonbonding e-)

Best Structure have the ______ formal charge. The ____ electronegative atom should have a _______ formal charge.

fewest; most; negative.

Resonance Structures

Needed when one Lewis Structures doesn’t accurately represent a molecule. Instead, the average of the structures best describes the molecule.

Resonance

Electrons are moved, same number of electrons, same atom arrangements.

Exceptions of Octet Rule

• Odd number of electrons.

• Fewer than eight electrons.

• More than eight electrons.

Odd Number of Electrons

Second is better because the oxygen is more electronegative.

Fewer Than Eight Electrons

H and Be can have 2 valence electrons. B can have 6 valence electrons (use formal charges). F will not form more than one bond (no positive formal charges for F).

More Than Eight Electrons

Elements on row 3 and below can have more than eight electrons because they have access to d orbitals (use formal charges).

Bond Enthalpy

Energy needed to break a bond. Takes energy to break bond.