Chapter 8 Covalent Bonding. Covalent bonds Atoms share their electrons When atoms share they create a molecule.

Chapter 8

Covalent Bonding

Covalent bonds

• Atoms share their electrons

• When atoms share they create a molecule.

Formation

• The electrons move about both nuclei and complete both valence shells.

Diatomic Molecules

• When two of the same type of atom join together

• H, N, O, F, Br, I, and Cl are the elements that form diatomic bonds.

Single Covalent Bonds

• Atoms share one pair of electrons.

• In a Lewis Structure the bond is represented by a line

Sigma Bond

• This is another name for the single covalent bonds

• It is the overlap of the two orbitals. Usually s and p or s and s.

Multiple Covalent Bonds

• Double bonds are when two pairs of electrons are being shared.

• Triple is three pairs

Pi bonds• These are multiple bonds.• Any additional electrons in a

bond after a single bond are pi bonds.

• Remember electrons repel each other, that is why a pi bond is formed.

8.2

Naming covalent compounds

Naming molecular compounds

• Put the most metallic element first

• The second element gets “ide”

• We use prefixes to show how many of each atom we have.

Prefixes

• Mono=1• Di=2• Tri=3 • Tetra=4• Penta=5• Hexa=6

• Hepta=7• Octa=8• Nona=9• Deca=10

Writing the Formula

• The name of the elements tell you the order

• The prefix is a subscript (it tells you how many)

8.4

Molecular Shapes

VSEPR

• Valence Shell Electron Pair Repulsion model

• Electrons push each other away making different angles.

VSEPR (cont’d)

• Electron pairs repel each other and cause molecules to be in fixed positions relative to each other.

• Unshared electron pairs also determine the shape of a molecule.

• Electron pairs are located in a molecule as far apart as they can be.

Hybridization

• This is when electrons move from one orbital type to another.

• It is a combination of two orbital types

• Carbon does this often

Bond Angle

• The number of bonds and pair electrons present influence the angle.

• More bonds/pair the smaller the angle.

Molecular Shapes

• Linear

• Trigonal planar

• Tetrahedral

• Trigonal pyramidal

• Bent

Molecular Shape cont’

• Bond angles of shapes

• Linear - 180˚

• Trigonal planar -120˚

• Tetrahedral – 109.5˚

Molecular Shape cont’

Molecular Shape cont’

Molecular Shape cont’

Section 8.5Electronegativity and

polarity

Electronegativity

• The attraction of electrons

• This determines the polarity of a molecule (along with shape)

Bond Character

• This tells you if it is ionic, covalent, polar, or nonpolar

Electronegativity Difference Bond Character

> 1.7 mostly ionic

0.4 - 1.7 polar covalent

< 0.4 mostly covalent

0 nonpolar covalent

Polar covalent bonds

• The unequal sharing of electrons

• Electrons are hanging out with one atom more than another

• Delta (δ) is used to say the partial charge or pole

Polarity and molecular shape

• If the bonds are polar but the molecule is symmetrical all around then the molecule will not be polar (you need to look for lone pairs of electrons too)

Examples

Solubility of polar molecules

• Solubility is the ability for a substance to dissolve in another substance.

• Molecules will dissolve in like substances (polar to polar)

Intermolecular Forces

• These forces hold substances together.

• They keep the atoms from just floating away.

Intermolecular Forces cont’

• Dispersion force – this is between two nonpolar molecules, it is very weak

• Dipole-dipole – this is between two polar molecules, it is strong.

Intermolecular Forces cont’

• Hydrogen bonds – these are very strong dipole-dipole forces

• It is H-O, H-F, or H-N

Forces and Properties

• The forces are what determine the properties.

• Weak forces will boil easier, or break apart easier.

Covalent Network Solids

• These create a lattice type solid

• They are typically hard, brittle, and non conductors.