Molecules

Objectives

• Write the electron dot structure for an atom.• Explain how covalent bonds form molecules.

Electron Dot Structures

Dots can be used to represent v/e.

Write electron dot structures for K, P, S, and Br.

Covalent Bonds

How do non-metal atoms bond together?

Example: chlorine gas, Cl2 2e-

Both nuclei are attracted to the

molecular orbital.

Both nuclei repel each other.

This equilibrium establishesthe bond length.

+ +

covalent bond: an attractionbetween non-metal atomssharing v/e in a molecular orbital

Molecules

molecule: group of neutral atoms held together with covalent bonds

molecular formula: indicates exact number and kinds of atoms in the molecule

IMPORTANTionic compoundformula: ratio of the ions in the crystalCaCl2 (1:2 ratio)sucrose: C6H12O6

Objective

• Be able to draw a Lewis diagram when given a molecular formula.

Lewis Structures

Draw Lewis structures for molecules containing the following elements: • nitrogen and fluorine• sulfur and chlorine• hydrogen and oxygen (water)

Lewis StructuresDouble and triple covalent bonds can also form:• H2CO (double)• ClCN (triple)

Stick Diagrams•C6H14

Objectives

• Understand the concept of VSEPR theory.• Use VSEPR theory to determine the shape(s)

present in a molecule.• Be able to draw Lewis diagrams for various

shapes.

VSEPR TheoryValence Shell Electron Pair Repulsion Theory: each pair of electrons (bonding pair or unshared pair) willrepel; molecule will adjust shape to maximize the anglesbetween each pair

Molecular Shapes

CH4

methanetetrahedral(4 single)

NH3

ammoniapyramidal(3 single)

H2Sdihydrogen

sulfide

bent(2 single)

Molecular Shapes

trignonal planar(2 single, 1 double)H2CO

formaldehdye

CO2

carbon dioxide

linear(2 double)

HCNhydrogencyanide

linear(1 single, 1 triple)

Objectives

• Be able to determine the polarity of a covalent bond using a table of electronegativities.

• Be able to determine the polarity of a molecule based on the shape(s) present and the polarity of the covalent bonds within the molecule.

• Understand and apply the concept of molecular symmetry.

Bond Polarity

Electrons are not always shared evenly.

H F

Atom with higher electronegativityattracts the e- pair more!

-+ polar bond: unevenlyshared covalent bond

Br Brnon-polar bond: evenlyshared covalent bond

2.2 4.0

3.0 3.0

Bond Polarity

What is the polarity of a H – C bond?H – O?K – Cl?

0.5 – 1.9 = polar covalent (uneven sharing)

Electronegativity difference determines bond polarity:

0.0 – 0.4 = non-polar covalent (even sharing)

2.0 – above = ionic bond (e- transfer)

Molecular Polarity

Bond polarity and molecular shape must be considered when determining molecular polarity.

dipole: a molecule that has an uneven distributionof charge (+ and - sides can be separated with a line)

+

-

-

+NH3 H2O

Molecular Polarity

A molecule with all non-polar bondsis called non-polar moleculeH-C bonds are always non-polar (example: hydrocarbons)

Due to symmetry, a moleculewith polar bonds can be anon-polar molecule.

Objectives

• Be able to determine the type of intermolecular bonding present in a molecular compound.

• Be able to predict the state of a molecular compound based on the type of intermolecular bonding present and the mass of the molecules.

Intermolecular Bonds

• intermolecular bond: a bond between molecules

• The state of a substance is determined by the strength of these bonds.

• Non-polar molecules tend to be gases (don’t attract): O2, N2, CO2

London Dispersion Forces

London dispersion force: caused by random distributions of electrons; brief partial charges cause attraction; more electrons = stronger bonds

non-polar moleculegas: < 70 e-liquid: 70 e- to 100 e-solid: > 100 e-

Dipole Interactions

• dipole interaction: dipoles attract each other; liquid or solid

• London force determines state:liquid: < 100 e-solid: > 100 e-

Hydrogen Bonds

• hydrogen bond: strong dipole interaction; occurs between molecules with –OH or –NH • always liquid or solid (depends on strength of London forces)

Hydrogen Bonds

H-bonds occur in both H2O and DNA.

Objectives

• Be able to name molecular compounds.• Know the diatomic elements.• Recognize elements that produce

macromolecules.

Molecular NamesCovalent Compound orMolecular Compound NomenclaturePrefixes1 mono- (mon-)2 di-3 tri-4 tetra- (tetr-)5 penta- (pent-)6 hexa- (hex-)7 hepta- (hept-)8 octa- (oct-)9 nona- (non-)10 deca- (dec-)

• use prefixes that match the number of atoms in the formula end the second name with “-ide”

• never start the first name with “mono-“; only use it for the second name

• The o or a at the end of a prefix is usually dropped when it precedes oxide.

Diatomic Elements• Many elements exist as paired atom molecules• H2, N2, O2, and halogens

Macromoleculesmacromolecule: a huge moleculeC and Si produce macromolecules because eachelement can form four bonds per atom

diamond (C)

quartz (SiO2)

Properties of Covalent Compounds

• Composed on non-metals only• Covalent bonding—exist as molecules.• Solid, liquid, or gas• Usually non-conductors of heat and electricity

Polyatomic Ions and Coordinate Covalent Bonds (extra-credit)

Polyatomic ions are essentially charged molecules(usually with additional electrons).

Draw Lewis structures for NO2– and OH–

coordinate covalent bond: a bond that forms whenone atom donates both electrons to a bond

Examples: CO, SO3, CO32–