c ovalent c ompounds. t wo t ypes of b onds ionic: electrons are transferred covalent: electrons...

COVALENT COMPOUNDS

TWO TYPES OF BONDS Ionic: Electrons are transferred

Covalent: Electrons are shared Non-polar covalent: equally shared

Polar Covalent: unevenly shared

BOND POLARITY

Which element is the most electronegative?

H F

Fluorine- Has 7 valence e- and wants 8

ability of an atom to attract electrons

REVIEW:WHAT IS ELECTRONEGATIVITY?

H F FH

electron richregion

electron poorregion

e- riche- poor

d+ d-

POLAR BOND :

covalent bond with greater electron density around one of the two atoms

1

2

3 4 5 6 7 8 9 10 11 12

13 14 15 16 17

18

NonpolarCovalentshare e-

Polar Covalentpartial transfer of e-

Ionictransfer e-

Increasing difference in electronegativity

Electronegativity Difference Bond Type

0 to 0.3 Nonpolar Covalent

0.4 to 1.6 Polar Covalent

1.7 Ionic

WHAT TYPE OF BOND IS IT?

Classify the following bonds as ionic, polar covalent,or covalent:

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

Cl – 3.0 N – 3.0 3.0 – 3.0 = 0 Nonpolar Covalent

Cs to Cl

H to S

Cl to N

DO YOU NOTICE A PATTERN FOR THE COMBO OF ELEMENTS THAT ARE IONIC VS COVALENT? Ionic bonds form between:

Covalent bonds form between:

Identify the following as ionic, covalent, or both:

CaCl2 BaSO4

CO2 AlPO4

SO3 H2O

PROPERTIES OF COVALENT COMPOUNDS

Usually soft and squishy

Not soluble in water

Does not conduct electricity

Soluble in organic solvents

Low melting points

Low boiling points

PROPERTIES OF IONIC COMPOUNDS Combination of ions (cation/anion)

Tightly packed solids in a crystal lattice

Hard and Brittle

Usually soluble in water

Conducts electricity when dissolved

High melting points

High boiling points

NAMING COVALENT COMPOUNDS

NAMING COMPOUNDS

Nonmetal – Nonmetal

USE PREFIXES!

1. Change the ending of the second word to -ide

2. No mono on the first word

3. Drop any double vowels

COVALENT PREFIXESNumber of Atoms Prefix

1 Mono-2 Di-3 Tri-4 Tetra-5 Penta-6 Hexa-7 Hepta-8 Octa-9 Nona-

10 Deca-

EXAMPLES

1. CO

2. CO2

3. SO2

4. SO3

5. N2H4

6. N2O3

1. Carbon Monoxide

2. Carbon Dioxide

3. Sulfur Dioxide

4. Sulfur Trioxide

5. Dinitrogen Tetrahydride

6. Dinitrogen Trioxide

EXAMPLES

1. disilicon hexafluoride

2. tricarbon octachloride

3. phosphorus pentabromide

4. nitrogen monoxide

5. selenium difluoride

6. dihydrogen monoxide

1. Si2F6

2. C3Cl8

3. PBr5

4. NO

5. SeF2

6. H2O

EMPIRICAL AND MOLECULAR FORMULAS

Define Empirical Formula:A chemical formula that gives the simplest whole-number ratio of the elements in the formula.

Which of the following is an empirical formula? CO2 C2O4

N2H4 NH2

Define Molecular Formula:A chemical formula that gives the actual number of the elements in the molecular compound. For the following molecular formulas, write the empirical formula:

Molecular: Empirical: C2H4 C6H12O6

C9H21O6N3

LEWIS STRUCTURES

OCTET RULE

Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE

This is called the octet rule

Bond formation follows the octet rule…Chemical compounds tend to form so that each atom:by gaining, losing, or sharing electrons, has an octet of electrons in its valence energy level.

LEWIS DOT DIAGRAMS

• an electron-configuration notation with only the valence electrons of an element are shown, indicated by dots placed around the element’s symbol.

• tracks the number of valence electrons

• the inner core electrons are not shown

LEWIS DOT PRACTICE

Li N F

Be O Ne

F F

LEWIS STRUCTURES FOR COMPOUNDS The pair of dots between two symbols

represents a shared pair. How many shared pairs does each fluorine have

below?

An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

F F

LEWIS STRUCTURES

The shared pair of electrons is often replaced by a long dash.

Each dash represents TWO electrons

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

WHY SHOULD TWO ATOMS SHARE ELECTRONS?

To get a valence of 8 electrons!

HC

HC

H

H

MULTIPLE COVALENT BONDS double bond:

covalent bond in which two pairs of electrons are shared between two atoms

shown by two side-by-side pairs of dots or by two parallel dashes

MULTIPLE COVALENT BONDS triple bond:

covalent bond in which three pairs of electrons are shared between two atoms

shown by three side-by-side pairs of dots or by three parallel dashes

Bond Type

Bond Length(pm)

C-C 154

CC 133

CC 120

C-N 143

CN 138

CN 116Bond Lengths

Triple bond < Double Bond < Single Bond

LENGTHS OF COVALENT BONDS

BOND LENGTH AND BOND ENERGY

As atomic size increases, bond length increases, and as a result bond energy decreases

As you increase the number of bonds between two atoms, energy increases, while bond length decreases.

BOND LENGTH AND BOND ENERGY EXAMPLES

1. Which bond is greater in length: Br2 or F2?

2. The HF bond is 570 pm, the H2 bond is 436 pm, which bond requires more energy to break?

3. Which bond would require more energy to break C-C single bond or C=C double bond? Which bond is longer?

WRITING LEWIS STRUCTURES

1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Complete an octet for all atoms except hydrogen

4. If structure contains too many electrons, form double and triple bonds on central atom as needed.

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

WRITE THE LEWIS STRUCTURE OF NITROGEN TRIFLUORIDE (NF3).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

WRITE THE LEWIS STRUCTURE OF THE CARBONATE ION (CO3

2-).

When there are two or more Lewis structures for a single molecule

O C O

O

- -O C O

O

-

-

OCO

O

-

-

What are the resonance structures of the carbonate (CO3

2-) ion?

RESONANCE STRUCTURE:

SOME ELEMENTS DO NOT FOLLOW THE OCTET RULE

H HBe

F B F

F

There can also be expanded octets!

MOLECULAR GEOMETRY

VSEPR THEORY

Lewis Dot Diagrams are 2D but we live in a 3D world.

How are molecules actually arranged??

Follows the Valance Shell Electron Pair Repulsion Theory or VSEPR

AB2 – LINEAR

Number of Surround Atoms Number of Lone Pairs Bond Angle

2 0 180˚

Cl ClBe

AB3 – TRIGONAL PLANAR

Number of Surround Atoms Number of Lone Pairs Bond Angle

3 0 120˚

AB2E1 – BENT

Number of Surround Atoms Number of Lone Pairs Bond Angle

2 1 <120˚

AB4 – TETRAHEDRAL

Number of Surround Atoms Number of Lone Pairs Bond Angle

4 0 109.5˚

AB3E1 – TRIGONAL PYRAMIDAL

Number of Surround Atoms Number of Lone Pairs Bond Angle

3 1 107˚

AB2E2 – BENT

Number of Surround Atoms Number of Lone Pairs Bond Angle

2 2 104.5˚

PREDICTING MOLECULAR GEOMETRY

1. Draw Lewis structure for molecule.

2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.

3. Use VSEPR to predict the geometry of the molecule.

What are the molecular geometries of SO2 and SF4?

SO O

AB2E

bent

C

F

F

F F

AB4

tetrahedral

INTERMOLECULAR FORCES

Intermolecular forces: attractive forces between molecules.Intramolecular forces: attractive forces within a molecule (the bonds)

intramolecular forces are much stronger than intermolecular forces

Intramolecular Forces

Intramolecular Forces

Intermolecular Forces

DIPOLES

What is a dipole?

A polar molecule

Uneven sharing of electrons so there is a separation of charge

DIPOLE-DIPOLE FORCES

Attraction between two polar molecules

— + — +

HYDROGEN BONDING

Special type of Dipole – Dipole

Attraction between:Hydrogen and Nitrogen/Oxygen/Fluorine

DIPOLE – INDUCED DIPOLE Attraction between one polar and one

nonpolar molecule

— +

— + — +

Electrons shift toward

positive end of dipole

LONDON DISPERSION FORCES Attraction between two nonpolar molecules

— + — +

Electrons become

uneven and form a dipole

STRENGTH OF IMF

Hydrogen Bond

Dipole – Dipole

Dipole – Induced Dipole

London Dispersion Forces

strongest

weakest

Which of the following molecules is polar?H2O, CO2, SO2, and CH4

O HH

dipole momentpolar molecule

SO

O

CO O

no dipole momentnonpolar molecule

dipole momentpolar molecule

C

H

H

HH

no dipole momentnonpolar molecule

SO

O

What type(s) of intermolecular forces exist between each of the following molecules?

HBrHBr is a polar molecule: dipole-dipole forces. There are

also dispersion forces between HBr molecules.

CH4

CH4 is nonpolar: dispersion forces.

SO2

SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

WHAT DOES IMF EFFECT? Viscosity

Surface Tension

Cohesion/Adhesion

Boiling Point

Stronger IMF Higher Viscosity

VISCOSITY

Measures a fluid’s resistance to flow

Stronger IMF Higher Surface Tension

SURFACE TENSION

result of an imbalance of forces at the surface of a liquid.

BOILING POINT

Point at which liquid particles escape the surface of the liquid into the gas phase

Stronger IMF Higher Boiling Point

Adhesion

Cohesion

ADHESION AND COHESION Cohesion:

intermolecular attraction between like molecules Adhesion:

intermolecular attraction between unlike molecules