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BICOL UNIVERSITY College of Science Department of Chemistry CHEM 1 GENERAL CHEMISTRY LECTURE HANDOUT 4 Ver. 1.1 α 20110310

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BICOL UNIVERSITY College of Science

Department of Chemistry

CHEM 1

GENERAL CHEMISTRY LECTURE HANDOUT 4

Ver. 1.1 α 20110310

2 http://buchem.weebly.com michaelvmontealegre

The attractive forces between atoms in molecules result in the

formation of chemical bonds.

1. A covalent bond is formed when two atoms share a pair of

electrons.

2. An ionic bond results from the electrostatic attraction of a cation

for an anion.

3. A metallic bond consists of the nuclei of metal atoms surrounded

by a "sea" of shared electrons.

Covalent compounds tend to form so that each atom, by sharing

electrons, has an octet of electrons in its highest occupied energy

level.

The octet rule states that when atoms undergo chemical reactions,

they tend to do so in a way that all of the atoms in the resulting

compound have eight electrons in their outer energy shells, even if

they have to share some of them.

Covalent compounds involve atoms of nonmetals only.

The term “molecule” is used exclusively for covalent bonding

The chemical symbol for the atom is surrounded by a number of dots

corresponding to the number of valence electrons.

Number of Valence

Electrons 1 2 3 4 5 6 7 8

Example H Group 1

Alkali metals

Helium

Group II Alkali earth

metals

Group 13

Group 14

Group 15

Group 16 Group

17 Halogen

Group 18

Noble Gases

Lewis Structure

Valence electrons are electrons in the outmost shell (energy level). They are the electrons available for bonding.

BONDING 3

Lewis Structures for Covalent Compounds

Electrons are shared between atoms to form a covalent bond in order that each atom in the compound has a share in the number of electrons required to provide a stable, Noble Gas, electronic configuration.

Electrons in the Lewis Structure (electron dot diagram) are paired to show the bonding pair of electrons.

Sometimes the bond itself is shown with a “-“ DRAWING LEWIS STRUCTURE Step 1 Arrange atoms symmetrically around the central atom (usually listed first in the

formula, not usually oxygen and never hydrogen). Step 2 Count the number of valence electrons of all atoms. For polyatomic ions, add

electrons corresponding to the negative charge, and subtract electrons corresponding to the positive charge on the ion.

Valence Electrons = Group Number

Valence Electrons = Group Number-10

(except He = 2)

The d and f block Valence Electrons are much more complicated and not covered

in the course

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Step 3 Place a bonding pair of electrons between the central atom and each of the surrounding atoms.

Step 4 Complete the octets of the surrounding atoms using lone pairs of electrons. (Remember hydrogen is an exception.) Any remaining electrons go on the central atom.

Step 5 If the central atom does not have an octet, move lone pairs from the surrounding

atoms to form double or triple bonds until the central atom has a complete octet.

Step 6 Draw the Lewis structure and enclose polyatomic ions within square brackets showing the ion charge.

Examples (for the purpose of the course, we will only deal with simple molecules with one to two central atoms) Hydrogen fluoride, HF

OR

Ammonia, NH3

OR

Oxygen molecule, O2

OR

BONDING 5

The HONC Rule

Hydrogen (and Halogens) form one covalent bond Oxygen (and sulfur) form two covalent bonds

One double bond, or two single bonds Nitrogen (and phosphorus) form three covalent bonds

One triple bond, or three single bonds, or one double bond and a single bond

Carbon (and silicon) form four covalent bonds. Two double bonds, or four single bonds, or a triple and a single, or a

double and two singles Exercise CH3Cl

HCN

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Exceptions to the Octet Rule 1. Electron Deficient Molecules

2. Expanded Octet

Lewis Structures for Ions of Elements

The chemical symbol for the element is surrounded by the number of valence

electrons present in the ion. The whole structure is then placed within square brackets, with a superscript to indicate the charge on the ion.

Atoms will gain or lose electrons in order to achieve a stable, Noble Gas (Group 18), electronic configuration.

Negative ions (anions) are formed when an atom gains electrons.

Positive ions (cations) are formed when an atom loses electrons.

Charge on Ion

1+ 2+ 3+ 4+ 4- 3- 2- 1-

No. electrons

gained or lost

1e lost 2e lost 3e lost 4e lost 4e

gained 3e

gained 2e

gained 1e gained

Example H+ Group I +

(Alkali metals)

Group II 2+

(alkali earth

metals)

Group 13 3+

Group 14 4+

Group 14 4-

Group 15 3-

Group 16 2-

Group 17 - (Halogens)

H- (hydride)

Lewis Structure (electron

dot diagram)

or H+ or Li+ or Be2+ or B3+ or C4+

BONDING 7

An ionic solid is made up of positive ions (cations) and negative ions (anions) held together by electrostatic forces in a rigid array or lattice.

Ionic bonding refers to the electrostatic attraction between cations and anions.

The physical properties of ionic compounds are: o High melting and boiling points o Ionic solids do not conduct electricity (they are insulators). o When molten (liquid) ionic compounds conduct electricity. o When dissolved in water to form an aqueous solution ionic

compounds conduct electricity. o Hard o Brittle

Lewis Structures for Ionic Compounds

The overall charge on the compound must equal zero, that is, the number of electrons lost by one atom must equal the number of electrons gained by the other atom.

The Lewis Structure (electron dot diagram) of each ion is used to construct the Lewis Structure (electron dot diagram) for the ionic compound.

Examples Lithium fluoride, LiF

Lithium atom loses one electron to form the cation Li+

Fluorine atom gains one electron to form the anion F- or

Li+ OR

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Lithium oxide, Li2O

Each lithium atom loses one electron to form 2 cations Li+ (2 electrons in total are lost)

Oxygen atom gains two electrons to form the anion O2- or

2Li+ OR Li+ Li+ OR

A metal is a lattice of positive metal 'ions' in a 'sea' of delocalised electrons.

Metallic bonding refers to the interaction between the delocalised electrons and the metal nuclei.

The physical properties of metals are the result of the delocalisation of the electrons involved in metallic bonding.

The physical properties of solid metals are: o conduct heat o conduct electricity o generally high melting and boiling points o strong o malleable (can be hammered or pressed out of shape without

breaking) o ductile (able to be drawn into a wire) o metallic lustre o opaque (reflect light)

BONDING 9

Valence Shell Electron Pair Repulsion

The structure around a given atom is determined principally by minimizing electron pair repulsions.

Predicting a VSEPR Structure

1. Draw Lewis structure. 2. Put pairs as far apart as possible. 3. Determine positions of atoms from the way electron pairs are shared 4. Determine the name of molecular structure from positions of the atoms.

Effect of unshared pair of electrons (NH3)

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Steric

No.

Basic Geometry

0 lone pair 1 lone pair 2 lone pairs 3 lone pairs

1 linear

2 linear

linear

3

trigonal

planar

bent /

angular

linear

4

tetrahedral

trigonal

pyramid

bent /

angular

linear

5

trigonal

bipyramid

sawhorse /

seesaw

t-shape

linear

6

Octahedral

square

pyramid

square

planar

7

pentagonal

bipyramidal

pentagonal

pyramidal

BONDING 11

Draw the lewis structure and determine the molecular geometry

Lewis structure Molecular geometry (draw)

CH2Cl2

CO2

XeF2

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Book References: Chang, Raymond. Chemistry. 9th edn. Digital Content Manager. 2007.

Manning, Phillip. Essential Chemistry: Chemical Bonds. Infobase Publishing

New York NY, USA. 2009

Masterton, William L. and Hurley, Cecile N. Chemistry Principles and Reactions

6th edn. Brooks/Cole Cengage Learning. CA USA. 2009

Silberberg, Martin S. Chemistry: The molecular nature of matter and change. 5th

edn. The McGraw-Hill Companies, Inc. New York, NY USA. 2009

Internet Resources: http://antoine.frostburg.edu/chem/senese/101/measurement/

http://www.sciencegeek.net/Chemistry/index.shtml

http://www.visionlearning.com/library/cat_view.php?cid=1

http://www.ausetute.com.au/index.html