A. The Basics scientists have studied the way

in order to

elements and compounds appear in nature categorize

most are combined with in nature (called ) and are

1.1 Forming and Representing Compounds

Chapter 1: Chemical Bonding

chemical bonding

metals non-metalsores solids

a few are found in theirmetals pure form…precious metals

metals (except ) in pure form are

Hgsolids

combine with one another to form

the only elements that are found in in nature are the

non-metals

solids, liquids or gases

nevercombined form noble gases

atoms in such a way that they create a

the can only hold and therefore satisfies the octet rule when it has

gain, lose or share electrons full outer energy

level… octet rule

first energy level

two e

two e in it

octet rule is a guideline …not all elements follow it at all times

called the

are the electrons in the

they are the only electrons involved in

for representative elements ( ) group number (ignore the “1” in front of groups above 10) tells you the number of

period number tells you the number of

valence e outermost

chemical bonding

groups 1,2 and 13-18 valence electrons

energy levels occupied by electrons

energy level of an atom

for many , the number of valence electrons

transition metals is not as predictable… it depends on the environment

eg)iron can be Fe3+ or Fe2+

the can be used to determine ion charge number of valence e

eg)Fe3+ had 3 valence electrons

around the ion

B. Electron Dot Diagrams you can’t see atoms and electrons,

therefore it is convenient to to show the structure and formation of

an is one such model

consists of the with

chemical bonds

electron dot diagram

symbol for the element dots

when drawing the diagrams, look up the , then place number of valence e dots around the symbol clockwise

draw models

representing thevalence e

for a maximum offour dots

if you have more electrons to place, go back to the start pairing up the e

Na

Al

Ca

Si

P

O

Cl

Ar

topof the symbol and

a orbital is called a and is (at this level)

a orbital contains a

the of an atom is the maximum number of that it can form (equals the number of )

full lone pairnot involved in bonding

half full bonding electron

bonding capacitysingle covalent bonds

bonding e

O

lone pairs

bonding e

Try These

H

He

Be

C

F

Draw the Lewis diagram (electron dot diagram) for each of the following:

S

Mg

K

Br

P

C. Ionic Bonding an is the

most have

they tend to these electrons and become

electrostatic attraction between oppositely charged ions

losepositive ions (cations)

ionic bond

metals three or fewer valence e

Na Na+

Crash course Lewis dot diagram

most have non-metals more than four valence e

they tend to and become

gain electrons negative ions (anions)

O2-

O

after ions form, the attraction between the

when drawing the electron dot diagrams for ionic compounds:

ions

the number of electrons by the must the number of electrons by the

the on the compound must be

you may have to have of the to balance out the

positive chargeandnegative chargedraws thetogether, forming anionic bond

lost metalequal gainednon-metal

net chargezero

more than onemetal and/or non-metal

charges

Examples

NaCl

MgO

Na Cl

OMg

Examples

CaF2

K2S

K

KS

F

F

Ca

O

O

O

Mg

Mg

Fe

Fe

N

MgN

Fe2O3

Mg3N2

notice the following about the diagrams:

the metal has (since they them)

the non-metal has the valence level

both ions have and the

charges = charges

no valence electronslose

filled

square bracketscharge

positive negative

D. Covalent Bonding a is formed when

compounds containing are called

two non-metals share a pair of electrons

covalent bond

covalent bonds molecular compounds

electron dot diagrams used to show molecular compounds are called Lewis

structures

ions are not formed!!!

Covalent Bond video

instead of transferring electrons, valence electrons are now to satisfy the

the electrons that are shared are called

sharing two or three pairs of electrons between two atoms results in a

shared

a bonding pair

double or triple bond,

octet rule

respectively

to draw the structures:

place the atom with the in the

arrange all other atoms around it as as possible

to make sure that all atoms have the (remember that hydrogen only needs electrons to be satisfied)

most bonding electronscentre

symmetrically

share electronsoctet rule satisfied

two

Covalent bonding video

P

eg) PH3

H

H

H

P

HH

H

Try These

1. HCl

2. CH4

3. F2

4. NBr3

Draw the Lewis diagram (electron dot diagram) for each of the following:

6. N2

5. C2H4

H Cl

F F

N Br BrBr

N N

C HHH

HC

H HC

H H

a is another way of drawing molecules

to draw them, figure out the Lewis structure then replace all

E. Structural Formulas

eg) PH3

Lewis Diagram Structural Diagram

P HH

H

P HH

H

structural formula (diagram)

shared pairs of e with a line and leave off the lone pairs

Try These

1. HCl

2. CH4

3. F2

4. NBr3

Draw the structural formula for each of the following:

6. N2

5. C2H4

H Cl

C HH

H

H

F F

N Br Br

Br

CH H

CH H

N N

most metals are at room temperature which means that there must be

metals form or with

in all the atoms

F. Metallic Bonding

solids

strong attractive forces

metallic bonding share all the valence e

the valence electrons are , which means they are from one atom to another

holding the atoms of a pure metal together

DO NOT covalent ionic bondsother metal atoms

delocalizedfree to move

metallic bonds are made up of a network of positive metal

ions a is the

this theory helps explain the of metals

eg) good conductors of electricity and heat, ductility, malleability

in a “sea” of electrons

metallic bond electrostatic force of attraction between the positive metal ions and the negative sea of electrons

properties

metal cations

“sea” of delocalized electrons

Metallic Bond Model

the of an element is the relative measure of the ability of an atom to

there is an attraction between the of an atom and the

A. Electronegativity electronegativity

valence e in an adjacent atom

1.2 The Nature of Chemical Bonds

attract electrons in a chemical bond

nucleus (protons)

nucleuselectrons

each element is designated a number to represent how strong it’s nucleus is at attracting another atom’s valence e

trend on periodic table – electronegativity anddecreases down group

electronegativity means

since do not readily react with other substances, electronegativities have been assigned to them

higher greater attraction(affinity)

increases across period

noble gasesnot

understanding electronegativity has contributed to the knowledge of bonding in ionic and molecular compounds

decreases

increases

Electronegativity and the Periodic Table

as you move from left to right across a period, both the and

B. Size & Electronegativity

electronegativity, however size of the atom

atomic number increasedecreases

here is why size decreases across a period:

the size of an atom depends on the of the containing the

in any given period, the valence e of each atom occupy the

as you move across the period, the and thus the in the nucleus

there is a between the and when there are more , therefore the atom is

radius energy level valence e

same energy level

atomic number increases number of protons increases

greater amount of attractionnucleus e

protonssmaller

3 p+

1 valence e

Li

6 p+

4 valence e

C

9 p+

7 valence e

F

Period 2 Elements

so, the next question is “why does electronegativity increase when atomic size across a period decreases?”

the strength of the attraction (and therefore electronegativity) between oppositely charged particles depends on two factors:

the between the charges – the attractive force between opposite charges with the between them

the of the charges – the attractive force is to the

distance

decreases square of the distance

magnitudedirectly proportional

amount of charge

this means that an atom that is and has lots of (like fluorine) will have a amount of electrostatic attraction (electronegativity) for the of another atom

big atoms have but they are by the therefore have a amount of attraction (electronegativity) for the of another atom

smallprotons

very largee

lots of protonsshielded inner levels of e

smalle

Cs SiF Si

nucleus of cesium

valence electrons of silicon

nucleus of fluorine

valence electrons of silicon

distance between nucleus of cesium and valence electrons of silicon

distance between nucleus of fluorine and valence electrons of silicon

electronegativities can be related to bond types:

ionic bonds occur between

C. Bond Type & Electronegativity

metals have electronegativities and will while non-metals have and will

the two ions that are formed will

metals and non-metals

lowlose e high

gain e

attract each other

electronegativities

and form achemical bond

covalent bonds occurs between

if you look at two atoms that have the electronegativity, like in H2(g), the two nuclei of the atoms will attract the

the electrons are

non-metallic atoms

same

electronswith exactly thesame strength

shared equally between the two atoms

when two non-metals that have electronegativities share electrons, the sharing is

the element with the electronegativity pulls the

different

no longer equal

higher

e closer to itself

this results in one end of the bond having a and the other end of the bond having a

+

slightly negative charge()slightly positive charge

(+)

bonds that have are called

unequal sharing of electronspolar covalent bonds

also called since the bonds have

bond dipoles

oppositely charged ends

H – F

“arrow” points towards element with higher electronegativity (-)

“+” at the end that is +

Bond Dipole Arrows

+ -

Try These:

Draw the bond dipole arrow, label the + and ends, and state the bond type (polar, nonpolar, ionic)

1. H – H

2. N – H

3. B – F

4. S – O

5. P – H

6. C – H

7. Cl – Cl

8. Si – Cl

9. O – H

10. Na – Cl

nonpolar polar

polar

polar

polar

nonpolar

nonpolar

polar

polar

ionic

+

+

+

+

+

+

-

-

-

-

-

-

0.8

2.0

0.8

0.4

1.3

1.2

you can use the difference in electronegativity between two atoms to determine

bond character 

mostly ionic polar covalent slightly polar covalent

non-polar covalent

3.3 1.7 0.5 0

Difference in Electronegativity

as you can see,

  

bonding is considered a and there is

continuum

bond classification is not simple

no clear distinctionbetween ionic and covalent bonding

ionic compounds have

they form so that are as as possible this is called a

crystal structure

3-D array of alternating positive and negative ions crystal lattice

A. Ionic Crystals

2.1 Three Dimensional StructuresChapter 2: Chemical Bonding

oppositely charged ionsclose together

since all the in the lattice are the

attractive forces same, you cannot call them

each positive ion is attracted to of the negative ions around it

the chemical formula is the lowest whole number ratio

molecules…all

(and vice versa)

for that type of crystal

sodium chloride

eg) NaCl has a 1:1 ratio of Na ions to Cl ions

shape also depends on the relative size of the ions and the charges on the ions

there are many different and they all depend on the way the ions

crystal shapespack together

B. Structure of Molecules in molecular compounds, covalent bonds exist between

these compounds exist as that have a and therefore are not necessarily written with the

specific pairs of atoms

molecules given number of atoms lowest whole number ratio

C. VSEPR & Structure of Molecular Compounds

electron pair repulsion is

the states that molecules adjust their shapes so that valence e-

shape is determined around the

valence shell electron pair repulsion (VSEPR) theory are as far away from each other as possible

not always equal… it is greatest between twoless between a and aand lowest between two

central atom

lone pairs (LP),LP bonding pair (BP),

BP’s

VSEPR

shapes can be classified into five categories:

1. linear – central atom is bonded to and has lone pairs, or there is only two atoms in the molecule eg) CO2(g), HCN(g), HCl(g)

C OO

C NH

ClH

two other atoms zero

2. trigonal planar – central atom is bonded to and has lone pairs eg) CH2O(l)

O

HH

C

3. tetrahedral – central atom is bonded to and has lone pairs eg) CH4(g)

H

H H

H

C

three other atoms zero

four other atomszero

4. pyramidal – central atom is bonded to and has lone pair

eg) NH3(g)

5. bent – central atom is bonded to and has either lone pairs eg) H2O(l), HNO(g)

H

H

HN

HH

O

three other atomsone

OH

N

two other atoms one or two

the code has two numbers:

we can use a to determine the shape of a molecule around the

code

1. the number of attached to the central atom 2. the number of on the central atom

central atom

atoms

lone pairs

eg) NH3(g)

H

H

HN

CH4

H

H H

H

C

3 - 1

pyramidal

4 - 0

tetrahedral

Code Shape Example

4 – 0

3 – 0

3 – 1

2 – 1

2 – 2

tetrahedral

trigonal planar

pyramidal

bent

bent

CH4

CH2O

NH3

HNO

H2O

***all other codes arelinear

VSEPR shapes

D. Polar Bonds & Polar Molecules a molecule that contains

can be overall

the individual are that can be to each other

if the bond dipoles are and , they each other out resulting in a

this canceling happens in symmetrical molecules

if the bond dipoles , the entire molecule will have a

polar covalent bondsnonpolar

bond dipoles vectorsadded

equal in strengthopposite in direction cancel

nonpolar molecule

do not cancel slightly positive and slightly negative end…called dipoles

general rules: tetrahedral: if all atoms

attached have the same pull (in or out), if different atoms attached

trigonal planar: if all atoms attached have the same pull (in or out), if different atoms attached

pyramidal: as long as there is a difference in electronegativity between the atoms

bent:

linear: …look at electronegativity difference

nonpolarpolar

nonpolarpolar

polar

polar

polar or nonpolar

Examples

1. H2O 2. HCl

4. C2HI3. C2H2

O

H H

H Cl

HH C C IH C Cnp np

polarpolar

nonpolarpolar

Try These

1. HF 2. CH4

4. PI3

H F

polar

polar

nonpolar

3. N2

5. C2H6

N Nnp

PI

II

CH

HH

H

nonpolar

nonpolar

HH C Cnp

H

H H

H

A. Types of Forces

responsible for state, melting point, boiling point etc.

are the forces of attraction intermolecular forces between molecules

are the forces of attraction intramolecular forces within molecules

(eg. ionic or covalent bonding)

2.2 Intermolecular Forces

they are theweakest of all forces

there are three types of intermolecular forces that we will look at:

1. Dipole-Dipole Forces electrostatic force of attraction between the

attract the in other molecules and vice versa

dipoles of polar molecules

+

+

+

+

+-

- -

-

-

slightly negative “poles” slightly positive “poles”

2. Hydrogen Bonding this is a special type of dipole-dipole

interaction that is

hydrogen bonding is the attraction between a

when hydrogen is bonded to a such as O, F or N, the electrons are pulled from it

very strong

hydrogen on one molecule which is bonded to O, F or N, to the O, F or N of an adjacent molecule

highly electronegative element

far away

since hydrogen doesn’t have any other its is basically exposed

this proton is then able to be attracted not only to the

electrons,

pole

OH

H

OH H

O

H

H

O

HH

proton

but also to the lone pairs

3. London (Dispersion) Forces the attractive force that occurs between

molecules is called

the result of the electrostatic attraction of

electrons in atoms and molecules are always in

for brief instances, the distribution of electrons becomes which produces very weak

induced dipoles

constant rapid motion

distorted

allLondon Dispersion force

dipoles

this induces a dipole in the when

this process throughout the substance, causing that

even though this force lasts and is , the overall effect in a substance is

temporary dipole like charges repel each other

“disperses”flickering dipoles

attract each other

only a momentvery weak

significant

adjacent molecule

LD forces are affected by two factors:

1.size of the atoms – means higher probability of creating

2.shape of the molecule –

the more between molecules, the the force of attraction

more e

temporary dipoles

contacthigher

Youtube videos

Intramolecular vs intermolecular

http://www.youtube.com/watch?v=S8QsLUO_tgQ&feature=related

Scale of Forces

very low

very high

Intermolecular Forces

(between)

Intramolecular Forces

(within) London Dispersion Dipole – Dipole Hydrogen Bonding

metallic ** wide range ionic

network covalent

eg) diamond, SiC, SiO2

covalent

LDDD

HB ionic covalent

networkcovalent

A. States of Matter – Read p. 72 – 74!!! the of a substance depends on the

between its particles

2.3 Relating Structures and Properties

solids have the forces of attraction, liquids have the and gases have intermolecular attractions between the particles

state strength of the attractive forces

greatestnext greatest

very few if any

B. Melting & Boiling Points both melting and boiling points are indicators

as to the

when melt or boil, must be broken

strength of attractions within or between molecules

metals metallic bonds

when melt or boil, must be broken

ionic compounds ionic bonds

when melt or boil, must be broken

these forces are much therefore it takes a lot less

as you the number of intermolecular forces, the melting and boiling points

molecular compoundsonly intermolecular forces

(covalent bonds DO NOT break)

weakerenergy to separate the molecules

increaseincrease

Order of bp’s

Using the scale of forces you can order compounds based on their relative bp’s

Ex. From Highest to Lowest Network covalent compound (ex. SiO2) Ionic compound Molecular compound with HB, DD, LD Molecular compound with DD, LD Molecular compound with LD (if 2 molecular

compounds have LD only then bigger molecule or molecule with more electrons has higher bp)

C. Mechanical Properties of Solids the mechanical properties of solids are

determined by the types of bonds in the substance delocalized e cause bonds to be non-directional, which allows a solid to be malleable and ductile

solids that do not have delocalized e have directional bonds, which causes them to be brittle and hard

eg) metals

eg) ionic compounds

D. Conductivity electric current is the directional flow of

electrons or ions

metals are good conductors of electricity because the delocalized valence e are free to move

solid ionic compounds have valence electrons that are held solidly in place therefore they cannot conduct electricity

when ionic compounds melt or dissolve in water, the ions are able to move past one another which allows them to carry an electric current

in most molecular compounds, valence electrons are not free to move through the molecule therefore they are not able to conduct electricity

when molecular compounds melt or dissolve in water, they do not form ions and therefore they do not carry an electric current