chemical bonding: bonding theory and lewis formulas

Chemical Bonding:Bonding Theory and

Lewis Formulas

Periodicity There are certain trends within

the periodic table which affect reactivity and the ability to form bonds.

Periodic Law When elements are arranged in

order of increasing atomic #, elements with similar properties appear at regular intervals.

0

50

100

150

200

250

0 5 10 15 20

Atom

ic R

adiu

s (p

m)

Atomic Number

FCl

Br

1 2 3 4 5 6 7

Chemical ReactivityFamilies

Similar valence e- within a group result in similar chemical properties

1 2 3 4 5 6 7

Chemical Reactivity Alkali Metals Alkaline Earth

Metals Transition

Metals Halogens Noble Gases

Atomic Radius

Properties that affect Reactivity

size of atom

0

50

100

150

200

250

0 5 10 15 20

Atom

ic R

adiu

s (p

m)

Atomic Number

Atomic Radius

Li

ArNe

K

NaHalogens

Noble Gases

1 2 3 4 5 6 7

Atomic RadiusAtomic radius increase as you move right to left.

Atomic radius increase as you move top to bottom.

Ionization energy is the amount of energy needed to turn a stable element into an ion.

0

500

1000

1500

2000

2500

0 5 10 15 20

1st I

oniz

atio

n En

ergy

(kJ)

Atomic Number

Ionization Energy

KNaLi

ArNeHe

Again when you compare the ionization energy of elements there is a trend..

1 2 3 4 5 6 7

Ionization energy increases as you move right and down.

Ionization Energy

Why are atomic radius and ionization energy opposite?

Ionization Energy

In small atoms, the negative electrons (e-) are closer to the positive nucleus where the attraction is stronger.

Small Atoms Large Atoms

50 mμ

20 mμ

Valence Electrons, Valence Energy Levels & Valence

Orbitals Periodic trends are related to the

number of valence electrons an element has.

Valance electrons are those electrons occupying the highest energy level of an atom (the outside shell).

How Many Valence Electrons?

***You can tell how many valence electrons are in each orbital by counting the number of elements on each row of the periodic table.***

First Level: 2e-

Second Level: 8e-

Third Level: 8e-

Ex. The magnesium atom has 12 protons and 12 electrons.

The maximum number of electrons in the first energy level is 2 e-

That leave 10 electrons left to place.The next energy level

cam only hold 8 electrons.

That only leaves 2 electrons that are TRUE valence electrons

Orbital: a region in space in which an electron with a given energy is likely to be found.

There are four valence orbitals within the valence energy level of an atom (1s and 3p’s)

There are few exceptions to this rule

Hydrogen Helium

Electrons will occupy all valence orbitals before forming electron pairs. Empty bus seat rule

Normally a maximum of 8 electrons may occupy a valence energy level. This is known as the octet rule.

urinaltest

2 electrons in first energy level.

\Leaving you with 4 more electrons to place.

You place the other 4 electrons in each of the four orbital.

Start at the top and go clockwise.

Electron Dot Diagrams (Lewis Symbols)

Electron dot diagrams can represent atoms (neutral) or ions (charged).

ONLY show the atom’s valence electrons! – These are the only electrons involved in a chemical reaction (the electrons in the outer most ring!)

Dot Diagram Trends

Follow these steps:1) Write the atomic symbol for the

atom. This symbol represents the nucleus and the core electrons that do not participate in the chemical bonding.

2) Dots () represent the electrons in the valence energy level of the atom. Arrange these dots around the atomic symbol.

3) One dot must be placed in each of the four orbitals before any electron pairing occurs.

4) Begin with the fifth electron to make lone pairs. (if you have to)

5) There is a maximum of 8 electrons that can be drawn.

Lets try some… Calcium

Oxygen

Bromine

Carbon

Ca

O

Br

C

Bonding Electrons versus Lone Pairs

Bonding electrons are unpaired electrons that are involved in bond formation.

Paired electrons are called lone pairs and are generally not involved in bond formation.

Bond Types There are 3 types of bonds that

can be formed

These are determined by which

elements combine

1. IONIC 2. COVALENTBond FormationType of Structure

Solubility in WaterElectrical ConductivityOtherProperties

e- are transferred from metal to nonmetal

high

yes (solution or liquid)

yes

e- are shared between two nonmetals

low

no

usually not

MeltingPoint

crystal lattice true molecules

Types of Bonds

Physical State solid liquid or gas

odorous

Metallic Bonding 2 metals share electrons but no

chemical reaction occurs Valence electrons are free to move

about between the atoms Positive ions surrounded by a “sea”

of mobile electrons Allows metals to be formed into any

shape

1. Ionic Bonding A complete transfer of electrons

occurs in an ionic bond.

Valence Electrons

1. Ionic Bonding

1 7 8 8

2. Covalent Bonds

Results in a mutual sharing of electrons between the two non-metals.

This Sharing can be:

Equal = nonpolar covalent Unequal = polar covalent

2. Covalent Bonding

e- are shared equally between both nucleus.

Electron “Cloud” is symmetrical.

Nonpolar Covalent Bond

+ -

Polar Covalent Bond e- are shared unequally asymmetrical e- density results in partial charges

(dipole)

+

Polar Covalent BondExample:

H2O(l)

OH

H

3.4 -2.2 =1.2

POLAR-

Bond Polarity Most bonds

are a blend of ionic and covalent characteristics.

Difference in electronegativity determines bond type.

ElectronegativityElectronegativity: The attraction an atom has for a

shared pair of electrons….aka….the strength an atom has to hold onto or take electrons.

Trends in Electronegativity

Name That Bond!!!!

Electronegativity Difference Between Two Atoms.

Type of Bond Between the Atoms

Descriptions of Electrons in the Bond

1.7 IonicTransfer of electrons between metal and nonmetal

0.4 < 1.7 Polar covalentElectrons shared unequally between unlike atoms

< 0.4 Nonpolar covalent

Electrons shared equally between identical nonmetals

Bond Polarity

Non-Polar

Polar

Ionic

(+) ( - )

Bond PolarityExamples: Cl2

HCl

NaCl

3.2-3.2=0.0Nonpolar

2.2-3.2=1.0Polar

0.9-3.2=2.3Ionic

Covalent bonds are classified as: 1. single (sigma bond),

2. double (1sigma and 1pi bond),

or 3. triple bonds (1sigma and 2 pi bonds) depending on the number of electrons shared between the two nuclei.

Electron Dot Diagrams for Ionic Compounds

Electrons are transferred from the metal to the nonmetal.

results in a net negative charge for the nonmetal.

And a net positive charge for the metal.

Electron Dot Diagrams for Ionic Compounds

Drawing

Na Cl [Na]+[ Cl ]-

Drawing

MgCl

Cl[Mg]2+

[ Cl ]-

Lewis Dot Diagrams

For Molecular Compounds

Octet Rule Remember…

Most atoms form bonds in order to have 8 valence electrons.

Exceptions:

Hydrogen 2 valence e-

Helium 2 Valence e-

Octet Rule

H O H

Drawing Lewis Diagrams1. Find total # of valence e-.2. Arrange atoms - singular atom is usually in the middle.3. Form bonds between atoms (2 e-).4. Distribute remaining e- to give each atom an octet (recall exceptions).5. If there aren’t enough e- to go around, form double or triple bonds.

Drawing Lewis Diagrams CF4

1 C × 4e- = 4e-

4 F × 7e- = 28e-

32e- FF C F

F

- 8e-

24e-

How many bonds can Carbon (C) form?

Drawing Lewis Diagrams BeCl21 Be × 2e- = 2e-

2 Cl × 7e- = 14e-

16e-

Cl Be Cl - 4e-

12e-

How many bonds can Beryllium (Be) form?

Drawing Lewis Diagrams CO21 C × 4e- = 4e-

2 O × 6e- = 12e-

16e-

O C O - 4e-

12e-

How many bonds can Carbon (C) form?

Polyatomic Ions To find total # of valence e-:

Add 1e- for each negative charge. Subtract 1e- for each positive charge.

Place brackets around the ion and label the charge.

Polyatomic Ions ClO4

-

1 Cl × 7e- = 7e-

4 O × 6e- = 24e-

31e- OO Cl O

O

+ 1e-

32e-

- 8e-

24e-

How many bonds can Carbon (C) form?

NH4+

1 N × 5e- = 5e-

4 H × 1e- = 4e-

9e- HH N H

H

- 1e-

8e-

- 6e-

2e-

Polyatomic IonsHow many bonds can Nitrogen (N) form?

Resonance Structures Molecules that can’t be correctly

represented by a single Lewis diagram.

Actual structure is an average of all the possibilities.

Show possible structures separated by a double-headed arrow.1. 2.

Resonance Structures

OO S O

OO S O

OO S O

SO3

1. 2.

3.

Don’t Panic!!!! Lewis dot diagrams can take some

time to figure out. If you cannot immediately determine

the right configuration, take a deep breath and try again.

Think of these as puzzles, keep on working with the pieces until they fit together

You Want What? So what is the difference between

a Lewis dot diagram? A chemical formula? And a structural diagram?

CH4

___________________ ___________________ ___________________Lewis Dot Diagram Chemical Formula Structural Diagram

I II III

3.3 Molecular Shapes and Dipoles

VSEPR TheoryValence Shell Electron Pair Repulsion

TheoryStereochemistry: The study of 3-D spatial

configurationElectron pairs (bond pairs and lone pairs)

orient themselves in order to minimize repulsive forces.

VSEPR TheoryIn non science speak:

The interaction of electrons is what forms bonds.

Electrons have a negative charge, so they want to be AS FAR AS POSSIBLE AWAY FROM EACH OTHER.

VSEPR TheoryTypes of e- Pairs

Bonding pairs - form bondsLone pairs - nonbonding e-

Lone pairs repel more strongly than

bonding pairs!!!

VSEPR TheoryLone pairs reduce the bond angle between

atoms.

Bond Angle

Determining Molecular Shape

Draw the Lewis Diagram.Count how many things attached to central atom

double/triple bonds = ONELone pairs = ONE

Shape is determined by the # of bonding pairs and lone pairs.

Know the base shapes & their bond angles as well as the

derived shapes

Base vs. Derived shapesBase Shapes: AKA electron pair geometry.

Include all lone pairs as a bond

Derived Shapes: AKA molecular geometry. Only worry about actual bonds.

1. Base Shape: Linear (AX2)2 total electron pairs2 bond pairs0 lone pairs180o bond angle BeH2

2. Base Shape: Trigonal Planar (AX3)3 total electron pairs3 bond pairs0 lone pairs120o bond angle

BH3

3. Derived Shape: Bent/Angular (AX2E1)3 total electron pairs2 bond pairs1 lone pair< 120o bond angle

SO2

4. Base Shape: Tetrahedral (AX4)

4 total electron pairs

4 bond pairs0 lone pairs109.5o Bond

angle

CH4

5. Derived Shape: Trigonal Pyramidal (AX3E1)4 total electron

pairs3 bond pairs1 lone pair< 109.5o bond

angle

NH3

6. Derived Shape: Bent/angular/v-shaped (AX2E2)4 total electron pairs2 bond pairs2 lone pairs<109.5o bond angle

H2O

7. Derived Shape: Linear (AX1E3)4 electron pairs1 bond pair3 lone pairs180o bond angle

HF

8. Derived Shape: Linear (AX2E3)5 total electron pairs2 bond pairs3 lone pairs180o bond anglesXeF2

ExamplesPF3

4 total3 bond1 lone

TRIGONAL PYRAMIDAL<109.5°

F P FF

AX3E

ExamplesCO2

O C O2 total2 bond0 lone LINEAR

180°AX2

Examples (You Try)NO2

-

C2H2

CH3OH

Molecular Polarity: Dipole TheoryPolar Molecule: The negative (electron charge) is not distributed symmetrically among the atoms

Dipole MomentDirection of the polar bond in a molecule.Arrow points toward the more e-neg atom.

H Cl+ -

Depends on:dipole momentsmolecular shape

Determining Molecular Polarity

Nonpolar MoleculesDipole moments are symmetrical and

cancel each other out.

F

F F

BBF3

Non-Polar

Polar MoleculesDipole moments are not symmetrical.

The bonds are between elements with an electronegativity difference of between 0.3-1.7.

***Presence of lone pairs is a good trick to decide if its polar*****

O

H H

H2OPolar

3.4

2.2 2.2

+

-

Group ActivityVSEPR shape and Polarity Flashcards.In groups of three.

Find yourself a place to work in the room with a flat surface.

Decide who is A, B, and C.

Group ActivityVSEPR shape and Polarity Flashcards.A’s Come up to the front and get

three markers of different colours.

B’s Come up to the front and get 5 coloured flashcards.

C’s You just sit tight…for now…..your time is coming.

Group ActivityVSEPR shape and Polarity Flashcards.When all the group members are back

and the group has all the materials you may do the following:

A’s: On the back (side without lines) of two cards you will draw the base shapes of Linear & Trigonal Planar.

Group ActivityVSEPR shape and Polarity Flashcards.B’s: On the back (side without lines) of

three cards you will draw the base shapes of Bent, derived shape of trigonal planar and Tetrahedral.

C’s: On the back (side without lines) of three separate cards you will draw the two derived shapes of Tetrahedral.

A,B & C: You’re the “Editors” to make sure their all correct pictures.

Group ActivityVSEPR shape and Polarity Flashcards.

LINEAR180°

TRIGONAL PLANAR120°

BENT<120°

TETRAHEDRAL<109.5°

TRIGONAL PYRAMIDAL<109.5°

BENT<109.5°

You have about 10 minutes.

Group ActivityVSEPR shape and Polarity Flashcards. Once all three group members have finished

drawing the base and derived molecules shapes on the flashcards….

Each group member will find a actual compound that is a match to the shapes they drew earlier

Each person will then draw the structural formula of the compound they found on the other side of the flashcard, along with the chemical formula of the compound.

You have 10 minutes to complete pictures.

Group ActivityVSEPR shape and Polarity Flashcards

EXAMPLE FRONT:

BACK:

TETRAHEDRAL<109.5°

CH4

VSEPR LABCH2Cl2

CH3ClCCl4

AlCl3 HClO

HCN NH3

PCl3

SiF4

BH3

C2H6

H2O2

C2H2Cl

NF3

C2H2

Cl2O

C2H4

C2Cl2

Group ActivityVSEPR shape and Polarity Flashcards. Once every member of your group has

completed their flashcards, share them in your group, everyone must be able to do ALL of the shapes by themelves!!

HINT: There are assignments in your assignment bundle that are very close to what you are doing…..

You have about 10 minutes to share in your group.

Meniscus?Capillary action?Is water attracted to a charged object?

Intramolecular bonds: attractions within a molecule.

Intermolecular forces: The weak forces or bonds between molecules.

*** Think of a highway in the states.

***What’s the highway that goes between different states called?

The Interstate.

•Intermolecular bonds involve the electrostatic attractive forces between molecules.

•Ionic substances do not form molecules.

• Therefore, intermolecular bonding only occurs in substances that form covalent bonds. (molecular Compounds)

Cl-

3.4

2.22.2

STRONG

WEAK

Types of Intermolecular Bonding Van der Waals forces can

be divided into three different types

London dispersion forces dipole-dipole forces hydrogen bonding STRONGEST

WEAKEST

London Dispersion ForcesWeak attractive forces that result when the

electrons of one molecule are attracted to the positive nuclei of a nearby molecule (random chance)

The movement of the electrons in a molecule generates temporary positive and negative regions in the molecule A temporary dipole

***Temporary***

Let us consider the Chlorine molecule, Cl2(g): •At a particular instant,

we may find that the two electrons that form the bond may be closer to one nucleus than the other.

•Results in a temporary dipole with one end more negative than the other.

- The positive end of one polar molecule will be attracted to the negative end of a neighbouring polar molecule. (Permanent)

Dipole-Dipole Forces

Occurs when hydrogen is bonded to a highly electronegative element (fluorine, oxygen and nitrogen) – chemistry is FON!!!

The hydrogen end of the bond takes on a strong positive charge because of the exposed positive nucleus, while the other element takes on a strong negative charge

This positive hydrogen will be attracted to nearby negative atoms.

It appears as though the hydrogen atom bonds to different molecules.

Hydrogen Bonding

Predicting Boiling PointsWhat affects boiling point?1. Number of Electrons

2. Number of Carbons

3. Polarity…electronegativity difference

Predicting Boiling PointsIn order to cause a substance or compound to boil,

You must provide enough energy to break the bonds in the compound.

More energy needed, higher temperature needed.

In order to describe how a solid looks and behaves, you first need to determine what class of substance it is.

Ionic Crystals

Metallic Crystals

Molecular Crystals

Covalent Network CrystalsVery hard (harder than ionic and molecular

crystals) Brittle with very high melting and boiling

pointsHigher than ionic and molecular

Insoluble nonconductors of electricity

Carbon based covalent network crystals

GROUP ACTIVITY :THINGS TO WATCH FOR

SATP: Standard Ambient (room) Temperature and Pressure.

“Compound” refers to the joining of elements (Metal-Nonmetal).

Ion charges of metals DO NOT have to be memorized, there on the periodic table.

Ion charges of non-metals must be memorized, as they are not on the periodic table.

“Molecule” refers to a molecular joining of elements (Nonmetal-Nonmetal).

INDIVIDUAL----->PAIR------>GROUP---->SHARE FIRST

Form pairs of two….quickly and quietly. Switch desks so that pairs are sitting next

to each other. Each pair decides who is going to be A

and who is going to be B.

INDIVIDUAL----->PAIR------>GROUP---->SHARE A: Reads Intro and all Ionic “Stuff” (p 119-

123).

B:Reads Intro and all Molecular/Covalent “Stuff” (p. 119(Intro) and 124-126).

As each partner reads over their pages, summarize the information (write it down).

Summarize in such a way that when you are done summarizing you can explain it to your partner.

First part is individual, so it should be pretty quiet.

You have about 25 minutes.

INDIVIDUAL----->PAIR------>GROUP---->SHARE Second Now each pair has the next 15 minutes to

explain their sections to each other…BOTH must have the summarized notes.

It should be a little louder now, but you should only be talking to your partners.

When you are done, bring up your summaries to me and so I can have a look.

INDIVIDUAL----->PAIR------>GROUP---->SHARE Third

The pairs should now join up with one other pair and make a “Place Mat” of their summarized ideas.

4 People in each group.

Divide into A, B, C, D, etc.

WAIT!

PLACE MAT ACTIVITY B’s get the markers. At the front.

C’s get the Paper. At the back.

A’s Divide paper into section with shared area in middle.

WAIT!

PLACE MAT ACTIVITY THIS TIME!

Each group member summarizes the information they have in their “spot” on the place mat.

You have 5 minutes

WAIT!

OK, NOW YOU CAN GO.

Four You then rotate the “Place Mat” and

check over and add (if necessary) information to the next persons summary.

You will keep rotating the “Place Mat” and checking/adding until you have rotated 4 times and are back to where you started (your own).

Use your time wisely. You have 15 minutes.

Summarize the information your group found in the middle.

Sort of a final copy..

Summarize the information your group found in the middle.

Sort of a final copy..