Lewis structuresvsepr theory

AP Chemistry Rapid Learning Series - 15

© Rapid Learning Inc. All rights reserved. :: http://www.RapidLearningCenter.com 1

Rapid Learning CenterChemistry :: Biology :: Physics :: Math

Rapid Learning Center Presents …p g

Teach Yourself AP Chemistry Visually in 24 Hours

1/86 http://www.RapidLearningCenter.com

Lewis Structures

AP Ch i t R id L i S i

Rapid Learning Centerwww.RapidLearningCenter.com/© Rapid Learning Inc. All rights reserved.

AP Chemistry Rapid Learning Series

Wayne Huang, PhDKelly Deters, PhDRussell Dahl, PhD

Elizabeth James, PhDDebbie Bilyen, M.A.2/86



Learning Objectives

Valence Bond Theory

Th O t t R l

By completing this tutorial you will learn…

The Octet Rule

Lewis Structures for:Elements

Covalent Compounds

Polyatomic Ions

Ionic Compounds

3/86

p

Valence Shell Electron Pair Repulsion Theory

Electron and Molecular Geometry

Concept MapChemistry

Studies

Previous content

New content

Matter

Compounds

One type is

ValanceBond

Theory

1 bonding theory is

4/86

MolecularGeometry

ElectronGeometry

StructuresLewis

Structures

Shown with

Used to determine



Valence Bond Theory and the Octet Rule

5/86

Definition: Valence Shell

Valence Shell – Outermost shell of electrons; the electrons with the highest principal energy level number; the electrons that form chemical bonds.

Cl: 1s2 2s2 2p6 3s2 3p5 7 valence electrons

6/86

4s 2 3d10 4p5Br: [Ar] 7 valence electrons



Definition: Valence Bond Theory

Valence Bond Theory– Bonds are formed by overlap of valenceby overlap of valence orbitals.

H HH H

Valence Bond

7/86

H HH Hs-orbital

Both atoms get to “count” the electrons that are being shared between the two.

Definition: Octet Rule

Octet Rule – Most atoms are more stable with a full valence shell (whichstable with a full valence shell (which is a noble gas configuration). A full shell has 8 electrons (“oct-” = 8).

More exceptions will be discussed soon, but for now, Hydrogen is an exception.

8/86

y g

Hydrogen’s valence shell only contains a 1s orbital, which can only hold 2 electrons.

Therefore, hydrogen is most stable with 2 electrons.



Determining # of Valence Electrons

9/86

Valence Electrons and the Periodic Table

The main groups of the periodic table have # of electrons = main group #.

1 2 3 4 5 6 7 8

10/86



The transition metals don’t have easy patterns. Here are some of the common elements:

Valence Electrons of Transition Metals

Element Valence Electrons

Al

Zn

Cd

3

2

2

11/86

Ag

Au

1

3

Lewis Structures

12/86



Definition: Lewis Structure

Lewis Structure – 2D visualization of how electrons are shared to formhow electrons are shared to form bonds between atoms.

Also called:- Electron Dot Structures- Dot Structures

13/86

- Lewis Dot Structures

Lewis Structures of Elements

Use the element’s symbol to represent the nucleus and core (non-valence) electrons.1

How to draw an element:

( )

Determine the number of valence electrons from the position on the Periodic Table.

Draw the electrons around the “nucleus”—one on each side before doubling up (Hund’s Rule—place one in each orbital before doubling).

2

3

14/86

Example: Draw the Lewis Structure for an oxygen atom.

OOxygen is in the 6th main group.

There are 6 valence electrons.



Lewis Structures—Lewis StructuresCovalent Compounds with 2 Elements

15/86

Binary Covalent Structures - 1

Arrange the atoms symmetrically.

Determine the # of valence electrons for each

1

For compounds with only 2 different non-metals:

Determine the # of valence electrons for each element.Draw the valence electrons—do not double up where 2 atoms are bonding.When atoms have 8 (2 for H), the structure is done.

2

3

4

Example: Draw the Lewis Structure for CH

16/86

Example: Draw the Lewis Structure for CH4

CCarbon is in the 4th main group.There are 4 valence electrons.

HHH

HHydrogen is in the 1st main group.There is 1 valence electron.

Carbon now has 8 electrons it’s sharing.

Each hydrogen has 2 electrons it’s sharing.



Binary Covalent Structure - 2



1

Another example:


2

3

4

Example: Draw the Lewis Structure for NH

17/86

Hydrogen is in the 1st main group.There is 1 valence electron.

Nitrogen is in the 5th main group.There are 5 valence electrons.

H

Nitrogen now has 8 electrons it’s sharing.Each hydrogen has 2 electrons it’s sharing.

Example: Draw the Lewis Structure for NH3

N HH

Lone Pairs and Bonding PairsLone Pair

Electrons not shared in a bond

Bonding PairElectrons shared between two atoms

N HHH

Lone Pairs are “counted” only by the one atom.

18/86

y yLone Pairs are important and must be drawn even though they aren’t bonding.

Bonding Pairs are “counted” by both atoms that are sharing them.



Lewis Structures - Multiple Bonds

19/86

Multiple Bond Example - 1a



1

Begin with the same steps:


2

3

4

Example: Draw the Lewis Structure for CO

20/86

Carbon is in the 4th main group.There are 4 valence electrons.

Oxygen is in the 6th main group.There are 6 valence electron.

O C OCurrently, carbon only has 6.

And each oxygen only has 7.

Example: Draw the Lewis Structure for CO2 .



When the previous steps do not result in full valences:

Move two unpaired electrons on adjacent atoms to bond together.

Multiple Bond Example - 1b

5 g

Example: Draw the Lewis Structure for CO

Repeat until all atoms have full valence shells.

21/86

O C ONow Carbon has 8.

And each oxygen also has 8.


Multiple Bond Example - 2a



1

Begin with the same steps:


2

3

4

Example: Draw the Lewis Structure for HCN

22/86

Carbon has 4 valence electronsHydrogen has 1 valence electronNitrogen has 5 valence electrons.

But the carbon and nitrogen each only have 6 electrons.

C NH

Example: Draw the Lewis Structure for HCN .




When the previous steps do not result in full valences:

Multiple Bond Example - 2b

5 g

Example: Draw the Lewis Structure for HCN


23/86

C NH

Example: Draw the Lewis Structure for HCN .

Now all valence shells are full.

Double and Triple Bonds

Double Bond

C OO

Triple Bond

C OO

C NH

24/86

A double bond is 2 pairs of electrons being shared.Double bonds are shorter and stronger than single bonds.

A triple bond is 3 pairs of electrons being shared.Triple bonds are shorter and stronger than double bonds.



Lewis Structures - CovalentLewis Structures Covalent Compounds with More Than 2 Elements

25/86

What Order do the Elements Go In?When there are more than two elements, how do you arrange them?

“COOH” is a carboxylic acid.1Hydrogen and halogens (F, Cl, Br, I) can’t go in the middle.Of the elements that can go in the middle, write in the order they’re given.Write the hydrogen and halogen atoms around what they’re next to in the formula.

2

3

4

26/86

Example: H5C2OH

Write in this order C C OHH

H H

HH



Finishing the Lewis StructuresOnce you’ve arranged the atoms, finish the process:

Arrange the atoms according to the formula.


1

Determine the # of valence electrons for each element.Draw the valence electrons - do not double up where 2 atoms are bonding.When atoms have 8 (2 for H), the structure is done.

2

3

4

27/86

Example: H5C2OH

C C OHH

H H

HH

Each carbon has 4 electronsEach hydrogen has 1 electronThe oxygen has 6 electrons.

Example 2 - 1

“COOH” is a carboxylic acid.

Hydrogen and halogens (F Cl Br I) can’t go in the1

Hydrogen and halogens (F, Cl, Br, I) can t go in the middle.Of the elements that can go in the middle, write in the order they’re given.Write the hydrogen and halogen atoms around what they’re next to in the formula.

2

3

4

28/86

Example: BrH2CCH2COOH

C C CBrH

H H

HOCarboxylic acid

OH



Example 2 - 2

Arrange the atoms according to the formula.


1


2

3

4

29/86


C C CBrH

H H

HO

O

HCurrently, some of the atoms are full…

But one carbon and one oxygen each only have 7.


Example 2 - 3

5 g


30/86

All valence shells are currently full.


C C CBrH

H H

HO

O

H



Moving Hydrogen Atoms AroundSometimes, in order to have all atoms with full valence shells, a hydrogen must be bonded in a different location.

Example:

Each carbon has 4 electrons.Each hydrogen has 1 electron.

C3H6

C C CH

H H

H

H

HTwo carbons do not have full valence shells.

31/86

They are not adjacent—they cannot double bond.If one hydrogen is moved to another carbon…

Now two carbons right next to each other have un-full shells and can double bond.

This move of the hydrogen is not prohibited by the given information (the formula C3H6).

Lewis Structures -Polyatomic Ions

32/86



Definition: Polyatomic Ion

Polyatomic Ion – Group of atoms covalently bonded that together havecovalently bonded that together have a charge.

e.g. NH4+, SO4

2-

33/86

Polyatomic CationA cation is a positively charge ion (loss of electrons).

Arrange the atoms according to the formula.Determine the # of valence electrons for each

1

2

In this case, there is no choice but to double up on a side with a bond.


2

3

4

34/86

Example: NH41+

N HHH

HNitrogen has 5 electronsEach hydrogen has 1 electron.The +1 charge means we can remove 1 electron!

H’s electron is removed to result in +1 charge

+1



Polyatomic AnionA anion is a negatively charge ion (gain of electrons).

Arrange the atoms according to the formula.Determine the # of valence electrons for each

1

2 Determine the # of valence electrons for each element.Draw the valence electrons—do not double up where 2 atoms are bonding.When atoms have 8 (2 for H), the structure is done.

2

3

4When placing electrons around the oxygen, do not place them where there is already a pair of electrons from sulfur.

35/86

Example: SO42-

S OOO

OSulfur has 6 electrons.Each oxygen has 6 electrons.The -2 charge means we can add 2 electrons!

-22 electrons are added to result in the -2 charge.

Lewis Structures -Another Approach

36/86



Another Approach - 1

Arrange the atoms symmetrically or according to the chemical formula.1

O C O


37/86

Determine the total # of valence electrons for each element.2


O C OCounting electrons:1 Carbon = 1 × 4 = 4 electrons2 Oxygens = 2 × 6 = 12 electrons


38/86

2 Oxygens 2 × 6 12 electronsTotal = 16 electrons



Put one pair of electrons in between each set of atoms.

3




39/86

2 Oxygens 2 × 6 12 electronsTotal = 16 electrons 161412

Place lone pairs around the most electronegative atom first (closest to F on the periodic table). Stop when you run out of electrons.

4



Example: Draw the Lewis Structure for CO2 .Oxygen is more electronegative than carbon

40/86

2 Oxygens 2 × 6 12 electronsTotal = 16 electrons 121086420



If any atoms do not have full valences, move a lone pair from an adjacent atom in to form multiple bonds.

5


O C O


Each oxygen has 8 electronsBut carbon only has 4 electrons

41/86

A Larger Compound - 1

Arrange the atoms symmetrically or according to the chemical formula.1


O

42/86

C CBrH

H H

HO

O

HC



Determine the total # of valence electrons for each element.2


Counting electrons:


O

43/86

1 Bromine = 1 × 7 = 7 electrons5 Hydrogens = 5 × 1 = 5 electrons3 Carbons = 3 × 4 = 12 electrons2 Oxygens = 2 × 6 = 12 electronsTotal = 36 electrons

C C CBrH

H H

HO

O

H

Put one pair of electrons in between each set of atoms.

3


Counting electrons:


O

44/86


3616

C C CBrH

H H

HO

O

H



Place lone pairs around the most electronegative atom first (closest to F on the periodic table). Stop when you run out of electrons.

4


Counting electrons:

Example:

Oxygen is most electronegative, followed by bromine.

BrH2CCH2COOH

O

45/861660


C C CBrH

H H

HO

O

H

If any atoms do not have full valences, move a lone pair from an adjacent atom in to form multiple bonds.

5


Example:

All have full valences except one carbon and one oxygen.

BrH2CCH2COOH

O

46/86Now all have full valences.

C C CBrH

H H

HO

O

H



Lewis Structures -Ionic Compounds

47/86

Definition: Ionic Compound

Ionic Compound – Metals transfers electrons to non-metals. Theelectrons to non metals. The resulting ions form an electrostatic attraction.

e.g. KCl, Na2SO4

48/86

e.g. KCl, Na2SO4



Ionic Compound ExampleAn ionic compound is between metals and non-metals.

Determine the # of valence electrons for each atom.1

Draw the valence electrons.

Transfer electrons from the metals to the non-metals to fill valence shells.

2

3

The metal will be left with no electrons in the electrons shell “drawn”.However, the next inner shell is full and it now the “valence” shell.

49/86

Example: KCl

ClKPotassium has 1 electron.Chlorine has 7 electrons.

+1 -1

Polyatomic Ionic Compound ExampleAn ionic compound is between metals and non-metals.

Determine the # of valence electrons for each atom.1

Draw the valence electrons.

Transfer electrons from the metals to the non-metals to fill valence shells.

2

3

The total charge should = 0

50/86

Example: Na2SO4 NaEach sodium has 1 electronSO4 is a polyatomic ion—it must be covalently bonded first.

+1

S OOO

O -2

Na+1

-1



Exceptions to the Octet Rule

51/86

Common Exceptions to the Octet Rule

# of Valence Electrons when “Full”

Element(s)

2

6

>8

H, He

B, Be

Any element in the 3rd period and below

52/86

Elements in period 3 and below have empty “d” orbitals that can be used to hold more than 8 valence electrons.



Exception Examples

“Full” with 6 electrons

B FF B FFF

S

53/86

Has 10 electrons

F

Electron andElectron and Molecular Geometry

54/86



VSEPR Theory

ValenceShellElectronP i

55/86

PairRepulsion Theory

What Does VSEPR Mean?

Bonds are made Negatively Bonds form asBonds are made of shared electrons (negatively charged subatomic particles)

Negatively charged things repel each other

Bonds form as far apart from each other (and other electrons) as possible

56/86

Valence shell electron pair repulsion



Definition: Electron & Molecular Geometry

Electron Geometry – Uses the VSEPR theory with the electron regions around the central atomregions around the central atom. An electron bond is a bond (single, double or triple…they all count as one region).

Molecular Geometry – Uses the VSEPR

57/86

Molecular Geometry – Uses the VSEPR theory with the atoms bonded around the central atom.

How to Count Electron Regions

21

Electron geometry depends on electron regions surrounding the central atom.

3

21

4

N HHH

C NH

58/86

1 2The triple bond is only one electron region.

C NH



Chemical Formulas for Geometry

Lone PairsCentral Atom

Each geometry has a “generic” chemical formula.

X EA

Lone PairsCentral Atom

Atoms bonded to central atom (“ligand”)

59/86

e.g. AX3E AXX

X

Linear GeometryAX2

2 electron regions2 electron regions

Bonds 180° apart

60/86

Examples: CO2, BeH2

Named: looks like a “line”



Trigonal Planar GeometryAX3

3 electron regions3 electron regions

Bonds 120° apart

61/86

Examples: BF3, C2H4

Named: It’s a flat (“planar”) triangle.

Tetrahedron GeometryAX4

4 electron regions

62/86

Bonds 109.5° apart

Examples: CH4, SO2Cl2

Named: If each plane (defined by 3 points) is covered, it’s a 4 (“tetra”) sided object (“hedron” ).



Trigonal Bipyramidal GeometryAX5

5 electron regions

Bonds: Inside “Triangle” = 120°

63/86

Bonds: Inside Triangle 120Between top/triangle/bottom = 90°

Examples: PCl5, AsF5

Named: If each plane (defined by 3 points) is covered, 2 pyramids with triangular bases are sitting base-to-base.

Octahedron GeometryAX6

6 electron regions

Bonds: 90°

64/86

Bonds: 90

Examples: SF6, PF6

Named: If each plane (defined by 3 points) is covered, it’s an 8 (“Octa”) sided object (“hedron”).



Determining Electron GeometryElectron geometry is determined by electron regions.

Electron regions

Name of geometry Angle between regions

Picture

Linear

Trigonal Planar

Tetrahedral

180°

120°

109.5°

2

3

4

65/86

Trigonal Bipyramidal

Octahedron

90° and 120°

90°

5

6

Determining Molecular GeometryMolecular geometry is determined by # of atoms bonded to the central atom.

Formula Name of geometry Angle between bonded atoms

Picture

Linear

Trigonal Planar

Tetrahedral

bonded atoms

180°

120°

109 5°

AX2

AX3

AX

66/86

Tetrahedral


Octahedron

109.5

90° and 120°

90°

AX4

AX5

AX6



Geometry with Lone PairsUse the formula to determine the electron geometry.

Each atom bonded to the central atom (ligand) counts as 1.Each lone pair counts as 1.ac o e pa cou ts as

e.g. AX2E2= 2 bonds & 2 lone pairs = 4 electron regions

To determine molecular geometry, first determine

67/86

the electron geometry and then remove atoms to form lone pairs and re-name the geometry.

Bent Geometry with 1 Lone PairAX2E

Start with 3 electron regionsTrigonal Planar

Remove 1 atom

Bent 120º

68/86

Examples: SO2, O3

Named: It looks like a “bent” line.

Trigonal Planar



Bent Geometry with 2 Lone PairsAX2E2

Start with 4 electron regionsTetrahedron

Remove 2 atoms

Bent 109.5º

69/86

Examples: H2O, SF2

Named: It looks like a “bent” line.

Tetrahedron

Trigonal Pyramidal GeometryAX3E

Start with 4 electron regionsTetrahedron

Remove 1 atom

Trigonal Pyramidal

70/86

Examples: NH3, SOCl2

Named: It looks like a pyramid with a triangular base.

Tetrahedron



See-Saw GeometryAX4E

Start with 5 electron regionsTrigonal Bipyramidal

Remove 1 atom

See-Saw

71/86

Examples: SF4, O2XeF2

Named: It looks like a see-saw


T-Shaped GeometryAX3E2

Start with 5 electron regionsTrigonal Bipyramidal

Remove 2 atoms

T-Shaped

72/86

Examples: ClF3, BrF3

Named: It looks like a “T”




Linear with 3 Lone PairsAX2E3

Start with 5 electron regionsTrigonal Bipyramidal.

Remove 3 atoms

Linear

73/86

Examples: XeF2, I3

Named: It looks like a “line”

Trigonal Bipyramidal.

Square Pyramidal GeometryAX5E

Start with 6 electron regionsOctahedron

Remove 1 atom

Square Pyramidal

74/86

Examples: BrF5, IF5

Named: It looks like a pyramid with a square base.

Octahedron



Square Planar GeometryAX4E2

Start with 6 electron regionsOctahedron

Remove 2 atoms

Square Planar

75/86

Examples: ICl4, XeCl4

Named: It looks a flat (“planar”) square.

Octahedron

Effect of Lone Pairs on Molecular Geometry

N HHH

Both molecules have 4 electron regions:

These electrons areN HHH

C HHH

These electrons are not being “controlled” by another nucleus.

Lone pairs take up more space than a bonding pair—they distort the bond angles slightly.

76/86

Instead of being 109.5° angles (tetrahedron), the angles are 107.3°

107.3°107.3°

N HHH



Determining Geometry with Lone PairsFormula Name of geometry

Bent

Picture

AX2E

Bent

Trigonal Pyramidal

See-Saw

T-Shaped

AX2E2

AX3E

AX4E

AX3E2

77/86

LinearAX2E3

Square PyramidalAX5E

Square PlanarAX4E2

Geometry ExampleExample: Give the electron and molecular geometry for H2O.

O HH21 BondLone pair

O HH

First, draw the Lewis Structure.

For electron geometry, determine # of electron regions

34

4 electron regions = tetrahedral electron geometry

Lone PairBond

78/86

g g y

For molecular geometry, count atoms bonded to and lone pairs around the central atom.

2 bonded atoms and 2 lone pairs = bent molecular geometry

Note - this problem is commonly answered incorrectly as it “looks” linear as it’s written above. But the lone pairs make it bent!



LewisLewis Structures & The AP Exam

79/86

Lewis Structures in the Exam

Draw Lewis Structures

Common Lewis Structure problems:

Identify VSEPR geometries

Draw isomers or resonance structures

Use Lewis Structures and bonding theory to explain behavior.

80/86



Multiple Choice QuestionsProblems often combine bonding theory with Lewis Structures.

E l Whi h VSEPR t h 3d h b idi ti ?Example: Which VSEPR geometry has sp3d hybridization?

A. Bent

B. Trigonal bipyramidal

C Tetrahedron

sp2 or sp3

sp3d

81/86

C. Tetrahedron

D. Octahedron

sp3

sp3d2

Answer: B

Free Response QuestionsThe Free Response questions often ask for you to explain certain behavior.

Example: Explain the following in terms of atomic or molecular structure:

A. The second ionization energy of Mg is much lower than the second ionization energy of Na

B. The atomic radius of N is greater than OC. The boiling point of H2O is higher than H2SD. The bond angle of NH3 is smaller than NH4

+

82/86

A Lewis Structure/Bonding theory question combined with other atomic structure, intermolecular forces sub-questions.



Answering Free Response QuestionsDrawing structures along with labeling and brief explanation is an excellent why to answer a question!

D The bond angle of NH is smaller than NH +D. The bond angle of NH3 is smaller than NH4+

N HH N HHH

Lone pairs take up more space than bonding pairs.

83/86

H H

This distorts the other bond angles in towards each other.

Valence Shell Electron Pair

Repulsion (VSEPR) th i d t

Valence Shell Electron Pair

Repulsion (VSEPR) th i d t

Elements bond to obtain a full

valence shell—

Elements bond to obtain a full

valence shell—

Electron geometry is determined by electron regions, while molecular

Electron geometry is determined by electron regions, while molecular

Learning Summary

Covalent compoundsCovalent compounds

theory is used to determine

geometry of molecules.

theory is used to determine

geometry of molecules.

for most elements, that means 8 (the octet rule).

for most elements, that means 8 (the octet rule).

while molecular geometry is

determined by atoms bonded to the central atom.

while molecular geometry is

determined by atoms bonded to the central atom.

Lewis Structures areLewis Structures are

84/86

Covalent compounds share electrons, while

ionic compounds transfer electrons from one atom to another.

Covalent compounds share electrons, while

ionic compounds transfer electrons from one atom to another.

Lewis Structures are used to show the

valence electrons and their arrangement in

compounds.

Lewis Structures are used to show the

valence electrons and their arrangement in

compounds.



Congratulations

You have successfully completed the core tutorial

Lewis Structures

85/86

Rapid Learning Center

Rapid Learning Center

Wh t’ N t

Chemistry :: Biology :: Physics :: Math

What’s Next …

Step 1: Concepts – Core Tutorial (Just Completed)

Step 2: Practice – Interactive Problem Drill

Step 3: Recap – Super Review Cheat Sheet

86/86

Go for it!

http://www.RapidLearningCenter.com

Date post:	13-Jan-2015
Category:	Technology
Upload:	timothy-welsh
View:	416 times
Download:	0 times

Lewis structuresvsepr theory

Technology