Covalent Bonding - NJCTLcontent.njctl.org/courses/science/chemistry/covalent... · 2015. 11. 16. · Covalent Bonding & Molecular Geometry Examine these two forms of the same compound,

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Covalent Bonding

Note: Students and classrooms with iPads should download the

free "Lewis Dots" App and can use that on all the slides where Lewis

Dot drawings are to be done.

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Table of Contents: Covalent Bonding

· Properties of Ionic and Covalent Materials

Click on the topic to go to that section

· Naming Binary Molecular Compounds

· VSEPR Theory

· Covalent versus Ionic Bonds

· Resonance Structures

· Molecular Geometry

· Lewis Structures

· Polarity

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Return toTable ofContents

Covalent versus Ionic Bonds

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Covalent Bonding & Molecular Geometry

Examine these two forms of the same compound, ibuprofen.

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This form of ibuprofen has virtually no

anti-inflammatory effect.

This form of ibuprofen is about 100x more effective at

alleviating pain than the other form.

Even though they consist of the exact same number and kinds of atoms, these two molecules have very different chemical properties.

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In this unit, we will explore what causes molecules to have various shapes. Later, we will then examine how molecular

geometry affects different chemical properties.


Take a look around you. The chemistry of everything you see, hear, feel, touch and taste is a result of not only what it's made of but also how it's put together.(Remember this for next year in biology!)

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Chemical Bonds

Ionic - The electrostatic attraction between ions

Covalent - The sharing of electrons between atoms

Metallic - Each metal atom bonds to other metals atoms within a "sea" of electrons (covered in a later unit)

Chemical bonds hold atoms together to create chemical compounds. There are three basic types of bonds:

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Chemical BondsHow ionic or covalent a bond is depends on the difference in electronegativity. The smaller the difference, the more likely electrons are "shared" and the bond is considered covalent, the greater the difference, the more likely electrons have been transferred and the atoms are ionized resulting in an ionic bond.

Li Be B C N O FElectronegativity 1.0 1.6 2.0 2.5 3.0 3.5 4.0

Bond Li-F Be-F B-F C-F N-F O-O F-F

Electronegativity 3 2.4 2.0 1.5 1 0.5 0

Increasing Covalent Character

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Chemical Bonds

We can make a few simplifications...

Ionic Bonding Ionic bonds occur when the difference in electronegativity between two atoms is more than 1.7. Na ---- F electronegativity = 3

Covalent Bonding If the difference of electronegativity is less than 1.7, neither atom takes electrons from the other; they share electrons. This type of bonding typically takes place between two non-metals or between two metals. H ---- Cl electronegativity = 1.1

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In the case of ionic bonding, a 3-D lattice of ions is the result . . . not individual molecules. The chemical formula for an ionic compound is just the ratio of each type of ion in the lattice, not a particular number of ions in a molecule.

In contrast, covalent bonding can result in individual molecules or 3-D lattices depending on the elements involved. The bonding and the shapes of these molecules help determine the physical and chemical properties of everything around us!

Ionic v. Covalent Bonding

click here for an animationabout ionic and covalent bonding

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http://njc.tl/bt

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Properties of Ionic and Covalent Materials

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Properties of Ionic Compounds

Boiling and Melting Points

Since the attractions between the ions span a short distance, these forces are quite strong resulting in high melting points and boiling points!

Na+ -- Cl- it takes a lot of energy to break an ionic lattice!

Compound Melting Point (C)

NaCl 801

MgO 2852

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Properties of Ionic Compounds

Conductivity

Since ionic compounds consist of ions, when these ions are free to move, the substance can conduct electricity. To move, they must be in the liquid or molten state.

NaCl (s) Molten NaCl(l)

Lattice is strong, no conductivity Lattice is broken, ions are free to move and conduct

+

- -

--

-+

++

+

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Properties of Metallic Substances

Melting and Boiling Points

Metallic compounds are held together by non-directional covalent bonds in which some electrons are shared but are loosely held and free to roam. The covalent bonds between the metal atoms are strong! This gives rise to high melting and boiling points!

Metallic Lattice

strong metallic covalent bonds

Metal Melting Point

Cu 1085 C

Fe 1585 C

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In order to obtain pure metals, the ancients had to melt the metal (metallic substance) out of the rock (an ionic compound).

Copper has a lower melting point so it could be obtained in furnaces at lower temperatures. Furnaces hot enough to extract

iron would come later. Move for answer

REAL WORLD APPLICATION

Why do you think the bronze age (copper mixed with tin) came before the

iron age?

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Since the electrons in metals are free to roam somewhat, metals are good conductors of electricity!

Silver is the most conductive metal and is roughly 5-10 times more conductive than steel (mostly iron).

Properties of Metallic Compounds

Conductivity

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Copper is often used in electrical cable rather than

silver even though it is roughly 10% less

conductive than silver.

Why?


Copper currently trades for roughly 3 dollars an ounce while silver trades for about 30 dollars a month. It's about the money!!!!

Move for answer

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5 Which of the following would NOT conduct electricity in the solid state?

A Al

B Al2O3

C NaClD Both A and BE Both B and C

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5 Which of the following would NOT conduct electricity in the solid state?

A Al

B Al2O3

C NaClD Both A and BE Both B and C

[This object is a pull tab]

Ans

wer

E

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Like ionic and metallic substances, covalent network solids are giant molecules arranged in 3-D crystalline shapes. Here, the atoms involved tend to semi-metals like Silicon or Germanium or elemental carbon. Since the bonds are covalent, they are quite strong! This gives rise to high melting and boiling points!

Properties of Covalent Network Substances

Melting Point and Boiling Point

Glass (75% SiO2) Diamond (pure C)

Melts at 1500 C Melts at 3500 C

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Since these substances have higher electronegativities, they keep good tabs on their electrons thereby preventing the electrons from moving. As a result they are largely non-conductive.

Diamond and graphite are both allotropes or different versions of carbon and vary somewhat in their conductivity.

Properties of Covalent Network Substances

Conductivity

Diamond (C) Graphite (C)

non-conductive a little conductive

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Diamond is notorious for being HARD! This is true for lots of covalent network crystals. Can you think of some applications where hardness is important?

Body Armor

B4C (boron carbide)Drill Bits

polycrystalline diamond


slide for answers

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6 Which of the following would be classified as a covalent network solid?

A NaClB HFC CO2

D Ge2O3

E Fe


6 Which of the following would be classified as a covalent network solid?

A NaClB HFC CO2

D Ge2O3

E Fe


Ans

wer

D

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Molecular Compounds

When atoms are bonded covalently, the atoms are held together by sharing electrons. This occurs between non-metals such as C,O,S,H,P,N, etc. Unlike in all of the other substances, the atoms form small individual molecules that then interact with each other and their environment. These are called molecular compounds.

P O H H O = C = O Cl Cl Cl

In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. Both atoms use the shared electrons to reach that goal.

Click here to view interactive website

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Since these substances contain lots of small molecules, the bonds holding these small molecules together are fundamentally different from the covalent bonds found inside the molecule.

weak inter-molecular forces between molecules

Properties of Molecular Substances

Melting and Boiling Points

They cover a much larger distance and are quite weak giving rise to LOW melting and boiling points!

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Molecular compounds contain electronegative non-metals and do not lose their electrons easily so they are non-conductive.

As a result they are excellent INSULATORS!

Properties of Molecular Substances

Conductivity

Rubber: (C5H9)250

http://www.teachersdomain.org/asset/lsps07_int_covalentbond/

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Summary of Substances

Ionic Metallic Cov. Network Molecular

metals and non-metals metals semi-metals and

pure carbon non-metals

Na2O Fe C(diamond) CH4

High MP High MP High MP Low MP

conduct as liquid conduct in all states non-conductive non-conductive

Brittle Malleable Brittle Brittle

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7 Which of the following would have the lowest melting point?

A N2

B C(graphite)C C(diamond)D WE LiF


7 Which of the following would have the lowest melting point?

A N2

B C(graphite)C C(diamond)D WE LiF


Ans

wer

A

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8 Which of the following will not conduct electricity in any state?

A CuB NaFC FeD CO2

E All of these will conduct


8 Which of the following will not conduct electricity in any state?

A CuB NaFC FeD CO2

E All of these will conduct


Ans

wer

D

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9 Which of the following consists of small individual molecules?

A C(diamond)B SiO2

C Cu2OD NaE SO3


9 Which of the following consists of small individual molecules?

A C(diamond)B SiO2

C Cu2OD NaE SO3


Ans

wer

E

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10 Which of the following substances has both ionic and covalent bonding within the crystal?

A CuB CuCO3

C LiClD BaE BaF2


10 Which of the following substances has both ionic and covalent bonding within the crystal?

A CuB CuCO3

C LiClD BaE BaF2


Ans

wer

B

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Naming Binary Molecular Compounds

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Use prefixes to indicate the number the atoms.

All end in "ide"

Examples

NO2 nitrogen dioxide

P2O5 diphosphorous pentoxide ( penta-oxide-->pentoxide)

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Look on your reference sheets for the prefixes.

The atom with the lower electronegativity is usually written first.

If there is only one of the first atom, the mono- is left off.

Examples

CO carbon monoxide

CO2 carbon dioxide

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11 Chlorine monoxide is

A ClO2

B ClO

C OCl

D O2Cl


11 Chlorine monoxide is

A ClO2

B ClO

C OCl

D O2Cl


Ans

wer

B

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12 Dinitrogen tetroxide is

A NO2

B N2O4

C NO3-

D N4O2


12 Dinitrogen tetroxide is

A NO2

B N2O4

C NO3-

D N4O2


Ans

wer

B

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13 H2O is

A Hydrogen monoxide

B Dihydrogen monoxide

C Hydrogen oxide

D Hydrogen dioxide


13 H2O is

A Hydrogen monoxide

B Dihydrogen monoxide

C Hydrogen oxide

D Hydrogen dioxide


Ans

wer

B

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14 SO3 is

A sulfate B sulfur oxide

C sulfur trioxide

D sulfite


14 SO3 is

A sulfate B sulfur oxide

C sulfur trioxide

D sulfite


Ans

wer

C

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15 MgO is

A monomagnesium monoxideB magnesium monoxideC monomagnesium oxideD magnesium oxide


15 MgO is

A monomagnesium monoxideB magnesium monoxideC monomagnesium oxideD magnesium oxide


Ans

wer

D

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16 P4O10 is

A Phosphorous pentoxide B Tetraphosphorous decoxide

C Phosphorous oxide

D Phosphate


16 P4O10 is

A Phosphorous pentoxide B Tetraphosphorous decoxide

C Phosphorous oxide

D Phosphate


Ans

wer

B

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Lewis Structures

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Lewis structures are diagrams that show valence electrons as dots. Lewis structures are also known as Lewis dot or electron dot diagrams.

Note that no electrons are paired until after the fourth one.

Lewis Structures

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17 How many valence electrons does nitrogen have?

A 2B 3

C 4

D 5

E 7


17 How many valence electrons does nitrogen have?

A 2B 3

C 4

D 5

E 7


Ans

wer

D

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18 The Lewis structure for nitrogen is N

True

False


18 The Lewis structure for nitrogen is N

True

False


Ans

wer

True

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Recall that atoms tend towards having the electron configuration of a noble gas.For most atoms, that means having 8 valence electrons. The Octet Rule also applies to molecular compounds.

In covalent bonding, an atom will share electrons in an effort to obtain eight electrons around it (except hydrogen which will attempt to obtain 2 valence electrons).

The Octet Rule

A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbonding pair.

Exceptions to the Octet Rule

H needs 2e-

Be needs 4e-

B needs 6e-

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How do electron dot structures represent shared electrons?

An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots.

H + H H H

Hydrogen atom

Hydrogen atom

Hydrogen molecule

Shared pair of electrons

H

H

1s

1sHydrogen molecule

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Structural Formulas

A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms. As in the example below, one shared pair of electrons is represented by one dash.

HH

Hydrogen molecule

Shared pair of electrons

H H

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19 How many electrons are shared by two atoms to create a single covalent bond?

A 2

B 1


19 How many electrons are shared by two atoms to create a single covalent bond?

A 2

B 1


Ans

wer

A

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The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example.

Single Covalent Bonds

F F F F F F+ # # # OR

Fluorineatom

Fluorine molecule

Fluorineatom

1s

2s

2p

1s

2s

2p

Fluorine molecule

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In a water molecule, each hydrogen and oxygen atom attains a noble-gas configuration by sharing electrons.

Lewis Structure of H2O

The water molecule has two unshared, or lone, pairs of electrons.

2 H + O --> O H or O HH

HHydrogen

atomsOxygen

atomWater

molecule

1s 2p2s

1s 1s

O

H H

Water molecule

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In the ammonia molecule, NH3, each atom attains a noble-gas configuration by sharing electrons.

This molecule has one unshared pair of electrons.

Lewis Structures of NH3

3 H + N --> N H or N H

H

H

HHydrogen atom

Nitrogen atom Ammonia

molecule

1s 2p2s

1s 1sH

N

H1s

Ammonia molecule

H

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Drawing Lewis Structures

2. The central atom is the least electronegative element (excluding hydrogen).

3. Connect the other atoms to it by single bonds.

P has an electronegativity of 2.1 and Cl has an electronegativity of 3.0

P will be the central atom.

The Cl atoms will surround the P atom.

The single bonds are shown as single lines.

Cl

P

ClCl

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4. Count each single bond as a pair (two) of electrons.

5. Add electons to the outer atoms to give each one 8 (a full shell), or just 2 electrons for hydrogen.

6. Do the same for the central atom.

7. Check: Does each atom have a full outer shell (8 except, 2 for hydrogen)?

Have you used up all the valence electrons? Have you used too many electrons?


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The N atom has 5 valence electrons

and

each of the three H atoms has 1 so the total number of valence electrons is,

NH3

5 + 3(1) = 8

1. Find the total number of valence electrons in the polyatomic ion or

molecule.

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2. The central atom is the least electronegative element (excluding hydrogen because it can only have one bond).


H can never be the central atom so N must be

The H atoms will surround the N atom.

The single bonds are shown as single lines.

HN HH

NH3

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HN HH

Each H already has two electrons, so that's done. But we have to add electrons to N to make 8.

HN HH

4. Count each single bond as a pair (two) electrons. Now add electons to the outer atoms to give each one a full shell (2 in the case of H).

5. Next, do the same for the central atom.

6. Check:Does each atom have a full outer shell ?

7. Have you used up all the valence electrons you started with? Have you used too many electrons?

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20 How many total valence electrons does H2O have?

A

B

C

D

8

10

12

14


20 How many total valence electrons does H2O have?

A

B

C

D

8

10

12

14


Ans

wer

A

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21 Which element in H 2O is the least electronegative?

A

B

H

O


21 Which element in H 2O is the least electronegative?

A

B

H

O


Ans

wer

A

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22 Which of the following is the correct Lewis Structure for H2O?

A

B

C

D

H O H

H H O

H H O

H H O

1. Find the total number of valence electrons:

2. Central atom is the least electronegative:


4. Count each single bond as a pair of electrons.

5. Add electrons to the outer atoms to give each one 8 (except H only gets 2).

6. Add electrons to the central atom to give it 8.

7. Check to make sure all valence electrons are used.


22 Which of the following is the correct Lewis Structure for H2O?

A

B

C

D

H O H

H H O

H H O

H H O

1. Find the total number of valence electrons:

2. Central atom is the least electronegative:


4. Count each single bond as a pair of electrons.

5. Add electrons to the outer atoms to give each one 8 (except H only gets 2).

6. Add electrons to the central atom to give it 8.

7. Check to make sure all valence electrons are used.


Ans

wer

D

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23

A

B

C

H H H H H H C C

C C H

H H

H

H H

Which of the following is the correct Lewis Structure for C2H6?

C C H H H H H H

C C H H H H H H D


23

A

B

C

H H H H H H C C

C C H

H H

H

H H

Which of the following is the correct Lewis Structure for C2H6?

C C H H H H H H

C C H H H H H H D [This object is a pull tab]

Ans

wer

B

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Lewis Structures for ions

If you are drawing the Lewis Structure for an ION...

A negative ion has extra electrons, add the charge of the ion to your valence electron count.

ClO2- has 1(7) + 2(6) + 1 = 20 electrons

A positive ion is missing electrons, subtract the charge of the ion to your valence electron count.

NH4+ has 1(5) + 4(1) -1 = 8 electrons

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24

A

B

12

18

How many valence electrons does CO32- have?

C

D

2426


24

A

B

12

18

How many valence electrons does CO32- have?

C

D

2426


Ans

wer

C

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25

A

B

8

9

How many valence electrons does H3O+ have?

C

D

1011


25

A

B

8

9

How many valence electrons does H3O+ have?

C

D

1011


Ans

wer

A

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Formal ChargeThe "Formal Charge" method tells us how the electrons are distributed within a molecule. For example, depending on how the electrons are shared, some atoms may have more electrons than others resulting in a semi-charged state for that atom.

O

P

O

O OFC for P: 5 - 4= +1 (count each bond as one)FC for each O: 6 -7= -1 (count each bond as one)

Note: The charges must add to the charge of the molecule. So for PO43-

1 P atom x +1 = +1 + 4 O atoms x -1 = -4 +1 + -4 = -3

Formal Charge = # of valence electrons - # of electrons atom possesses within the lewis structure.

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Formal ChargeThe best Lewis structure will have the formal charge = 0 on each atom. However, if the molecule carries a charge, the more electronegative atoms should carry a charge as they have the greater attraction for electrons!

Each bond is counted as one in a formal charge calculation as each atom forming part of the bond contributes just one electron to that bond.

[ O - H ]-1 FC on O = 6-7 = -1 FC on H = 1-1 = 0

O H

The oxygen is more electronegative so it makes sense that it carries the negative charge.

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Example: Below are two possible lewis structure for the phosphate ion, PO4

3-. Which Lewis structure is considered to more closely represent the actual molecule based on formal charge calculations?

O

P

O

O O

O

P

O

O O

Structure 2 is superior as all formal charges = 0 whereas in structure 1, the P carries a +1 charge and each oxygen

carries a -1 charge

Structure 1 Structure 2

slide for answer

Formal Charge

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26 Which of the following would be the formal charge on the N in the ammonium ion?

A +1

B 0

C -1

D -2

E -3


26 Which of the following would be the formal charge on the N in the ammonium ion?

A +1

B 0

C -1

D -2

E -3


Ans

wer

A

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27 In which of the following molecules would N carry a non-zero formal charge?

A HCNB NH3C NO3-D NO2-E NH4+


27 In which of the following molecules would N carry a non-zero formal charge?

A HCNB NH3C NO3-D NO2-E NH4+


Ans

wer

E

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Lewis Structures

Draw the Lewis dot structure for the sulfate ion, SO4 2-,

and find the formal charge on each atom.

FC on S = 6-4 = +2

FC on O = 6-7 = -1

--------------------------

1(+2) + 4(-1) = -2

slide for answer

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Lewis StructuresDraw the Lewis dot structure for the hydronium ion, H3O+ and find the formal charge on each atom.

FC on O = 6-5 = +1

FC on H = 1-1 = 0

* note how in this case the more electronegative atom (O) is carrying a + charge relative to H. This demonstrates the

theory is imperfect.

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C

NCl

F

OSB

P

I

H

CO OSi

SeXe

CO2Draw a Lewis Structure

We ran out of electrons, but carbon does not have an octet

yet!

Now What?

Slide for Answer

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Double and Triple Covalent Bonds

Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons.

A bond that involves two shared pairs of electrons is a double covalent bond.

A bond formed by sharing three pairs of electrons is a triple covalent bond.

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Carbon Dioxide, CO2

1. Determine the # of valence electrons.

1 (4) + 2 (6) = 16 e-

This leaves 12 electrons, 6 pairs

3. Place lone pairs on oxygen atoms to give each 8.

Double and Triple Covalent Bonds

O C O

O C O2. Form Single Bonds

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O C O

Carbon Dioxide, CO2

4. Check: We had 16 electrons to work with; how many have we used?

5. There are too many electrons in our drawing. We must form DOUBLE BONDS between C and O.

Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond.

O C O

O C O

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Covalent Bond Length

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Covalent Bond Energy

It requires more energy to break double and triple bonds compared to single bonds.

Triple bonds are the strongest of the three.

Bond Type Bond Energy

C C

C C

C C

348 kJ

614 kJ

839 kJ

Bond Type Bond Energy

N N 163 kJ

418 kJ

941 kJ

N N

N N

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Covalent Bond Energies

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Covalent Bonds Comparison

Type of Bond

Electrons shared

BondStrength

BondLength

2

4

6

weak

intermediate

strong

long

intermediate

short

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28 As the number of bonds between a pair of atoms increases, the distance between the atoms:

A increases

B decreases

C remains unchanged

D varies, depending on the atoms


28 As the number of bonds between a pair of atoms increases, the distance between the atoms:

A increases

B decreases

C remains unchanged



Ans

wer

B

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29 As the number of bonds between a pair of atoms increases, the strength of the bond between the atoms:

A increases

B decreases

C remains unchanged



29 As the number of bonds between a pair of atoms increases, the strength of the bond between the atoms:

A increases

B decreases

C remains unchanged



Ans

wer

A

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30 As the number of bonds between a pair of atoms increases, the energy of the bond between the atoms:

A increases

B decreases

C remains unchanged



30 As the number of bonds between a pair of atoms increases, the energy of the bond between the atoms:

A increases

B decreases

C remains unchanged



Ans

wer

A

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31 How many electrons are shared by two atoms to create a single bond?


31 How many electrons are shared by two atoms to create a single bond?


Ans

wer

2

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32 How many electrons are shared by two atoms to create a double bond?


32 How many electrons are shared by two atoms to create a double bond?


Ans

wer

4

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33 How many electrons are shared by two atoms to create a triple bond?


33 How many electrons are shared by two atoms to create a triple bond?


Ans

wer

6

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34 Using Lewis structure drawings, determine which molecule below would have the shortest bond length between atoms?

A O2

B F2

C Cl2D COE I2


34 Using Lewis structure drawings, determine which molecule below would have the shortest bond length between atoms?

A O2

B F2

C Cl2D COE I2


Ans

wer

D

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35 Which of the following molecules would have the longest C-O bond length? Use Lewis structures.

A COB CO2C H2COD CH3OHE The lengths are all the same


35 Which of the following molecules would have the longest C-O bond length? Use Lewis structures.

A COB CO2C H2COD CH3OHE The lengths are all the same


Ans

wer

D

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If you run out of electrons before the central atom has an octet……form multiple bonds until it does.

Writing Lewis Structures

Slide 93 / 186

Oxygen molecule

Bonding of O2

1s

2s

2p

1s

2s

2p

O + O --> O O or O O

O

O

Oxygenatom

Oxygenatom

Oxygenmolecule

Oxygenmolecule

Slide 94 / 186

C

NCl

F

OSB

P

I

H

CSi

SeXe

CODraw a Lewis Structure

Carbon has the lower electronegativity, so we will consider it the "central" atom...

O

Slide for Answer

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Coordinate Covalent Bonds

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Coordinate Covalent BondsIn carbon monoxide, oxygen has a stable

configuration but the carbon does not.

1s 2p2s

2s1s 2p

C + O # # # C O

Carbonatom Oxygen

atom

Carbonmonoxide

C

OCarbon monoxide molecule

Slide 97 / 186

A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons.

In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them.

In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms.

Carbon has 4 valence electrons, oxygen has 6.

Coordinate Covalent Bonds

Slide 98 / 186

C

NCl

F

OSB

P

I

H

Si

SeXe

F2 Draw a Lewis Structure

F F

Slide for AnswerF

Slide 99 / 186

A molecule is a neutral group of atoms joined together by covalent bonds. Air contains oxygen molecules.

A diatomic molecule is a molecule consisting of two atoms. Certain elements do not exist as single atoms; they always appear as pairs.

When atoms turn into ions, this NO LONGER HAPPENS!

HydrogenNitrogenOxygenFluorineChlorineBromineIodine

Remember:HONClBrIF

Diatomic Molecules

H H

N N

O O

H2

N2

O2

Slide 100 / 186

36 On the periodic table below, mark which elements exist as diatomic molecules. Note the pattern.

Slide 101 / 186


There are three types of ions or molecules that do not follow the octet rule:

#1 Ions or molecules with an odd number of electrons

#2 Ions or molecules with less than an octet

#3 Ions or molecules with more than eight valence electrons (an expanded octet)

Slide 102 / 186

Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

NO is an example:

Exception 1: Odd Number of Electrons

Slide 103 / 186

Exception 2: Fewer Than Eight Electrons

Beryllium (Be) - this metal is shown to form molecular compounds, rather than ionic compounds as expected; only needs 4 electrons to be stable

Boron (B) - only needs 6 electrons to be stable

Memorize these exceptions

B

Be

Slide 104 / 186

The only way PCl5 exists is if phosphorus has 10 electrons around it.

This is called an expanded octet.

Atoms on the third energy level or higher are allowed to expand their octet to 10 or 12 electrons.

These atoms are larger and can accommodate more electrons.

Exception 3: Expanded Octet

Slide 105 / 186

How many electrons do these central atoms have around them?

Exception 3: Expanded Octet

Slide 106 / 186

Draw the Lewis dot structure for sulfur hexaflouride, SF6:


Move for answer

Slide 107 / 186

Draw the Lewis dot structure for the xenon tetrafluoride, XeF4.


Move for answer

Slide 108 / 186


Draw the Lewis dot structure for boron trifluoride, BF3:

Move for answer

Slide 109 / 186

Draw the Lewis dot structure for the iodine tricholoride, ICl3.


Cl - I - Cl

ClMove for answer

Slide 110 / 186

37

A Boron and Beryllium

B Boron and Helium

C Boron, Beryllium, and Hydrogen

D Boron, Beryllium, Hydrogen and Helium

E Boron, Beryllium, Hydrogen, Helium and Oxygen

[*] Which of the following need fewer than 8 valence electrons to be stable?


37

A Boron and Beryllium

B Boron and Helium

C Boron, Beryllium, and Hydrogen

D Boron, Beryllium, Hydrogen and Helium

E Boron, Beryllium, Hydrogen, Helium and Oxygen

[*] Which of the following need fewer than 8 valence electrons to be stable?


Ans

wer

C

Slide 111 / 186

38 The correct lewis structure for BeCl2 is

Cl - Be - Cl

True

False


38 The correct lewis structure for BeCl2 is

Cl - Be - Cl

True

False


Ans

wer

False

Slide 112 / 186

39 Elements in the first two rows of the periodic table cannot have expanded octets because their atoms do not have enough space.

True

False


39 Elements in the first two rows of the periodic table cannot have expanded octets because their atoms do not have enough space.

True

False


Ans

wer

True

Slide 113 / 186


Resonance Structures

Slide 114 / 186

C

NCl

F

OSB

P

I

H

Si

SeXe

O3Draw a Lewis Structure and use that to determine the VSEPR number

For the central oxygen:Electron domains = 3Bonding domains = 2

Unpaired electrons = 1

Its VSEPR number is 3 2 1

OO

OSlide for Answer

Slide 115 / 186

Consider the Lewis structure we would draw for ozone, O3:

We would expect the double bond to have a shorter bond length than the single bond.

However, the true, observed structure of ozone shows that both O-O bonds are the same length. How can this be?

Resonance

O

OO

O

O

O

[*]

Slide 116 / 186

One Lewis structure cannot accurately depict a molecule like ozone. Therefore, we use multiple structures, called resonance structures, to describe the molecule.

Ozone has two resonance structures.

Resonance

O

O

OO

O

O

[*]

Slide 117 / 186

ResonanceThe actual ozone molecule is a synthesis of these two resonance structures.

The bond length for both outer oxygen atoms falls somewhere between the single and double bond length.

O

O

OO

O

O

Resonancestructure

Resonancestructure

Ozone molecule

[*]

Slide 118 / 186

Resonance

The nitrate ion, NO31- also requires resonance structures to explain

its covalent bonding.

There are three resonance structures for the nitrate ion:

[*]

page53svg

Slide 119 / 186

Draw the Lewis dot structure for SO3:

Resonance Structures

move for answer

[*]

Slide 120 / 186

40 How many resonance structures can be drawn for the carbonate ion, CO32- ?

A 1B 2C 3

D 4E 5

[*]


40 How many resonance structures can be drawn for the carbonate ion, CO32- ?

A 1B 2C 3

D 4E 5

[*]


Ans

wer

C

Slide 121 / 186

The benzene molecule is a regular hexagon of carbon atoms with a hydrogen atom bonded to each one. There are two resonance structures for benzene.

Benzene

Benzene, C6H6, is obtained from the distillation of fossil fuels. More than 4 billion pounds of benzene is produced annually in the United States. Because benzene is a carcinogen, its use is closely regulated.

[*]

Slide 122 / 186

Localized v. Delocalized electronsIn truth, the shared pairs of electrons do not always remain between adjacent C atoms. They are not localized.

Instead, the electrons are said to be delocalized, meaning that they they can move around the 6-carbon ring.

Benzene is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring... we will talk more about this at the end of the year when we study organic chemistry.

# # # # or

[*]

Slide 123 / 186


VSEPR Theory

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Slide 124 / 186

VSEPR Theory

Valence Shell Electron Pair Repulsion

According to VSEPR theory, the molecules will adopt a shape/geometry so as to reduce the repulsion between the bonded electrons.

Click here to view a PhET simulation

Slide 125 / 186

The VSEPR number of a molecule is a three digit number that can be used to determine a molecule's shape.

Here's how you find it:

1. Draw the Lewis structure for the molecule. Locate the central atom, if applicable.

2. The first digit of the VSEPR number is the total number of electron-domains around the central atom.

VSEPR Numbers

Electron domains are either shared pairs of electrons or lone pairs of electrons

Multiple bonds (i.e. double or triple bonds) count as only ONE electron domain.

Slide 126 / 186

3. The second digit of the VSEPR number is the total number of bonding-domains around the central atom.

4. The third digit of the VSEPR number is the total number of lone pairs around the central atom.

5. Check your work - the first digit is equal to the sum of the second and third.

VSEPR Numbers (cont)

Bonding domains are single, double or triple bonds.

Each pair of electrons that are not involved in bonds counts as one lone pair.

Slide 127 / 186

41 How many electron domains does CH4 have?

A 1B 2C 3D 4E 5


41 How many electron domains does CH4 have?

A 1B 2C 3D 4E 5


Ans

wer

D

Slide 128 / 186

42 How many electron domains does H2O have?

A 1B 2C 3D 4E 5

H H

O

http://phet.colorado.edu/en/simulation/molecule-shapes


42 How many electron domains does H2O have?

A 1B 2C 3D 4E 5

H H

O


Ans

wer

D

Slide 129 / 186

43 How many electron domains does CO2 have?

A 1B 2C 3D 4E 5

C OO


43 How many electron domains does CO2 have?

A 1B 2C 3D 4E 5

C OO


Ans

wer

B

Slide 130 / 186

C

NCl

F

OSB

P

I

H

Si

SeXe

CH4Draw a Lewis Structure and use that to determine the VSEPR number

H

H

HC

H

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #.

Electron domains = 4Bonding domains = 4Lone pairs of electrons = 0


Slide 131 / 186

C

NCl

F

OSB

P

I

H

Si

SeXe

NF3

Draw a Lewis Structure and use that to determine the VSEPR number

N F

F

F




Slide for Answer

Slide 132 / 186

C

NCl

F

OSB

P

I

H

Si

SeXe

SiF4Draw a Lewis Structure and use that to determine the VSEPR number

F

Si

F

F F




Slide for Answer

Slide 133 / 186

C

NCl

F

OSB

P

I

H

Si

SeXe

PO43-Draw a Lewis Structure and use that to determine the VSEPR number

O

P

O

O O




Slide for Answer

Slide 134 / 186

FC

NCl

F

OSB

P

I

H

Si

SeXe

IF5Draw a Lewis Structure and use that to determine the VSEPR number

FI

FF F




Slide for Answer

Slide 135 / 186


Molecular Geometry

Slide 136 / 186

VSEPR and molecule shape prediction

According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.

The shape of a molecule plays an important role in determining its chemical and physical properties.

To determine a molecule's shape, i.e. its molecular geometry, we must first determine its electron-domain geometry.

Slide 137 / 186

Recall:

Electron domains are either shared pairs of electrons or lone pairs of electrons

Bonding domains are single, double or triple bonds.

Each pair of electrons that are not involved in bonds counts as one lone pair.

To determine the electron-domain geometry, look at the first number and use the following chart...

How does VSEPR theory help predict the shapes of molecules?

Slide 138 / 186

Electron Domain Geometry

page53svg

Slide 139 / 186

Electron-Domain Geometry (EDG)

The EDG (2,3,4,5,or 6) gives us the general shape of the molecule, as shown here.

However, these domains do not have to be bonds.

The molecular geometry tells us if there is a bond or lone pair of electrons present, thereby specializing the general shape.

Let's take a closer look...

Slide 140 / 186

Linear Electron-Domain Geometry

Linear

Two atoms around a central one will form a linear shape with bond angles of 180o

Slide 141 / 186

Linear Molecular Geometry

There is only one molecular geometry for linear electron-domain: linear molecular geometry (220).

Slide 142 / 186

Trigonal Planar Electron-Domain Geometry

trigonal planar

Three atoms around a central one will form a trigonal planar shape with bond angles of 120o

Slide 143 / 186

Trigonal Planar Molecular Geometry

There are two molecular geometries:· Trigonal planar, if all the electron domains are bonding (330)

· Bent, if one of the domains is a nonbonding pair (321)

Slide 144 / 186

120

trigonal planar(330)

117

bent(321)

Trigonal Planar Molecular Geometry

It is very important to note that unbonded pairs of electrons repel more strongly than bonded electrons thereby shrinking the bond

angle between atoms

Slide 145 / 186

Tetrahedral Electron-Domain Geometry

Four atoms around a central one will form a tetrahedral shape with bond angles of 109.5o

tetrahedral

Slide 146 / 186

Tetrahedral Molecular Geometry

There are three molecular geometries:Tetrahedral, if all are bonding pairs (440)

Trigonal pyramidal, if one is a nonbonding pair (431)Bent, if there are two nonbonding pairs (422)

Slide 147 / 186

Tetrahedral Molecular Geometry

tetrahedral (440)

trigonal pyramidal(431)

bent (422)

109.5

107104.5

Again, note the decrease in bond angle as the number of high repelling unbonded pairs of electrons increase.

Slide 148 / 186

Five atoms around a central one will form a trigonal bipyramidal shape with bond angles of 120o and 90o

trigonal bipyramidal

Trigonal Bipyramidal Electron-Domain Geometry

Slide 149 / 186

Trigonal Bipyramidal Molecular Geometry

Trigonal bipyramidal

Seesaw

T-shaped

Linear

Slide 150 / 186

Trigonal Bipyramidal

(550)See-Saw

(541)T-Shape

(532)Linear (523)

Trigonal Bipyramidal Molecular Geometry

There are four molecular geometries for the trigonal bipyramidal electron domain geometry:

Slide 151 / 186

Six atoms around a central one will form an octahedral shape with bond angles of 90o

octahedral

Octahedral Electron-Domain Geometry

Slide 152 / 186

Octahedral Molecular Geometry

Square Planar

Octahedral

Square Pyramidal

Slide 153 / 186

Octahedral (660)

Square Pyramidal

(651)

Square Planar (642)

Octahedral Molecular Geometry

There are only three molecular geometries for the octahedral electron domain geometry:

Slide 154 / 186

VSEPR and molecular geometry

Using VSEPR numbers, you can determine molecular geometry.

VSEPR numbers are a set of 3 numbers.

1) the total number of electron domains2) the number of bonding domains*3) the number of unshared pairs of electrons

Electron-domain geometry has the same name as the first shape.

(*Remember that multiple bonds count as ONE domain)

Slide 155 / 186

Draw the Lewis structure for ammonia, NH3.

What are the VSEPR numbers for NH3? 4,3,1

What is the electron-domain geometry of NH3? tetrahedral

What is the molecular shape of NH3?###

What is the N-H bond angle in the molecule? 107

What is the formal charge on the N atom? 5-5 = 0

VSEPR Numbers and Molecular Geometries

triangular pyramidal

slide for answer

Slide 156 / 186

Draw the Lewis structure for ClF3.

What are the VSEPR numbers for ClF3? 5,3,2

What is the electron-domain geometry of ClF3? trigonal bipyramidal

What is the molecular shape of ClF3? T

What would be the Cl-F bond angle(s)? 180, 90

What would be the formal charge on Cl? 7-7 = 0

VSEPR Numbers and Molecular Geometries

slide for answer

Slide 157 / 186

44 The methane molecule (CH4) has which geometry?

A linear

B trigonal bipyramidal

C trigonal planar

D tetrahedral


44 The methane molecule (CH4) has which geometry?

A linear

B trigonal bipyramidal

C trigonal planar

D tetrahedral


Ans

wer

D

Slide 158 / 186

45 Give the VSEPR number for this molecule.[*]


45 Give the VSEPR number for this molecule.[*]


Ans

wer

541

Slide 159 / 186

46 Give the VSEPR number for this molecule.




Ans

wer

642

Slide 160 / 186


F Xe F



F Xe F


Ans

wer

523

Slide 161 / 186

48 Which compound below contains an atom that is surrounded by more than an octet of electrons?

A PF5

B CH4

C NBr3

D OF2

[*]


48 Which compound below contains an atom that is surrounded by more than an octet of electrons?

A PF5

B CH4

C NBr3

D OF2

[*]


Ans

wer

A

Slide 162 / 186

49 Which of the following molecules would have a bent shape?

A SO2

B SO3

C CH4

D C2H2

E HF


49 Which of the following molecules would have a bent shape?

A SO2

B SO3

C CH4

D C2H2

E HF


Ans

wer

A

Slide 163 / 186

50 Which of the following molecules would have a 104.5 degree bond angle between atoms?

A H2SB CF3ClC CO2

D PCl3

E NO3-


50 Which of the following molecules would have a 104.5 degree bond angle between atoms?

A H2SB CF3ClC CO2

D PCl3

E NO3-


Ans

wer

A

Slide 164 / 186

51 The molecular shape and geometry of the nitrate ion (NO3-) would be:

A bentB linearC trigonal planarD trigonal bipyramidalE tetrahedral


51 The molecular shape and geometry of the nitrate ion (NO3-) would be:

A bentB linearC trigonal planarD trigonal bipyramidalE tetrahedral


Ans

wer

C

Slide 165 / 186

According to carbon's orbital diagram, it should only be able to form two bonds...

__ __ __ __ __

1s 2s 2p

HYBRIDIZATION THEORY

But we know carbon forms 4 bonds, not 2!!!

[*]

Slide 166 / 186

Scientists propose that the outermost s and p orbitals are actually combined to create 4 "hybrid" orbitals of equal energy.

Carbon __ ___ ___ ___ ___

1s sp3 hybrid orbitals

This explained how carbon could form 4 bonds

HYBRIDIZATION THEORY[*]

Slide 167 / 186

To predict the hybridization involved in a compound, simply look at the first VSEPR numbers, this tells you how many electron

domains(orbitals) need to be hybridized.

For example: = 4 electron domains

sp3

Carbon requires 4 hybrid orbital so it hybridizes it's outermost "s" orbital and all three of the "p" orbitals to give 4 sp3 hybrids.


Slide 168 / 186

Example: Find the hybridization of the N atom in NH3?

VSEPR # = 4 so the hybridization is sp3


Slide 169 / 186

Example: What is the hybridization of C in CO2?

VSEPR # = 2 Only the s orbital and 1 p orbital are needed to be hybridized so the hydridization is sp

Note: The other 2 p orbitals not involved in hybridization are used to form the double bonds (called Pi bonds)


Slide 170 / 186

52 Which of the following would require sp2 hybridization?[*]

A BF3

B H2OC PCl3D F2

E N2


52 Which of the following would require sp2 hybridization?[*]

A BF3

B H2OC PCl3D F2

E N2


Ans

wer

A

Slide 171 / 186

53 What would be the hybridization found on O in OF2?[*]

A spB sp2

C sp3

D s2p3

E s3p3


53 What would be the hybridization found on O in OF2?[*]

A spB sp2

C sp3

D s2p3

E s3p3


Ans

wer

C

Slide 172 / 186


Polarity

Slide 173 / 186

Polarity of BondsThough atoms often form compounds by sharing electrons, the electrons are not always shared equally. In a covalent bond, one atom has a greater ability to pull the shared pair toward it.

Slide 174 / 186

Polarity of Bonds

Identical atoms will have an electronegativity difference of ZERO. As a result, the bond is NONPOLAR.

Slide 175 / 186

Bonds and Electronegativity

Bond Type

Non-Polar Covalent

Polar Covalent

Ionic

Electronegativity Difference

very small or zero

about 0.2 to 1.6

above 1.7 (between metal & non-metal)

Slide 176 / 186

Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

Polarity of Bonds

H F

We use the symbol to designate a dipole (2 poles). The "+" end is on the more positive end of the molecule and the arrow points towards the more negative end.

page53svg

Slide 177 / 186 Slide 178 / 186

Polarity of Bonds

Compound Bond Electronegativity Dipole length (A0) Difference Moment (D)

HF 0.92 1.9 1.82HCl 1.27 0.9 1.08HBr 1.41 0.7 0.82HI 1.61 0.4 0.44

Bond lengths, Electronegativity, Differences and Dipole Moments of the Hydrogen Halides

Slide 179 / 186

But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

Polarity of Molecules

For instance, in the case of CO2:

The polar bond is shown as a dipole, the arrow points to the more negative atom. Dipoles add as vectors.

[*]

Slide 180 / 186


By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

For a molecule to be polar, it must a) contain one or more dipoles AND b) have these polar bonds arranged asymmetrically

[*]

In other words, if all the dipoles are symmetrical, they will cancel each other out and the molecule will be

NONPOLAR.

Many molecules with lone pairs of electrons will be POLAR.

Slide 181 / 186

These are some examples of polar & nonpolar molecules. What are their VSEPR numbers?


330, nonpolar

440, nonpolar

440, polar

431, polar110(?), polarSlide for Answer

[*]

Slide for Answer

Slide for Answer

Slide for Answer

Slide for Answer

Slide 182 / 186

54 Which of these are polar molecules?

A a, bB a, b, cC a, cD a, c, d

E c, e

[*]


54 Which of these are polar molecules?

A a, bB a, b, cC a, cD a, c, d

E c, e

[*]


Ans

wer

C

Slide 183 / 186

55 Sulfur trioxide (SO3) is polar.

TrueFalse


55 Sulfur trioxide (SO3) is polar.

TrueFalse


Ans

wer

False

Slide 184 / 186

56 Hydrogen sulfide gas (H2S) is non-polar.

TrueFalse


56 Hydrogen sulfide gas (H2S) is non-polar.

TrueFalse


Ans

wer

True

Slide 185 / 186

57 Which of the following contains polar bonds but is a non-polar molecule?

A CH4

B CS2

C H2SD CF4

E All of these are polar


57 Which of the following contains polar bonds but is a non-polar molecule?

A CH4

B CS2

C H2SD CF4

E All of these are polar


Ans

wer

D

Slide 186 / 186

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