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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.

Slide 3 / 186 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

Slide 4 / 186

Return to Table of Contents

Covalent versus Ionic Bonds

Slide 5 / 186 Covalent Bonding & Molecular Geometry

Examine these two forms of the same compound, ibuprofen.

Slide 6 / 186 Covalent Bonding & Molecular Geometry

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.

Slide 7 / 186

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.

Covalent Bonding & Molecular Geometry

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!)

Slide 8 / 186 Chemical Bonds

Ionic - The electrostatic attraction between ions Covalent - The sharing of electrons between atoms Metallic - Each metal atom bonds to

• ther 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:

Slide 9 / 186 Chemical Bonds

How 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 F

Electronegativity 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

Slide 10 / 186 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 animation about ionic and covalent bonding

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Return to Table of Contents

Properties of Ionic and Covalent Materials

Slide 17 / 186 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

Slide 18 / 186 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

+

• +

+ + +

Slide 19 / 186 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 Slide 22 / 186

Copper is often used in electrical cable rather than silver even though it is roughly 10% less conductive than silver. Why?

REAL WORLD APPLICATION

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 Slide 23 / 186

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

A Al B Al2O3 C NaCl D

Both A and B

E

Both B and C

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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

REAL WORLD APPLICATION

slide for answers

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

A NaCl B HF C CO2 D

Ge2O3

E

Fe

Slide 28 / 186 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

Slide 31 / 186 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

W

E

LiF

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

A Cu B NaF C Fe D

CO2

E

All of these will conduct

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9 Which of the following consists of small individual molecules?

A C(diamond) B SiO2 C Cu2O D

Na

E

SO3

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

A Cu B CuCO3 C LiCl D

Ba

E

BaF2

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Return to Table of Contents

Naming Binary Molecular Compounds

Slide 37 / 186 Naming Binary Molecular Compounds

Use prefixes to indicate the number the atoms. All end in "ide" Examples NO2 nitrogen dioxide P2O5 diphosphorous pentoxide ( penta-oxide-->pentoxide)

Slide 38 / 186 Naming Binary Molecular Compounds

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

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12 Dinitrogen tetroxide is A NO2 B

N2O4

C

NO3-

D N4O2

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13 H2O is

A

Hydrogen monoxide B

Dihydrogen monoxide

C Hydrogen oxide D Hydrogen dioxide

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14 SO3 is

A sulfate B sulfur oxide C sulfur trioxide D sulfite

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15 MgO is

A monomagnesium monoxide

B magnesium monoxide C monomagnesium oxide

D magnesium oxide

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16 P4O10 is

A Phosphorous pentoxide B

Tetraphosphorous decoxide

C

Phosphorous oxide

D Phosphate

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Return to Table of Contents

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 Slide 47 / 186

17 How many valence electrons does nitrogen have?

A 2 B 3

C

4

D

5 E 7

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18 The Lewis structure for nitrogen is N True False

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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

• btain 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-

Slide 50 / 186 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

• f electrons

H H

1s 1s

Hydrogen molecule

Slide 51 / 186 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.

H H

Hydrogen molecule

Shared pair

• f 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

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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

Fluorine atom Fluorine molecule Fluorine atom

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 H H H

Hydrogen atoms Oxygen atom Water molecule

1s 2p 2s 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 H

Hydrogen atom Nitrogen atom Ammonia molecule 1s 2p 2s 1s 1s

H N H

1s Ammonia molecule

H

Slide 56 / 186 Slide 57 / 186 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 Cl Cl

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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
• uter shell (8 except, 2 for

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

Drawing Lewis Structures Slide 59 / 186 Drawing Lewis Structures

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.

Slide 60 / 186 Drawing Lewis Structures

• 2. The central atom is the least

electronegative element (excluding hydrogen because it can only have one bond).

• 3. Connect the other atoms to it

by single bonds. 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.

H N H H

NH3

Slide 61 / 186 Drawing Lewis Structures

H N H H

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

H N H H

• 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 H

2O have?

A B C D

8 10 12 14 Slide 63 / 186

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

A B

H O Slide 64 / 186

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

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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

Slide 66 / 186 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

Slide 67 / 186

24

A B

12 18

How many valence electrons does CO3

2- have?

C D

24 26 Slide 68 / 186

25

A B

8 9

How many valence electrons does H3O+ have?

C D

10 11 Slide 69 / 186 Formal Charge

The "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 O FC 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 PO4

3-

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.

Slide 70 / 186 Formal Charge

The 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.

Slide 71 / 186

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

Slide 72 / 186

26 Which of the following would be the formal charge

• n the N in the ammonium ion?

A +1 B 0 C -1 D -2 E -3

Slide 73 / 186

27 In which of the following molecules would N carry a

non-zero formal charge?

A

HCN

B

NH3

C

NO3-

D

NO2-

E

NH4+

Slide 74 / 186 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 Slide 75 / 186

Lewis Structures

Draw 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.

Slide 76 / 186

C N Cl F O S B P I H

C O O

Si Se Xe

CO2

Draw a Lewis Structure We ran out of electrons, but carbon does not have an octet yet!

Now What?

Slide for Answer

Slide 77 / 186 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.

Slide 78 / 186

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 O

• 2. Form Single Bonds

Slide 79 / 186 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
• ur 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 Slide 80 / 186 Covalent Bond Length Slide 81 / 186 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

Slide 82 / 186 Covalent Bond Energies Slide 83 / 186 Covalent Bonds Comparison

Type of Bond Electrons shared Bond Strength Bond Length 2 4 6 weak intermediate strong long intermediate short

Slide 84 / 186

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

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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

D

varies, depending on the atoms

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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

D

varies, depending on the atoms

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31 How many electrons are shared by two

atoms to create a single bond?

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32 How many electrons are shared by two

atoms to create a double bond?

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33 How many electrons are shared by two

atoms to create a triple bond?

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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 Cl2 D

CO

E

I2

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35 Which of the following molecules would have the

longest C-O bond length? Use Lewis structures.

A CO B CO2 C H2CO D CH3OH E The lengths are all the same

Slide 92 / 186

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

Oxygen atom Oxygen atom Oxygen molecule

Oxygen molecule

Slide 94 / 186

C N Cl F O S B P I H

C

Si Se Xe

CO

Draw a Lewis Structure Carbon has the lower electronegativity, so we will consider it the "central" atom... O

Slide for Answer

Slide 95 / 186 Coordinate Covalent Bonds Slide 96 / 186 Coordinate Covalent Bonds

In carbon monoxide, oxygen has a stable configuration but the carbon does not.

1s 2p 2s 2s 1s 2p

C + O # # # C O

Carbon atom Oxygen atom Carbon monoxide

C O

Carbon 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 N Cl F O S B P I H Si Se Xe

F2

Draw a Lewis Structure

F F Slide for Answer F

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!

Hydrogen Nitrogen Oxygen Fluorine Chlorine Bromine Iodine

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 Exceptions to the Octet Rule

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:

Exceptions to the Octet Rule Move for answer Slide 107 / 186

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

Exceptions to the Octet Rule Move for answer Slide 108 / 186 Exceptions to the Octet Rule

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.

Exceptions to the Octet Rule Cl - I - Cl Cl Move 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?

Slide 111 / 186

38 The correct lewis structure for BeCl2 is

Cl - Be - Cl

True 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

Slide 113 / 186

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Resonance Structures

Slide 114 / 186

C N Cl F O S B P I H Si Se Xe

O3

Draw a Lewis Structure and use that to determine the VSEPR number For the central oxygen: Electron domains = 3 Bonding domains = 2 Unpaired electrons = 1 Its VSEPR number is 3 2 1

O O O Slide 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 O O 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 O O O O

[*]

Slide 117 / 186 Resonance

The 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 O O O O

Resonance structure Resonance structure

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:

[*]

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 1 B

2

C 3 D 4

E

5

[*]

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 electrons

In 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

• rganic chemistry.

# # # #

• r

[*]

Slide 123 / 186

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VSEPR Theory

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

• r lone pairs of electrons

Multiple bonds (i.e. double or triple bonds) count as

• nly ONE electron domain.

Slide 126 / 186

• 3. The second digit of the VSEPR number is the total number
• f 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 1 B 2 C 3 D 4 E 5

Slide 128 / 186

42 How many electron domains does H2O have?

A 1 B 2 C 3 D 4 E 5

H H O Slide 129 / 186

43 How many electron domains does CO2 have?

A 1 B 2 C 3 D 4 E 5

C O O

Slide 130 / 186

C N Cl F O S B P I H Si Se Xe

CH4

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

H H H C H

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Its VSEPR number is 4 4 0

Slide 131 / 186

C N Cl F O S B P I H Si Se Xe

NF3

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

N F F F

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 3 Lone pairs of electrons = 1 Its VSEPR number is 4 3 1

Slide for Answer

Slide 132 / 186

C N Cl F O S B P I H Si Se Xe

SiF4

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

F Si F F F

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Its VSEPR number is 4 4 0

Slide for Answer

Slide 133 / 186

C N Cl F O S B P I H Si Se Xe

PO43-

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

O P O O O

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Its VSEPR number is 4 4 0

Slide for Answer

Slide 134 / 186

F C N Cl F O S

B P I H Si Se Xe

IF5

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

F I F F F

Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 6 Bonding domains = 5 Lone pairs of electrons = 1 Its VSEPR number is 6 5 1

Slide for Answer

Slide 135 / 186

Return to Table of Contents

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

• r 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

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)

109.5 107 104.5

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

• ctahedral

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 domains 2) the number of bonding domains* 3) the number of unshared pairs

• f 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

Slide 158 / 186

45 Give the VSEPR number for this molecule.

[*]

Slide 159 / 186

46 Give the VSEPR number for this molecule.

Slide 160 / 186

47 Give the VSEPR number for

this molecule.

F Xe F

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

[*]

Slide 162 / 186

49 Which of the following molecules would have a bent

shape?

A SO2 B SO3 C CH4 D

C2H2

E

HF

Slide 163 / 186

50 Which of the following molecules would have a 104.5

degree bond angle between atoms?

A H2S B CF3Cl C CO2 D

PCl3

E

NO3-

Slide 164 / 186

51 The molecular shape and geometry of the nitrate ion

(NO3-) would be:

A bent B linear C trigonal planar D

trigonal bipyramidal

E

tetrahedral

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"

• rbital and all three of the "p" orbitals to give 4 sp3 hybrids.

HYBRIDIZATION THEORY

[*]

Slide 168 / 186

Example: Find the hybridization of the N atom in NH3? VSEPR # = 4 so the hybridization is sp3

HYBRIDIZATION THEORY

[*]

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)

HYBRIDIZATION THEORY

[*]

Slide 170 / 186

52 Which of the following would require sp2 hybridization?

[*]

A BF3 B H2O C PCl3 D

F2

E

N2

Slide 171 / 186

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

[*]

A sp B sp2 C sp3 D

s2p3

E

s3p3

Slide 172 / 186

Return to Table of Contents

Polarity

Slide 173 / 186 Polarity of Bonds

Though 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.

Slide 177 / 186 Slide 178 / 186 Polarity of Bonds

Compound Bond Electronegativity Dipole length (A0) Difference Moment (D) HF 0.92 1.9 1.82 HCl 1.27 0.9 1.08 HBr 1.41 0.7 0.82 HI 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 Polarity of Molecules

By adding the individual bond dipoles,

• ne 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

• ther 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?

Polarity of Molecules

330, nonpolar 440, nonpolar 440, polar 431, polar 110(?), polar

Slide 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, b

B

a, b, c

C

a, c D

a, c, d

E

c, e

[*]

Slide 183 / 186

55 Sulfur trioxide (SO3) is polar.

True False

Slide 184 / 186

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

True False

Slide 185 / 186

57 Which of the following contains polar bonds but is a

non-polar molecule?

A CH4 B CS2 C H2S D

CF4

E

All of these are polar

Slide 186 / 186