Chemical BondingChapter 3

Bonding Theory & Lewis FormulasSection 3.1 (pg. 78-84)

Homework: Lewis Symbols Pg. 82 #2 – 4 Pg. 84 # 2, 4, 5, 7-10

Objectives:1) Define valence electron, electronegativity, and ionic bond

2) Use the Periodic Table and Lewis structures to support and explain ionic bonding

3) Explain how an ionic bond results from the simultaneous attraction of oppositely charged ions.

Bonding Theory: Valence Electrons & Orbitals

To describe where electrons exist in the atom, chemists created the concept of an orbital.

Orbital – region of space around an atom’s nucleus where an electron may exist

An “orbital” is not a definite race track, it is a 3-D space that defines where an electron may be (like a rain drop in a cloud)

For bonding study we are only concerned with an atom’s valence orbitals (the volume of space that can be occupied by electrons in an atom’s highest energy level)

WHY? Bonding only involves valence e-’s because lower energy levels are held so strongly by their positively charged nucleus

FYI Read pg. 78-79 for the history on Bonding Theory

According to bonding theory, valence electrons are classified in terms of orbital occupancy.

(0 = empty, 1 = half filled , 2 = full)

An atom with a valence orbital that has a single electron can theoretically share that electron with another atom Such an electron is called a BONDING

ELECTRON

An atom with a full valence orbital (2 e-’s), repels nearby orbitals and wants to be alone Such a pairing is called a LONE PAIR

Bonding Theory: Valence Electrons & Orbitals

1. The first energy level has room for only one orbit

- can only hold 2 e-’s max

2. Energy levels above the first have room for four orbitals = 8 electrons max

- Noble gases have this structure; their lack or reactivity indicates that eight electrons filling a valence orbital is very stable

(Remember the OCTET RULE)

The Four Rules of Bonding Theory 2e-

2 p+

He

8e-

8e-

2e-

18 p+

Ar

2e-

2e- 2e- 2e- 2e-

FYI: Only C, N, O, and F atoms always obey the octet rule when bonding

EXCEPTIONS:B = stable with 6 valence e- (3 orbitals)P = stable with 10 valence e- (5 orbitals)S = stable with 12 valence e- (6 orbitals)

The Four Rules of Bonding Theory 3. An orbital can be unoccupied, or it may contain one or two electrons – but never more than two (Pauli Exclusion Principle)

4. Electrons “spread out” to occupy any empty valence orbitals before forming electron pairs

Never more than 2e- in an orbital

“Aluminum has three half-filled orbitals and one vacant orbital.” How would you describe Sulfur?

Atomic Models: LEWIS SYMBOLS(aka Lewis Dot Diagrams, Electron Dot Diagrams, LDD, Lewis Models

• Named after Gilbert Lewis who is responsible for the Octet Rule. He

reasoned that all atoms strive to be like the nearest noble gas.

• Used dots or ‘x’ to represent the valence electrons

• The inner electrons and the nucleus are represented by the element symbolHow to draw Lewis Symbols:

1. Write the element symbol2. Add a dot to represent each valence e-

3. Start by placing valence e-’s singly into each of the four valence orbitals (4 sides)

4. If additional e-’s need to be placed, start filling each of the orbitals with a second e- up to 8

Q: Which element has 4 bonding e-’s? Which has 3 lone pairs and 1 bonding e-?

PracticeDraw the Lewis Symbols for the elements in Period 3

For each one indicate how many lone pairs or bonding electrons are present

It is important to remember that the Lewis symbols do not mean that electrons are dots or that they are stationary.

The four sides represent the four orbitals that may be occupied by electrons; it is a simplistic 2-D diagram of a complex 3-D structure

Electronegativity- A measure of the force that an atom exerts on electrons of other atoms; (the “pull” on bonding

electrons)

- Each atom is assigned a value between 0.0 – 4.0; the larger the number

the greater the “pulling” force

- Example: Fluorine has an EN = 4.0 and francium has an

EN = 0.7- This means fluorine wants to pull on other electrons very

strongly- This means francium doesn’t want to pull on other

electrons

Q: Does lithium (EN = 1.0) want to lose or gain an electron to be stable?

Q: Does fluoride (EN = 4.0) want to lose or gain an electron to be stable?

Do you see any relation to their electronegativity numbers?

So how do we assign each atom an electronegativity number?

a) The farther away from the nucleus that electrons are, the weaker their attraction to the nucleus

b) Inner electrons shield valence electrons from the attraction of the positive nucleus

c) The greater the number of protons in the nucleus, the greater the attraction for more electrons

EN = 0.8

EN = 2.6

Cesium's valence electrons are not held as tightly by its nucleus because the atom is larger

EN = 0.8

EN = 3.0

Potassium’s valence electrons are not attracted to its nucleus as much as Nitrogen’s valence electrons because their are more inner electrons present in K

1 e-8e-8e-2e-19p+K

5e-2e-7p+N

EN =1.9

EN = 3.0

14p+Si

35p+Br

Bromine has more protons (+ charge) which attracts the negative charge of electrons more so than silicon’s 14 protons

.

In this 3-D image, the electronegativity scale is vertical.

Q: Which element has the highest EN? Give three reasons why?

Electronegativity

Q: What is the EN trend within a period and a group?

Why do we care about electronegativity??

Imagine that two atoms, each with an orbital containing one bonding electron, collide in such a way that these half-filled orbitals overlap.

As the two atoms collide, the nucleus of each atom attracts and attempts to “capture” the bonding electrons of the other atoms

A “Tug of War” over the bonding electrons occurs

Which atom wins?By comparing the electronegativities of the two atoms

we can predict the result of the contest = 3 different types of bonds result

Covalent BondingBoth atoms have a high EN so neither atom “wins”

The simultaneous attraction of two nuclei for a shared pair of bonding electrons = covalent bond

EN difference can be zero = Cl – Cl EN = 3.2 EN = 3.2

EN difference can be small = H - Cl EN = 2.2 EN = 3.2

This is called a polar covalent bond – because one side pulls on the electrons more but we will learn more about this in Section 3.3

Cl2 = diatomic

Ionic BondingThe EN of the two atoms are quite different

The atom with the higher EN will remove the bonding e- from the other atom electron transfer occurs

Positive and negative ions are formed which electrically attract each other

EN = 0.9

EN = 3.2

Metallic Bonding Both atoms have a relatively low EN so atoms

share valence electrons, but no actual chemical reaction takes place

In metallic bonding:a) e-’s are not held very strongly by their

atomsb) the atoms have vacant valence orbitals

- This means the electrons are free to move around between

the atoms and the (+) nuclei on either side will attract them

Analogy: The positive nuclei are held together by a glue of

negative e-’s

Metallic bonding visual This diagram represents Mg atoms that have released their electrons and are embedded in a sea (or glue) of electrons.

Note: These metal atoms don`t have to be in a particular arrangement to attract each other therefore they are flexible, malleable and ductile = useful alloys (Brass, Stainless Steel, etc.)

Summary of Bonding Theory:Chemical Bond = competition for bonding electrons

1) Atoms with equal EN = electrons shared equally

If both have high EN = covalent bond (equal = non-polar)

If both have a low EN = metallic bond

2) Atoms with unequal EN = covalent bond (unequal = polar)

3) Atoms with unequal EN = ionic bond

metallic

The nature of chemical bonds changes in a continuous way,

creating a broad range of characteristics.

PRACTICE

Copy pg. 84 – Bonding Theory Summary into your Notes

Pg. 82 #2 - 4

Pg. 84 # 2, 4, 5, 7-10

Section 3.2 (pg. 85-90)

Explaining Molecular Formulas

Pg. 89 #5 (a-f), 6 (a-e)

Pg. 90 #1-4, 6

Section 3.2 (pg. 85-90)

Explaining Molecular Formulas

Objectives:

1) Draw electron dot diagrams of atoms and molecules, writing structural formulas for molecular substances using Lewis structures to predict bonding in simple molecules

2) Illustrate, by drawing or building models, the structure of simple molecular substances

3) Explain why the formulas for molecular substances refer to the number of atoms of each constituent element

Molecular ElementsMany molecular elements are

diatomic and some are polyatomic

You will need to memorize the formulas of the 9 molecular elements as they will not be given to you:

Name Symbol

hydrogen H2(g)

nitrogen N2(g)

oxygen O2(g)

fluorine F2(g)

chlorine Cl2(g)

iodine I2(g)

bromine Br2(g)

phosphorous

P4(g)

sulfur S8(g)

Why are they diatomic?Remember fluorine has 7 valence e-’s and

needs 1 more e- to be stable?Well 2 fluorine atoms could obtain a stable

octet of e-’s if they shared a pair with each other

F - F F2

Remember: This is a simplified 2-D version, not where the electrons actually are

In structural formulas, lone pairs

are not shown

Diatomic ElementsWhat about oxygen and nitrogen?

Each oxygen atom only has 6 valence electrons. So by sharing 2 electrons each, the two oxygen atoms can create a full octet. This creates a double bond

O = O

Each nitrogen atom only has 5 valence electrons. So by sharing 3 electrons each, the two nitrogen atoms can create a full octet. This creates a triple bond

N Ξ N

Would sharing only 1 electron each work?

Molecular CompoundsBackground:

Molecular compounds have covalent bonds (shared electrons) between non-metals and non-metals

Can be solid, liquid or gas as SATP

May or may not be soluble in water (more later)

Don’t ever conduct electricity - even when (aq)

Generally have lower m.p. and b.p than ionic compounds

Review from Section 1.5 Notes

Molecular CompoundsBackground:

Empirical Formulas – show the simplest whole number ratios of atoms in a compound Very useful for IONIC compounds

Formula Unit – the ratio of ions that repeats in a pattern within the crystal; the chemical formula of ionic compounds represents the formula unit

Not useful for MOLECULAR compounds

Na242Cl242 Na16Cl16 NaCl

CH C2H2 acetylene C6H6 benzene C8H8 octene

All are extremely different compounds but the empirical formula would be the same

Molecular CompoundsBackground:

Molecular Formulas – a molecular formula shows the actual number of atoms that are covalently bonded to make-up each molecule

We use this because chemical formulas for molecular compounds result from sharing electrons, therefore a variety of compounds are possible (which we determine empirically through experiments) Often the symbols are written in a sequence that helps you determine how

the atoms are bonded

C2H4O2 CH3COOH

Empirical formula: CH2O incorrect for molecular

compounds

Review of Molecular Compound Formulas

SeePg. 88

Q: How do we know how molecular compounds bond?

(Aka: How do we draw Lewis Formulas?)

Where do these come from???

Determining Lewis Formulas Bonding Capacity: the maximum number of single

covalent bonds that an atom can form

REMINDER: How many e-’s does an atom want in its valence energy level to be satisfied?

H = 2 e- C, N, O, F, P, S, Cl, etc. = 8 e-

REMINDER: What types of covalent bonds are possible?

F – F single = sharing one e- pairO = O double = sharing two e- pairsN Ξ N triple = sharing three e- pairs

So why do we care about bonding capacity?If we know how many bonding e-’s an atom has, we can

predict what structure a molecular compound will have

Determining Lewis Formulas

Atom Number ofvalence electrons

Number ofbonding electrons

Bonding capacity

carbon 4 4 4

nitrogen 5 3 3

oxygen 6 2 2

halogens 7 1 1

hydrogen 1 1 1H

I.e. Carbon can form 4 single bonds, 2 double bonds, 1 triple and 1 single, or 1 double and 2 singles

Determine the Lewis formula and structural formula for sulfur trioxide, SO3(g)

1. Count the number of valence electrons there are in total? (If polyatomic ions are included, subtract or add electrons to account for the net charge)

Oxygen = 3 atoms x 6 valence e-’s each = 18 valence e-’s Sulfur = 1 atom x 6 valence e-’s each = 6 valence e-’s

24 valence e-’s

Lewis Formulas- Guided Ex. #1

2. Choose your central atom In our course, we will limit our formulas to ones with one

central atom (unless extra info is provided)

So how do you know which is the central atom? Usually the one in lesser quantity (SO3(g))

OR The one with the higher bonding capacity

Carbon usually – because 4 is the highest bonding capacity

So which is the central atom?

Lewis Formulas – Guided Ex. #1

3. Arrange peripheral atoms around central atom and place one pair of valence e-’s between them

4. Place lone pairs on all peripheral atoms to complete their octet

5. Place any remaining valence e-’s on the central atom as lone pairs.

For this example, all 24 have been assigned

Lewis Formulas – Guided Ex. #1

6. If the central atom’s octet is not complete, move a lone pair from a peripheral atom to a new position between the peripheral and central atom.

7. Show the structural formula but omit lone pairs and replace every bond with a line(Count the electrons around each atom to confirm the octet rule. Each atom should have 8 e-’s around it; exception: H)

Lewis Formulas – Guided Ex. #1

Determine the Lewis formula & structural formula for the nitrate ion, NO3

-

1. Count the valence electrons (*look for a net charge if an ion).

nitrogen = 1 x 5 valence e-’s = 5

oxygen = 3 x 6 valence e-’s = 18 23 + 1 (b/c net charge is -1) = 24

Lewis Formulas – Guided Ex. #2

N

2. Which is the central atom? Nitrogen (in lesser quantity)

3. Arrange peripheral atoms around central atom and place 1 pair of valence e-’s between them

4. Place lone pairs on all peripheral atoms to complete their octet

5. Place any remaining valence e-’s on the central atom as lone pairs.

6. If the central atom’s octet is not complete, move a lone pair from a peripheral atom to a new position between the peripheral and central atom.

7. If the entity is a polyatomic ion, place square brackets around the entire Lewis formula and then write the net charge outside the bracket on the upper left.

Lewis Formulas – Guided Ex. #2

N N

N

PracticePg. 89 #5 (a-f), 6 (a-e)

Watch 5 (f) there is an exception noted. The central atom does not follow the octet rule.

We will go through these answers as a class.

Pg. 90 #1-4, 6

The theory presented today is not absolute – there are exceptions. But rather than presenting a more detailed theory, your textbook will always note such exceptions.

Tomorrow...Molecular Model InvestigationThought Lab InvestigationMorse Code Assignment

VSEPR TheorySection 3.3 – Part A

Pg. 91-96

Objective:

1) Apply VSEPR theory to predict molecular shapes

Stereochemistry – is the study of the 3-D spatial configuration of molecules and how this affects their reactions.

The shape of molecules is determined by the repulsion that happens between electron pairs

The theory behind molecular shapes is called VSEPR Theory (Valence Shell Electron Pair Repulsion)

Molecular Shapes

Solid = in plane of page Dashed = behind (away) Wedge = ahead (toward)

General Rule: Pairs of electrons in the valence shell of an atom

stay as far apart as possible because of the repulsion of their negative charges

The type, number and direction of bonds to the central atom of a molecule determine the shape of the resulting molecule.

So how do we predict these molecular shapes?

VSEPR

We will be using the following compounds to analyze the 6 shapes possible

BeH2(s), BH3(g), CH4(g), NH3(g), H2O(l), HF(g)

To start, draw a Lewis formula for each of the molecules and then consider the arrangement of all pairs of valence electrons.

(Remember – all pairs of valence e-’s repel each other and want to get as far apart as possible)

Using VSEPR to Predict Molecular Shapes

Shape #1 = LinearLewis Formula

Bond Pairs

Lone

Pairs

Total Pair

s

General Formula

Electron Pair Arrangement

Stereochemical Formula

2 0 2 AX2 linearX – A – X

linear

Be

• This Lewis formula indicates that BeH2(s) has two bonds and no lone pairs on the central atom.

• VSPER theory suggests that the two bond pairs will be farthest apart by moving to opposite sides to a bond angle of 180°

• This gives the molecule a linear orientation

* A is the central atom; X is another atom

*Exception* Beryllium does not follow OCTET RULE

Shape #2 = Trigonal PlanarLewis Formula

Bond Pairs

Lone

Pairs

Total Pair

s

General Formula

Electron Pair Arrangement

Stereochemical Formula

3 0 3 AX3trigonal planar

• This Lewis formula indicates that BH3(g) has three bonds and no lone pairs on the central atom.

• VSPER theory suggests that the three bond pairs will be farthest apart by moving to a bond angle of 120° to each other.

• This gives the molecule a trigonal planar orientation.

* A is the central atom; X is another atom

B

*Exception* - Boron Does not follow OCTET RULE

Draw the Lewis Formula for BF3

Practice

Does not obey the octet rule

Trigonal Planar

F

F

F

Shape #3 =TetrahedralLewis

FormulaBond Pairs

Lone

Pairs

Total Pair

s

General Formul

a

Electron Pair Arrangement

Stereochemical Formula

4 0 4 AX4tetrahedr

al

• This Lewis formula indicates that CH4(g) has four bonds and no lone pairs on the central atom.

• VSPER theory suggests that the four bond pairs will be farthest apart by arranging in three dimensions so that every bond makes an angle of 109.5° with each other.

• This gives the molecule a tetrahedral orientation.

* A is the central atom; X is another atom

Draw the Lewis Formula for SiH4

Practice

H

H

H

H

Tetrahedral

Shape #4 =Trigonal PyramidalLewis

FormulaBond Pairs

Lone

Pairs

Total Pair

s

General Formul

a

Electron Pair Arrangement

Stereochemical Formula

3 1 4AX3

Etetrahedra

lTrigonal

pyramidal

• This Lewis formula indicates that NH3(g) has three bonds and one lone pair on the central atom.

• VSPER theory suggests that the four groups of e-’s should repel each other to form a tetrahedral shape (bond angle = 109.5°)

• But the lone pair is very repulsive, thus pushes the atoms more to a 107.3° bond angle

• This gives the molecule a trigonal pyramidal orientation.

* A is the central atom; X is another atom, E is a lone pair of electrons

Draw the Lewis Formula for PCl3

Practice

Cl

Cl

Cl

Trigonal pyramidal

Shape #5 =Angular (Bent)Lewis

FormulaBond Pairs

Lone

Pairs

Total Pair

s

General Formul

a

Electron Pair Arrangement

Stereochemical Formula

2 2 4AX2E

2

tetrahedral Angular

(Bent)

• This Lewis formula indicates that H2O(l) has two bonds and two lone pairs on the central atom.

• VSPER theory suggests that the four groups of e-’s should repel each other to form a tetrahedral shape (bond angle = 109.5°)

• But the TWO lone pairs are very repulsive, thus pushes the atoms more to a 105° bond angle

• This gives the molecule an angular (bent) orientation.

* A is the central atom; X is another atom, E is a lone pair of electrons

Draw the Lewis Formula for OCl2

Practice

Angular (bent)

Shape #6 =Linear (Tetrahedral)Lewis

FormulaBond Pairs

Lone

Pairs

Total Pair

s

General Formul

a

Electron Pair Arrangement

Stereochemical Formula

1 3 4 AXE3

Linear(Tetrahedra

l)

• This Lewis formula indicates that H2O(l) has two bonds and two lone pairs on the central atom.

• VSPER theory suggests that the four groups of e-’s should repel each other to form a tetrahedral shape (bond angle = 109.5°)

• But since there are only two atoms with one covalent bond holding them together, by definition, the shape is linear, as is the shape of every other diatomic molecule.

* A is the central atom; X is another atom, E is a lone pair of electrons

FH

Draw the Lewis Formula for HCl

Practice

VSEPR theory describes, explains, and predicts the geometry of molecules by counting pairs of electrons that repel each other to minimize repulsion. The process for predicting the shape of a molecule is summarized below:

Step 1: Draw the Lewis formula for the molecule, including the electron pairs around the central atom.

Step 2: Count the total number of bonding pairs (bonded atoms) and lone pairs of electrons around the central atom.

Step 3: Refer to Table 7, and use the number of pairs of electrons to predict the shape of the molecule.

Summary

Pg. 95

Draw the Lewis and stereochemical formulas for a sulfate ion, SO4

2- and predict the shapeSee pg. 95

Draw the Lewis and stereochemical formulas for a chlorate ion, ClO3

- and predict the shapeSee pg. 96

On your own: Pg. 96 #3

Practice

It is important to remember that a double or triple bond is one bond, and to treat it as such, when predicting the VSEPR shapes of molecules.

Example: Predict the shape of C2H4(g)

Draw the Lewis formula for the molecule

Count the # of pairs of e-’s around the central carbon atoms. The carbon atoms have 3 bonds (2 single, 1 double) and no

lone pairs. This is the same as a trigonal planar configuration.

Practice: Predict the shape for C2H2(g).

Multiple Bonds in VSEPR Models

H

H

H

H

Answer: See pg. 97

1) Finish pg. 96 #1-3

2) Pg. 98 #6-7 (Multiple Bond Practice)For 7 c, d, e - If there is more than one central atom

involved, tell me the shape around each of the central atoms Example:

3) Pg. 104 #1, 2, 3 #2: If there is more than one central atom involved, tell

me the shape around each of the central atoms

Homework

trigonal planar—first two carbonstetrahedral—third carbon

Draw the Lewis Formula for PCl3

Practice

Molecular PolarityDipole Theory

Section 3.3 – Part BPg. 98 - 104

1) Determine the polarity of a molecule based on simple structural shapes and unequal charge distribution

2) Describe bonding as a continuum ranging from complete electron transfer to equal sharing of electrons.

PolarityChemists believe that molecules are made up of charged

particles (electrons and nuclei).

A polar molecule is one in which the negative (electron) charge is not distributed symmetrically among the atoms making up the molecule. Thus, it will have partial positive and negative charges on

opposite sides of the molecule.

A molecule with symmetrical electron distribution is a nonpolar molecule.

The existence of polar molecules can be demonstrated by running a stream of water past a charged object. Demo: See Figure 9

TESTING A LIQUID WITH A CHARGED OBJECT:

In a liquid, molecules are able to rotate freely.

Polar molecules in a liquid will rotate so that their positive sides are closer to a negatively charged material.

Near a positively charged material they become oriented in the opposite direction.

EMPIRICAL RULES FOR POLAR AND NONPOLAR MOLECULES

Type Description of molecule

Examples

Polar AB diatomic with different atoms

HCl(g), CO(g)

NxAy containing nitrogen and other atoms

NH3(g), NF3(g)

OxAy containing oxygen and other atoms

H2O(l), OCl2(g)

CxAyBz containing carbon and two other kinds of atoms

CHCl3(l), C2H5OH(l)

Nonpolar

Ax all elements Cl2(g), N2(g)

CxAy containing carbon and only one other kind of atom (except CO(g))

CO2(g), CH4(g)

When the water test was repeated with a large number of pure liquids, this provided the set of empirical rules above.

Pg. 99

PREDICTING AND EXPLAINING POLARITY

Linus Pauling explained polarity by creating the concept of electronegativity. Introduced in Section 3.1

Electronegativity increases as you go up or to the right on the periodic table.

Pauling explained the polarity of a covalent bond as the difference in electronegativity of the bonded atoms.

If the bonded atoms have the same electronegativity, they will attract any shared electrons equally and form a nonpolar covalent bond.

If the atoms have different electronegativities, they will form a polar covalent bond.

The greater the electronegativity difference, the more polar the bond will be.

For a very large electronegativity difference, the difference in attraction may transfer one or more electrons resulting in ionic bonding.

PREDICTING AND EXPLAINING POLARITY

Cl2(g)

We use the Greek symbol delta to show partial charges

PREDICTING AND EXPLAINING POLARITY Pauling liked to think of chemical bonds as being different

in degree rather than different in kind.

According to him, all chemical bonds involve a sharing of electrons, with ionic bonds and nonpolar covalent bonds being just the two extreme cases

The bonding in substances therefore ranges anywhere along a continuum from nonpolar covalent to polar covalent to ionic.

For polar covalent bonds, the greater the electronegativity difference of the atoms, the more polar the bond.

~EN difference:nonpolar (<

0.4)polar (0.4+)ionic (m +

nm)

PRACTICESee pg. 100 Sample Problem 3.4

Try on your own pg. 100 #9 (a-c), 10, 11(a only)

DOES BOND POLARITY = MOLECULAR POLARITY?? NO! Chemists have found that the existence of polar bonds

in a molecule does not necessarily mean that you have a polar molecule.

Example: Carbon dioxide is found to be a nonpolar molecule, although

each of the CO bonds is a polar bond. WHY??

According to VSEPR (two bonds, no lone pairs) = linear arrangement

We will start showing bond polarity as arrows pointing in the negative direction (where e-’s want to go) = bond dipole Points from lower to higher electronegativity

The arrows are vectors and when added together, the equal but opposite bond dipoles equal zero.

Non-polar molecules are ones where the bond dipoles balance each other; producing a molecular dipole (vector sum) of zero

δ– δ+ δ–

O = C = O 3.4 2.6 3.4

Prediction Molecular Polarity Step 1: Draw a Lewis formula for the molecule.

Step 2: Use the number of electron pairs and VSEPR rules to determine the shape around each central atom.

Step 3: Use electronegativities to determine the polarity of each bond.

Step 4: Add the bond dipole vectors to determine whether the final result is zero (nonpolar molecule) or nonzero (polar molecule).

Guided Practice #1Predict the polarity of the water molecule.

1) Draw the Lewis formula

2) VSEPR: Draw the stereochemical formula

3) Assign the EN of the atoms, assign δ– and δ+ to the bonds

4) Draw in the bond dipoles

O

H HAngular (bent)

• The bond dipoles (vectors) do not balance. • Instead, they add together to produce a

nonzero molecular dipole (shown in red).• This results in a polar molecule (explains

bending water)

Go to Learning Tip

pg. 102

Guided Practice #2Predict the polarity of the methane

molecule.

1) Draw the Lewis formula

2) VSEPR: Draw the stereochemical formula

3) Assign the EN of the atoms, assign δ– and δ+ to the bonds

4) Draw in the bond dipoles• Notice how all the bond dipoles point into the central atom.

• There are no positive or negative areas on the outer part of

the molecule. • A tetrahedral molecule is symmetrical in 3-D and

four equal tetrahedral bond diploes always sum to zero

Tetrahedral

PracticePredict the bond polarity of the

ammonia, NH3(g) molecule. Include your reasoning.

Answer: See pg. 102

FYI: DO YOU REMEMBER “LIKE DISSOLVES LIKE”?

Means:“Polar substances are soluble in polar

substances; Non-polar substances are soluble in non-polar substances”

Mixing non-polar and polar substances results in them forming layers, with the least dense one on top.This occurs because polar molecules attract

each other more strongly; thus they stay close together excluding nonpolar molecules

Two clear liquids formed layers in this tube: nonpolar hexane C6H14(l)on top, and polar water, H2O(l) below.

Nonpolar dark orange liquid bromine, Br2(l) was then added. The bromine dissolves much more readily in the nonpolar hexane.

WHY DO WE CARE ABOUT POLAR MOLECULES?

Cleaning!! Water (polar) is useless at removing oil (nonpolar) so detergents are artificially created molecules that overcome this problem

Detergents have long, nonpolar sections which are attracted to (dissolve in) a tiny oil droplet.

The polar end of each of these detergent molecules helps form a polar “layer” around the droplet, which attracts polar water molecules.

This allows them to pull the oil droplet away from a stained area of fabric and hold it suspended in the wash water.

Homework: Pg. 102-103 #13-16

Pg. 104 # 4, 5, 10

Intermolecular Forces

Section 3.4 Pg. 105-117

1) Explain intermolecular forces, London (dispersion) forces, dipole-dipole

attractions and hydrogen bonding

2) Relate properties of substances to the predicted intermolecular bonding in the

substance.

BACKGROUND• All chemical changes (reactions) are accompanied by energy

changes

▫ Energy is mostly heat, light, or electrical energy

▫ Energy can be released slowly (battery) or quickly (fireworks)

▫ Two types of energy changes are possible:

EXOTHERMIC – energy is released into the surroundings - the product’s bonds have less energy than the reactant’s bonds

ENDOTHERMIC – energy is absorbed from the surroundings - the product’s bonds have more energy than the reactant’s

bonds

▫ Bond Energy – the energy required to break a chemical bond or the energy released when a bond is formed

BACKGROUND• There are three types of forces in matter:

1) Intranuclear force (bond) – bonds within the nucleus between protons and neutrons (very strong)

2) Intramolecular force (bond) – bonds between atoms within the molecule or between ions within the crystal lattice (quite strong)

3) Intermolecular force (bond) – bonds between molecules (quite weak); are electrostatic (involve positive and negative charges)

There are 3 types of intermolecular bonds:

a) Dipole-Dipole Forces (a.k.a. Polar Forces)

b) London Force (a.k.a. London Dispersion Force, Dispersion Force)

c) Hydrogen BondingNote: “Van der Walls force” – includes London and dipole-dipole forces

Weakest

Medium

Strongest

1) Dipole-Dipole Force• The simultaneous attraction between

oppositely charged ends of polar molecules.

▫ Simply put, the attraction between diploes

▫ Dipole-dipole forces are among the weakest intermolecular forces, but still control important properties (i.e. Solubility because water is polar))

Dipole: a partial separation of positive and negative charges within a molecule, due to electronegativity differences

1) Dipole-Dipole ForceIn a liquid, polar molecules can move and rotate to maximize attractions and minimize repulsions. The net effect is greater overall attraction.

The strength of the dipole-dipole force is dependent on the overall polarity of the molecule

Note: If a molecule is polar it will be soluble in water? Why?

1) Dipole-Dipole Forces In a liquid: In a solid:

2) London ForceSimultaneous attraction between a momentary dipole

in a molecule and the momentary dipoles in surrounding molecules

momentary dipole: an uneven distribution of electrons around a molecule, resulting in a temporary charge difference between its ends

They last for just the instant

that the electrons are

not distributed perfectly even.

2) London Force• Fritz London also showed that momentary dipoles

occurring in adjacent molecules would result in an overall attraction

• The strength of the London force is directly related to the number of electrons in the molecule, and inversely related to the distance between the molecules.

▫ Increase electrons = Increase force (directly related)..▫ Increase distance = Decrease force (inversely related)

2) London Force• The key point is that:

▫ the more electrons a molecule has, the more easily momentary dipoles will form, and the greater the effect of the London force will be.

• London forces are present between all molecules, whether any other type of attraction is present.

Why do we care about intermolecular forces?We can use Dipole-Dipole and London Forces to predict Boiling

PointsCompound (at SATP) Electrons Boiling Point (°C)

CH4(g) 10 -164

SiH4(g) 18 -112

GeH4(g) 36 -89

SnH4(g) 54 -52

Remember (if all other factors are equal):

1) The more polar the molecule = The stronger the dipole-dipole force

2) Increase the number of electrons = Increase the strength of London Force

A higher boiling point temperature means more energy has to be added, thus we assume the intermolecular

forces are stronger. (see Learning Tip pg. 109)

Example #1• Use Intermolecular force theory to predict which of the

following hydrocarbons has the highest boiling point:

▫ methane (CH4), ethane (C2H6), propane (C3H8), butane (C4H10)

1) Are the molecules polar or non-polar? non-polar (no dipole-dipole force)

2) Which has more electrons? butane: greatest # of e-’s = greatest London force

Check:Alkane Boiling Point (°C)

methane -162

ethane -89

propane -42

butane -0.5

Example #2• Use Intermolecular force theory to predict which of the following

has the highest boiling point:

▫ bromine (Br2 ) or iodine monochloride (ICl)

1) Which has more electrons?

They are isoelectronic: have the same number of electrons (70 e-’s)

-Therefore the London force is the same (or nearly the same)

2) Are the molecules polar or non-polar?

-Bromine is non-polar (has no dipole-dipole force; only London forces)

- Iodine monochloride is polar (has dipole-dipole forces and London forces)

- This extra attraction between ICl molecules produces a higher boiling point

Check: Substance Electrons Boiling Point (°C)

bromine 70 59

iodine monochloride

70 97

You cannot predict boiling points when:One molecule has a stronger dipole-dipole force and

the other has a stronger London force

The two molecules differ significantly in shape

The central atom of either molecule has an incomplete octet

PracticePg. 109 #1-4

4) a) Boron is stable with 6 valence e-

b) chloromethane (CH3Cl)

3) Hydrogen Bonding• Occurs when a hydrogen atom bonded to a

strongly electronegative atom, (N, O and F) is attracted to a lone pair of electrons in an adjacent molecule.▫ Hydrogen nucleus (proton) is simultaneously

attracted to two pairs of electrons; one closer (in the same molecule) and one further away (on the next molecule)

Why do you need a strongly

electronegative atom?

It pulls the hydrogen’s electron away making it

“unshielded”, so the lone pair on the other side can come much

closer

•• •

•

••

••

3) Hydrogen Bonding• Hydrogen bonds are momentary attractive

forces between passing mobile molecules but are the strongest of the intermolecular forces.

• Hydrogen bonds only act as continuous bonds between molecules in solids, where the molecules are moving slowly enough to be locked into position.

• Hydrogen force would have been a better name.

3) Hydrogen BondingIn ice, hydrogen

bonds between the molecules result in a regular hexagonal crystal structure.

The ···H– represents a hydrogen nucleus (proton) being shared unequally between two pairs of electrons

3) Hydrogen Bonding• Do lakes freeze from the

bottom-up or the top-down?

• Top–down, because water is unique in that its solid form (ice) is less dense than its liquid form. Why??

• The hydrogen bonds hold water molecules in a hexagonal lattice with open space in the center, which explains the low density (mass/volume) of ice.

Hydrogen Bonding in DNA• FYI: The double helix of the

DNA molecule owes its unique structure largely to hydrogen bonding.

• The red bonds are hydrogen bonds.

• If the helix were held together by covalent bonds, the DNA molecule would not be able to unravel and replicate and life could not continue!!

• Explains surface tension, shape of a meniscus, volatility and capillary action

1) Surface Tension

▫ Molecules within a liquid are attracted by other molecules in all directions equally, but right at the surface, molecules are only attracted downwards and sideways. This means the net pull is downward so the surface tends to stay intact

▫ The stronger the intermolecular force the stronger the surface tension.

Why do we care about intermolecular forces?This shows water adhering to the faucet gaining mass until it is stretched to a point where the surface tension can no longer bind it to the faucet.

It then separates and surface tension forms the drop into a sphere.

2) Capillary Action – due to adhesion (attraction between unlike molecules) and cohesion (attraction of like molecules)

▫ The adhesion between water and glass is greater than the cohesion between water molecules.

▫ The cohesion between mercury molecules is greater than the adhesion between mercury and glass

Why do we care about intermolecular forces?

In a sense, water is pulled up the tube by

the intermolecular forces between

water and glass

Hg clip

Meniscus

PracticePg. 117 # 1, 4, 5

#1 – use pg. 99 table to determine polarity#1 – look for NH2, NH, OH2, OH, to determine if hydrogen

bonding is possible

Ex. CH3CHOHCH3 will it have hydrogen bonding?