Polarity, IMFs, Solids

Chapter 8 & 9

Chapter 11- section 1 &2Chapter 12 – section 1

Lewis Structures• A Lewis structure is a representation of a

molecule that shows BOTH the bonding between atoms AND how the valence electrons are arranged among the atoms in the molecule.

• Octet Rule - all atoms (except H, Be, B) try to get 8 valence electrons around themselves.– H wants 2 electrons– Be wants 4 electrons– B wants 6 electrons

RULES FOR MAKING LEWIS STRUCTURES

1. Obtain the sum total of the VALENCE electrons of ALL the atoms in the molecule.

2. For ion structures, add or subtract electrons to the total as indicated by the charge on the ion. Examples: -2 ion, add 2 electrons; +1 ion, subtract 1 electron

3. Determine the central atom– A single atom is often the central atom– The atom that can form the most bonds is often the central

atom

4. Start by making single bonds – In most cases, there is a symmetrical

arrangement of the atoms.

5. Arrange the remaining electrons to satisfy the octet rule. These electrons are called "lone pairs".

6. If you do not have enough electrons to satisfy all atoms, then double or possibly triple bonds will be required.

7. If you have too many electrons the extras make an “expanded octet” on the central atom.

Valence Shell Electron Pair Repulsion Theory

(VSEPR)• Used to determine the SHAPE of a

molecule - shape helps determine the properties

• ALL electron pairs (bonding and lone pairs) try to get as far apart as possible - all electrons repel each other

Electronegativity• The measure of an atom’s attraction for a pair of

BONDING electrons

- low electronegativity is characteristic of metals- high electronegativity is characteristic of non-

metals- Electronegativity generally decreases going

down a group and increases across a series.- Fluorine has the highest electronegativity

Note: • You will not necessarily have electronegativity values on the

test. – Know the general trends on the periodic table– Know that:

• If there is a metal, it’s ionic• Any bond with fluorine is polar• Common polar bonds: H-N, H-O, H-F, H-Cl, C-N, C-Cl, S-O and C-O• Common nonpolar bonds: C-H, anything bonded to itself

Electronegativity Difference

The bond is: Example (difference)

0-0.4 Nonpolar Cl-Cl (0.0)

0.4-1.0 Polar H-Cl (0.9)

1.0-2.0 Very polar H-F (1.9)

>2.0 Ionic NaCl (2.1)

Molecular Polarity (p.343)Nonpolar – Molecule contains only nonpolar

bonds OR contains polar bonds that are arranged symmetrically.

Polar – Molecule must contain at least 1 polar bond and the bonds are not arranged symmetrically. This causes a charge distribution within the molecule.

*You must know the shape of the molecule!

Predicting Molecular PolarityWhen there are no polar bonds in a molecule, there is no permanent charge difference between one part of the molecule and another, and the molecule is nonpolar. For example, the Cl2 molecule has no polar bonds because the electronegativity is identical for both atoms.It is therefore a nonpolar molecule.

Cl - Cl

Predicting Molecular PolarityA molecule can possess polar bonds and still be nonpolar. If the polar bonds are evenly (or symmetrically) distributed, the bond dipoles cancel and do not create a molecular dipole. For example, the three bonds in a molecule of BF3 are significantly polar, but they are symmetrically arranged around the central boron atom. No side of the molecule has more negative or positive charge than another side, and so the molecule is nonpolar:

Predicting Molecular PolarityA water molecule is polar because (1) its O-H bonds are significantly polar, and (2) its bent geometry makes the distribution of those polar bonds asymmetrical. The side of the water molecule containing the more electronegative oxygen atom is partially negative, and the side of the molecule containing the less electronegative hydrogen atoms is partially positive.

+ side of molecule

- side of molecule

EXAMPLES – Predicting Molecular Polarity:  Decide whether the molecules represented by the following formulas are polar or nonpolar.a. CO2     b. OF2     c. CCl4     d. CH2Cl2     e. HCN

a. The C-O bonds are polar.The Lewis structure and shape for CO2 is

If we put arrows into the sketch for CO2,

This is a symmetrical arrangement so this molecule is NONPOLAR

this side negativethis side negative

EXAMPLES – Predicting Molecular Polarity:  Decide whether the molecules represented by the following formulas are polar or nonpolar.a. CO2     b. OF2     c. CCl4     d. CH2Cl2     e. HCN

b. The O-F bonds are polar .The Lewis structure for OF2 is

The molecular geometry of OF2 is bent and adding arrows gives:

This asymmetrical distribution of polar bonds would produce a POLAR molecule.

this side +

this side -

EXAMPLES – Predicting Molecular Polarity:  Decide whether the molecules represented by the following formulas are polar or nonpolar.a. CO2     b. OF2     c. CCl4     d. CH2Cl2     e. HCN

c. The C-Cl bonds are polar. The Lewis structure for CCl4 is

The molecular geometry is tetrahedral and adding arrows gives:                

This is a symmetrical arrangement and produces aNONPOLAR molecule

this side -

this side -

EXAMPLES – Predicting Molecular Polarity:  Decide whether the molecules represented by the following formulas are polar or nonpolar.a. CO2     b. OF2     c. CCl4     d. CH2Cl2     e. HCN

d.  The C-Cl bonds are polar and the C-H bonds are nonpolar.

The Lewis structure for CH2Cl2 is

The molecular geometry is tetrahedral and adding arrows gives:

This is an asymmetrical arrangement that produces a POLAR molecule.

this side +

this side -

EXAMPLES – Predicting Molecular Polarity:  Decide whether the molecules represented by the following formulas are polar or nonpolar.a. CO2     b. OF2     c. CCl4     d. CH2Cl2     e. HCN

e.  The H – C bonds are nonpolar and the C – N bond is polar.

The Lewis structure is:

The molecular geometry is linear and adding arrows gives:

This is an asymmetrical arrangement and the molecule is POLAR.

this side +

this side -

• These are forces that exist BETWEEN particles of a substance (for example: the forces between 2 hydrogen molecules - NOT the bond holding the 2 hydrogen atoms together to make the molecule)

Intermolecular Forces (IMF)

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Intermolecular Forces

Weakest IMFsStrongest IMFs

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Dispersion ForcesAn instantaneous dipole can form when electron densitybecomes unbalanced because electrons move as seen

in the helium atom below

Frame 1 Frame 2 Frame 3

When both electrons end up on one side of the atom (like in frame 3), that side of the

atom acquires a δ- charge and the other side becomes δ+

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Dispersion Forces

• These forces are present between particles of ALL substances

IN GENERAL:• Their relative strength is related to the total number of

protons & electrons in a single particle (a larger molar mass indicates more protons & electrons - stronger dispersion forces)

• London dispersion forces increase as surface area increases. Increasing surface area makes the molecule more polarizable (able to form instantaneous dipole easier)

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Higher molar mass means larger dispersion forces and higher mp and bp

Higher molar mass because more protons, neutrons and electrons – more charged particles – more attraction.

We can use the molar mass to determine the relative strength of dispersion forces BUT the REASON is the increase in the number of charged particles NOT the increase in mass

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same formula with same molar mass BUT branching decreases the dispersion forces (decreases the polarizability of the molecule’s electron cloud) so branched structure has lower boiling point (bp)

Dipole Forces Only found between molecules that are polar

• Polar molecules form dipoles (regions of partial charges formed by the unequal sharing of the electrons in the polar bond)

Dipole forces are stronger than dispersion forces

Relative strength is related to the polarity of the molecule

Hydrogen Bonding• Hydrogen bonding is a

special form of dipole forces.

• It is the strongest dipole force

• It only occurs when molecules have a hydrogen attached to a VERY electronegative element (nitrogen, oxygen or fluorine)

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Examples of H-bondingNotice that H “bonds” to the lone pair on the O or N

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To maximize H-bonding, water molecules actually get farther apart when changing from the liquid state to the solid state. This makes iceless dense than water.

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Showing the strength of hydrogen bonding using BP data

High BP due to presence of

H bonding

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Ion-Dipole Force• An ion-dipole force is an interaction between an ion (e.g.,

Na+) and the partial charge on the end of a polar molecule (e.g., water).

• Not really an intermolecular force because there is more than 1 substance involved

• important for aqueous solutions of ionic compounds • the strength of the ion-dipole attraction is one of the main

factors that determines the solubility of ionic compounds in water

• Example: NaCl (aq)

Ionic Bonds • Are only found between positive and

negative ions in an ionic compound

IMF Flow Chart

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Comparing IMFs• Dispersion forces are found in all substances.

– Their strength depends on molecular shapes and the number of charged particles (molar mass).

• Dipole-dipole forces add to the effect of dispersion forces. – They are found only in polar molecules.– H-bonding is a stronger special case of dipole-dipole interactions.

H-bonding is only significant when H is bonded directly to N, O or F.

• Ion-dipole and ionic bonding are possible when there are ions.– Ion-dipole interactions are slightly stronger than H-bonds. – Ionic bonding attractions are stronger than ion-dipole forces.

IN GENERAL Ionic bonding > Ion-dipole > H-bonding > dipole-dipole > dispersion

SolidsTypes of solids:• Ionic

• All ionic compounds are solids• Ions held by the attraction between cations and

anions (ionic bond)• Generally brittle, not electrical conductors, held in

crystal structures

• Metallic Solids– Held together by a “sea” of electrons– Electrical conductors– Alloys (p.473)

• Mixtures of metals• Bronze = copper + tin• Metals are mixed to form materials with desirable

properties.

• Covalent-Network Solids– Held together by covalent bonds– Often very hard (diamond)

• Polymers– long chains of atoms (often carbon)

• Molecular Solids– Held together by strong IMF– Usually softer than other types of solids