week 3 intermolecular forces&lewis structure

Prepared by:Mrs Faraziehan Senusi

PA-A11-7C

David P. White Prentice Hall ©

2003

Quantum Theory

Atomic Orbitals

Electronic Configuration

Chapter 1Atoms, Molecules & Chemical bonding

Molecular Orbitals

Bonding and Intermolecular Compounds

Introduction

• Physical properties of substances understood in terms of kinetic molecular theory:– Gases are highly compressible, assumes shape and volume

of container: • Gas molecules are far apart and do not interact much with

each other.– Liquids are almost incompressible, assume the shape but

not the volume of container:• Liquids molecules are held closer together than gas molecules,

but not so rigidly that the molecules cannot slide past each other.

– Solids are incompressible and have a definite shape and volume:• Solid molecules are packed closely together. The molecules are

so rigidly packed that they cannot easily slide past each other.

A Molecular Comparison of Liquids and Solids

• Converting a gas into a liquid or solid requires the molecules to get closer to each other:– cool or compress.

• Converting a solid into a liquid or gas requires the molecules to move further apart: – heat or reduce pressure.

• The forces holding solids and liquids together are called intermolecular forces.


• The covalent bond holding a molecule together is an intramolecular forces.

• The attraction between molecules is an intermolecular force.

• Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).

• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

Intermolecular Forces


Ion-Dipole Forces• Interaction between an ion and a dipole (e.g. water).• Strongest of all intermolecular forces.• ion-dipole attractions become stronger as either the charge on

the ion increases, or as the magnitude of the dipole of the polar molecule increase.


Dipole-Dipole Forces• Dipole-dipole forces exist

between neutral polar molecules.

• Polar molecules need to be close together.

• Weaker than ion-dipole forces.• There is a mix of attractive and

repulsive dipole-dipole forces as the molecules tumble.

• If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity.


Dipole-Dipole Forces


London Dispersion Forces• Weakest of all intermolecular forces.• It is possible for two adjacent neutral molecules to

affect each other.• The nucleus of one molecule (or atom) attracts the

electrons of the adjacent molecule (or atom).• For an instant, the electron clouds become

distorted.• In that instant a dipole is formed (called an

instantaneous dipole).



London Dispersion Forces• One instantaneous dipole can induce another

instantaneous dipole in an adjacent molecule (or atom).

• The forces between instantaneous dipoles are called London dispersion forces.


London Dispersion Forces• Polarizability is the ease with which an electron cloud

can be deformed.• The larger the molecule (the greater the number of

electrons) the more polarizable.• London dispersion forces increase as molecular weight

increases.• London dispersion forces exist between all molecules.• London dispersion forces depend on the shape of the

molecule.• The greater the surface area available for contact, the

greater the dispersion forces.


London Dispersion Forces


Hydrogen Bonding• Special case of dipole-dipole forces.• By experiments: boiling points of compounds with

H-F, H-O, and H-N bonds are abnormally high.• Intermolecular forces are abnormally strong.


Hydrogen Bonding

Hydrogen Bonding• H-bonding requires H bonded to an

electronegative element (most important for compounds of F, O, and N).– Electrons in the H-X (X = electronegative element) lie

much closer to X than H.– H has only one electron, so in the H-X bond, the + H

presents an almost bare proton to the - X.– Therefore, H-bonds are strong.


Hydrogen Bonding

Hydrogen Bonding• Hydrogen bonds are responsible for:

– Ice Floating• Solids are usually more closely packed than liquids;• Therefore, solids are more dense than liquids.• Ice is ordered with an open structure to optimize H-bonding.• Therefore, ice is less dense than water.• In water the H-O bond length is 1.0 Å.• The O…H hydrogen bond length is 1.8 Å.• Ice has waters arranged in an open, regular hexagon.• Each + H points towards a lone pair on O.


Hydrogen Bonding



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Lewis Structures, Octet Rule, Resonance & Formal Charge

The number and arrangements of electrons in the outermost shells of atoms determine the chemical and physical properties of the elements as well as the kinds of chemical bonds they form.

We write Lewis dot formulas (or Lewis dot representations, or just Lewis formulas) as a convenient bookkeeping method for keeping track of these “chemically important electrons.”

“In most compounds, the representative elements in the compounds achieve noble gas configurations”

• This statement is usually called the Octet Rule, because the noble gas configurations have 8 e in their outermost shells (except for He, which has 2 e).

• Many Lewis formulas are based on this idea

The Octet Rule

How to write Lewis structures for molecules & polyatomic ions?

• Step 1 : Count the total number of valence electrons in the molecule

Doesn’t matter which element it comes from, just the total of all elements.

• Step 2 : Set up a skeleton structure least electronegative element, or carbon in center. H is always at the outskirts. Most of the time (but not always) halogen will also be at the

outskirts.

• Step 3 : Place the valence electrons until every atom has a stable octet by:

i) drawing bond (either ‘’ , ‘’ or ‘’) and/or

ii) drawing lone pair (non-bonding pair of electrons) (:)

**Make sure every atom has a stable octet, which means, every atom needs to have 8 valence electrons around it. Except for H of course, which satisfy its stability with 2 electrons.

**Don’t forget to show all lone pair electrons.

Now, let’s practise writing Lewis structures!!!

Draw Lewis structures for:

1) CCl2F2

2) NF3

3) N2

4) CO32-

Exceptions to the Octet Rule

1) Electron deficient structures Gaseous molecules containing either beryllium or boron as central atom ~ they have

fewer than eight electrons around Be or B atom. Lewis structure for BeCl2 and BF3

Multiple bonds to the central atoms give unlikely structures:

Halogens are much more electronegative than Be or B.

2) Odd – electrons molecules A few molecules contain a central atom with odd number of valence electrons

~ free radicals, contain a lone (unpaired) electron. Resonance for nitrogen oxide, NO2

But, the form with the lone electron on N (left) must be impotant because of the way NO2 react.

When two NO2 molecules collide, the lone electron pair up to form N–N bond in dinitrogen tetraoxide (N2O4) and each N attains an octet.


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3) Expanded valence shells/octet Molecules and ions have more than eight valence electrons

around the central atom. Expanded valence shells occur only with a central nonmetal atom

in which d orbitals are available, one from period 3 or higher. Example: Sulfur hexafluoride, SF6. The central sulfur is

surrounded by six single bonds, one to each fluorine for a total of 12 electrons.


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• A molecule or polyatomic ion for which two or more Lewis formulas with the same arrangements of atoms can be drawn to describe the bonding is said to exhibit resonance.

• The three structures below are resonance structures of the carbonate ion:

• The relationship among them is indicated by the double-headed arrows.The double-headed arrow indicates that the structures shown are resonance structures.

Resonance structures

• Draw resonance structures for CO32- and NO3

– .


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Now, let’s practise!!!

• Formal charge is the hypothetical charge on an atom in a molecule or polyatomic ion.

• The concept of formal charges helps us to write correct Lewis formulas in most cases.

• The most energetically favorable formula for a molecule is usually one in which the formal charge on each atom is zero or as near zero as possible.

Formal Charge


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Formal Charge

e) valenceshared ofnumber (2

1e valenceunshared ofnumber totalatom free in e valenceofnumber

or

e) bonding ofnumber total(2

1e nonbonding ofnumber totalatom free in e valenceofnumber total

atom of charge Formal


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• Write formal charges for the carbonate ion, CO32-.

• Write formal charges for the ammonia molecule, NH3 and to the ammonium ion, NH4

+.


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Now, let’s practise!!!

week 3 intermolecular forces&lewis structure

Technology

intramolecular forces

repulsive dipoledipole

liquids molecules

gas molecules

solid molecules

iondipole attractions

molecules tumble

intermolecular compounds