UNIT 3 CHAPTERS 6 &10 Chemical Bonding & Intermolecular Forces

UNIT 3CHAPTERS 6 &10

Chemical Bonding & Intermolecular Forces

Chemical Bonding

Chapter 6

Chemical Bonding

• Valence electrons are the electrons in the outer shell (highest energy level) of an atom.

• A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

• During bonding, valence electrons are redistributed in ways that make the atoms more stable.

Chapter 6 – Section 1: Introduction to Chemical Bonding

The Three Major Types of Chemical Bonding

• Ionic Bonding results from the electrical attraction between oppositely-charged ions.

• Covalent Bonding results from the sharing of electron pairs between two atoms.

• Metallic Bonding results from the attraction between metal atoms and the surrounding sea of electrons.


Ionic or Covalent?

• Bonding is usually somewhere between ionic and covalent, depending on the electronegativity difference between the two atoms.

• In polar covalent bonds, the bonded atoms have an unequal attraction for the shared electron.


0.30 1.7 3.3

Bonding Morebetween Electroneg. negativesulfur and: difference Bond typeatom

Ionic or Covalent?Sample Problem

Use electronegativity values (in table on pg 161)to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?Solution:


hydrogencesiumchlorine

2.5 – 2.1 = 0.42.5 – 0.7 = 1.83.0 – 2.5 = 0.5

polar-covalentpolar-covalent

ionicsulfursulfurchlorine

Molecules

• A covalent bond is formed from shared pairs of electrons.

• A molecule is a neutral group of atoms held together by covalent bonds.

Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds

Visual Concept

Why Do Covalent Bonds Form?• When two atoms form a covalent

bond, their shared electrons form overlapping orbitals.

• This gives both atoms a stable noble-gas configuration.


The Octet Rule

• Atoms are the most stable whenthey have completely full valenceshells (like the noble gases.)

• The Octet Rule – Compounds tend to form so that each atom has an octet (group of eight) electrons in its highest energy level.

• Hydrogen is an exception to the octet rule since it can only have two electrons in its valence shell.


Visual Concept

Electron-Dot Notation

• Electron-dot notation is indicated by dots placed around the element’s symbol. Only the valence electrons are shown.Inner-shell electrons are not shown.


Visual Concept

Electron-Dot NotationSample Problem

a. Write the electron-dot notation for hydrogen.b. Write the electron-dot notation for nitrogen.Solution:a.Hydrogen is in group 1. It has one valence electron.

a.Nitrogen is in group 15. It has 5 valence electrons.


H

N

•

•••

••

Lewis Structures

• Electron-dot notations of two or more atoms can be combined to represent molecules.

• Unpaired electrons will pair up to form a shared pair or covalent bond.


Lewis Structures (continued)

• The pair of dots representing the shared pair of electrons in a covalent bond is often replaced by a long dash.

• An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.


Shared pair (covalent bond)

Lone pair

••

How to Draw Lewis Structures1. Draw the electron-dot notation for each type of

atom, and count the valence electrons.2. Put the least electronegative atom in the

center (except H.)3. Use electron pairs to form bonds between all

atoms.4. Make sure all atoms (except H) have octets.5. Count the total electrons in your Lewis

structure. Does it match the number you counted in step 1? If not, introduce multiple bonds.


Lewis StructuresSample Problem A

Draw the Lewis structure of iodomethane, CH3I.

Solution:Step 1 - Draw the electron-dot

notation for each type of atom, and count the valence electrons.


H•C••

•• I •

•

•• ••

•

C 1 x 4 e- = 4 e- 3H 3 x 1 e- = 3 e- I 1 x 7 e- = 7 e- 14 e-

Total

Lewis StructuresSample Problem A (continued)

Step 2 – Put the least electronegative atom in the center (except H).Step 3 – Use electron pairs to form bonds between all atoms.Step 4 – Make sure all atoms (except H) have octets.Step 5 – Count the total electrons. Does it match your beginning total?


H C I••

••

H

H

14 Total e-

••

Multiple Covalent Bonds

• In a single covalent bond, one pair ofelectrons is shared between two atoms.

• A double bond is a covalent bond in which two pairs of electrons are shared between two atoms.

• A triple bond is a covalent bond in which three pairs of electrons are shared between two atoms.

• Multiple bonds are often found in molecules containing carbon, nitrogen,

and oxygen.


Single Bond

Double Bond

Triple Bond

Lewis StructuresSample Problem B

Draw the Lewis structure for methanal, CH2O.

Solution:Step 1 - Draw the electron-dot

notation for each type of atom, and count the valence electrons.


H•C••

•• O•

•

•• ••

C 1 x 4 e- = 4 e- 2H 2 x 1 e- = 2 e- O 1 x 6 e- = 6 e- 12 e-

Total

Lewis StructuresSample Problem B (continued)

Step 2 – Put the least electronegative atom in the center (except H).Step 3 – Use electron pairs to form bonds between all atoms.Step 4 – Make sure all atoms (except H) have octets.Step 5 – Count the total electrons. Does it match your beginning total?If not, introduce multiple bonds (remove 2 lone pairs to make 1 shared pair.)Now does it match?


C O••

••

H

H

14 Total e-

••

••

12 Total e-

Formation of Ionic Compounds• Sodium and other metals easily lose

electrons to form positively-charged ions called cations.

• Chlorine and other non-metals easily gain electrons to form negatively-charged ions called anions.

Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds

Ionic Bonding

• Cations (+) and anions (-)are attracted to each other because of their opposite electrical charges.

• An ionic bond is a bondthat forms betweenoppositely-charged ionsbecause of their mutualelectrical attraction.


Visual Concept

Ionic Bonding and the Crystal Lattice• In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice.

• A formula unit is the smallest repeating unit of an ionic compound.


Sodium Chloride crystal lattice (many Na and Cl atoms)

Formula Unit = NaCl

Comparing Ionic and Covalent Compounds

• Covalent compounds have relatively weak forces of attraction between molecules, but ionic compounds have a strong attraction between ions. This causes some differences in their properties:


Ionic Covalentcrystals molecule

svery high melting points

low melting pointshard, but brittle usually gas or

liquidEx: NaCl, CaF2, KNO3

Ex: H2O, CO2, O2

Polyatomic Ions

• A charged group of covalently bonded atoms is known as a polyatomic ion.

• Draw a Lewis structure for a polyatomic ion with brackets around it and the charge in the upper right corner.


hydroxide ion, OH-

ammonium ion, NH4

+

The Metallic Bond

• In metals, overlapping orbitals allow the outer electrons of the atoms to roam freely throughout the entire metal.

• These mobile electrons form a sea of electrons around the metal atoms, which are packed together in a crystal lattice.

• A metallic bond results from the attraction between metal atoms and the surrounding sea of electrons.

Chapter 6 – Section 4: Metallic Bonding

Properties of Metals

• The characteristics of metallic bonding gives metals their unique properties, listed below.

electrical conductivity thermal (heat) conductivity malleability (can be hammered into thin sheets)

ductility (can be pulled or extruded into wires)

luster (shiny appearance)

Chapter 6 – Section 4: Metallic Bonding

Visual Concept

VSEPR Theory

• The abbreviation VSEPR (say it “VES-pur”) stands for “valence-shell electron-pair repulsion.”

• VSEPR theory – repulsion between pairs of valence electrons aroundan atom causes the electron pairs tobe oriented as far apart as possible.

• Treat double and triple bonds the same as single bonds.

Chapter 6 – Section 5: Molecular Geometry

Visual Concept

VSEPR Theory (continued)

• VSEPR theory can also account for the geometries of molecules with unshared electron pairs.

• VSEPR theory postulates that the lone pairs occupy space around the central atom just like bonding pairs, but they repel other electron pairs more strongly than bonding pairs do.



• 2 electron pairs around acentral atom will be 180o

apart, and the molecule’sshape will be linear.

• 3 bonding pairs around acentral atom will be 120o

apart, and the molecule’sshape will be trigonal planar.If one of the pairs is a lonepair, the shape will be bent.



• 4 bonding pairs around acentral atom will be 109.5o

apart, and the molecule’sshape will be tetrahedral.If one of the pairs is a lonepair, the shape will be trigonal pyramidal. If two of the pairs are lone pairs, the shape will be bent.

• Unshared pairs repel electrons more strongly and will result in smaller bond angles.


VSEPR TheorySample Problem A

Use VSEPR theory to predict the molecular geometry of water, H2O.Solution:Draw the Lewis Structure for H2O:How many total electron pairs aresurrounding the central atom?

How many are unshared pairs?

The shape is bent.


Total Electrons: 8 e-

O••

••HH

Octets

4

2 OHH

••••

VSEPR TheorySample Problem B

Use VSEPR theory to predict the molecular geometry of carbon dioxide, CO2.

Solution:Draw the Lewis Structure for CO2:

How many total electron pairs aresurrounding the central atom?

The shape is linear.


Total Electrons: 16 e-

C••

••OO

Octets

2 (double or triple bonds count the same as single)

••

••

••

••••

••

Molecular Polarity

• Molecular Polarity depends on both bondpolarity and molecular geometry. If all bonds are

non-polar, the moleculeis always non-polar.

If bonds are polar, but there is symmetry in the molecule so that the polarity of the bonds cancels out, then the molecule is non-polar. (Ex: CO2, CCl4)

If bonds are polar but there is no symmetry such that they cancel each other out, the overall molecule is polar. (Ex: H20, CH3Cl)


Intermolecular Forces

• The forces of attraction between molecules are called intermolecular forces.

• Intermolecular forces vary in strength but are

generally weaker than any of the three typesof chemical bonds (covalent, ionic or

metallic.)


Intermolecular Forces (continued)

• The strongest intermolecular forces exist between polar molecules.

• Because of their uneven charge distribution, polar molecules have dipoles.

• A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole.


OHH

••••

Visual Concept

Types of Intermolecular Forces• 3 types of intermolecular forces (strongest to

weakest):1. Dipole-dipole – between 2 polar

molecules. The - side of 1 dipole attracts the + side of another.• Hydrogen Bonding – a very strong

type of dipole-dipole force. Only existsbetween atoms of H and N, O or F.

2. Induced dipole – between a polar and a non-polar molecule.

3. London dispersion forces – instantaneous dipoles created by the constant motion of electrons.


Visual Concept

States of Matter

Chapter 10

The Kinetic-Molecular Theory• The kinetic-molecular theory of matter

states: Particles of matter (atoms and

molecules) are always in motion. We measure this energy of motion

(kinetic energy) as temperature. If temperature increases, the

particles will gain more energy and move even faster.

Molecular motion is greatest in gases, less in liquids, and least in solids.

Chapter 10 – Section 1: The Kinetic-Molecular Theory of Matter

Gases

• An Ideal Gas is a hypothetical gas that perfectly fits all the assumptions of the kinetic-molecular theory.• Many gases behave nearly

ideally if pressure is not veryhigh and temperature is not very low.• Fluidity – Gas particles glide easily

past one another. Because liquids and gases flow, they are both referred to as fluids.


Gases (continued)

• Low Density – Gas particles are very far apart. The density of a gas is about 1/1000 the density of the same substance in the liquid or solid state.• Expansion – A gas will expand to fill its

container.• Compressibility – The volume of a gas

can be greatly decreased by pushing the particles closer together.


Liquids

• Surface Tension – Strong cohesive forces at a liquid’s surface act to decrease the surface area to the smallest possible size. The higher the force of attraction between the particles of a liquid, the higher the surface tension.

Chapter 10 – Section 2: Liquids

Liquids (continued)

• Vaporization – A liquid or solidchanging to a gas. Evaporation – particles escape from the surface of a liquid andbecome a gas. This occurs because liquid particles havedifferent kinetic energies.

Boiling – bubbles of vapor appear throughout a liquid. Will not occur below a certain temperature (the boiling point.)

• A volatile liquid is one that evaporates readily.

Chapter 10 – Section 2: Liquids

Solids

• There are two main types of solids: Crystalline Solids – Made up

of crystals. Particles are arranged in an orderly, geometric, repeating pattern.

Amorphous Solid – Particles are arranged randomly.

Chapter 10 – Section 3: Solids

Solids (continued)

• Melting Point – The temperature at which a solid becomes a liquid. At this temperature, the kinetic energies of the particles within the solid overcome the attractive forces holding them together.

Chapter 10 – Section 3: Solids

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UNIT 3 CHAPTERS 6 &10 Chemical Bonding & Intermolecular Forces