1. Chemical Bonding & Intermolecular Forces
UNIT 3 Chapters 6 &10
Chemical Bonding & Intermolecular Forces

2. Chapter 6
Chemical Bonding

3. Chapter 6 – Section 1: Introduction to 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.

4. The Three Major Types of Chemical Bonding
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.

5. Chapter 6 – Section 1: Introduction to Chemical Bonding
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.
0.3
1.7
3.3

6. Ionic or Covalent? Sample Problem
Chapter 6 – Section 1: Introduction to Chemical Bonding
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:
Bonding More between Electroneg negative
sulfur and: difference Bond type atom
hydrogen
cesium
chlorine
2.5 – 2.1 = 0.4
polar-covalent
sulfur
2.5 – 0.7 = 1.8
ionic
sulfur
3.0 – 2.5 = 0.5
polar-covalent
chlorine

7. Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Molecules
A covalent bond is formed from shared pairs of electrons.
A molecule is a neutral group of atoms held together by covalent bonds.
Visual Concept

8. Why Do Covalent Bonds Form?
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
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.

9. Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
The Octet Rule
Atoms are the most stable when they have completely full valence shells (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

10. Electron-Dot Notation
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
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

11. Electron-Dot Notation Sample Problem
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Electron-Dot Notation Sample Problem
a. Write the electron-dot notation for hydrogen.
b. Write the electron-dot notation for nitrogen.
Solution:
Hydrogen is in group 1. It has one valence electron.
Nitrogen is in group 15. It has 5 valence electrons. H • • N • • • •

12. Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
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.

13. Lewis Structures (continued)
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
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

14. How to Draw Lewis Structures
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
How to Draw Lewis Structures
Draw the electron-dot notation for each type of atom, and count the valence electrons.
Put the least electronegative atom in the center (except H.)
Use electron pairs to form bonds between all atoms.
Make sure all atoms (except H) have octets.
Count the total electrons in your Lewis structure. Does it match the number you counted in step 1? If not, introduce multiple bonds.

15. Lewis Structures Sample Problem A
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures Sample 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. C • I • H •
C x 4 e- = 4 e-
3H x 1 e- = 3 e-
I x 7 e- = 7 e-
14 e- Total

16. Lewis Structures Sample Problem A (continued)
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures Sample 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 
14 Total e- • • H C I • • • • H

17. Multiple Covalent Bonds
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Multiple Covalent Bonds
In a single covalent bond, one pair of electrons 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

18. Lewis Structures Sample Problem B
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures Sample 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. C • O • H •
C x 4 e- = 4 e-
2H x 1 e- = 2 e-
O x 6 e- = 6 e-
12 e- Total

19. Lewis Structures Sample Problem B (continued)
Chapter 6 – Section 2: Covalent Bonding and Molecular Compounds
Lewis Structures Sample 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?  
14 Total e- H • • C • O •  • •
12 Total e- • • H

20. Formation of Ionic Compounds
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
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.

21. Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
Cations (+) and anions (-) are attracted to each other because of their opposite electrical charges.
An ionic bond is a bond that forms between oppositely-charged ions because of their mutual electrical attraction.
Visual Concept

22. Ionic Bonding and the Crystal Lattice
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
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

23. Comparing Ionic and Covalent Compounds
Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
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
Covalent
crystals
molecules
very high melting points
low melting points
hard, but brittle
usually gas or liquid
Ex: NaCl, CaF2, KNO3
Ex: H2O, CO2, O2

24. Chapter 6 – Section 3: Ionic Bonding and Ionic Compounds
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+

25. Chapter 6 – Section 4: Metallic Bonding
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.

26. 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)
Visual Concept

27. Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory
The abbreviation VSEPR (say it “VES-pur”) stands for “valence-shell electron-pair repulsion.”
VSEPR theory – repulsion between pairs of valence electrons around an atom causes the electron pairs to be oriented as far apart as possible.
Treat double and triple bonds the same as single bonds.
Visual Concept

28. VSEPR Theory (continued)
Chapter 6 – Section 5: Molecular Geometry
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.

29. VSEPR Theory (continued)
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory (continued)
2 electron pairs around a central atom will be 180o apart, and the molecule’s shape will be linear.
3 bonding pairs around a central atom will be 120o apart, and the molecule’s shape will be trigonal planar. If one of the pairs is a lone pair, the shape will be bent.

30. VSEPR Theory (continued)
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory (continued)
4 bonding pairs around a central atom will be 109.5o apart, and the molecule’s shape will be tetrahedral. If one of the pairs is a lone pair, 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.

31. VSEPR Theory Sample Problem A
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory Sample 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 are surrounding the central atom? How many are unshared pairs? The shape is bent. 
Total Electrons: 8 e- 
Octets • • H H O 4 • • • • • 2 O • H H

32. VSEPR Theory Sample Problem B
Chapter 6 – Section 5: Molecular Geometry
VSEPR Theory Sample 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 are surrounding the central atom? The shape is linear. 
Total Electrons: 16 e- 
Octets • • • • • • • O C O • • • • • • • • •
2 (double or triple bonds count the same as single)

33. Chapter 6 – Section 5: Molecular Geometry
Molecular Polarity
Molecular Polarity depends on both bond polarity and molecular geometry.
If all bonds are non-polar, the molecule is 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)

34. Intermolecular Forces
Chapter 6 – Section 5: Molecular Geometry
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 types of chemical bonds (covalent, ionic or metallic.)

35. Intermolecular Forces (continued)
Chapter 6 – Section 5: Molecular Geometry
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. O H •
Visual Concept

36. Types of Intermolecular Forces
Chapter 6 – Section 5: Molecular Geometry
Types of Intermolecular Forces
3 types of intermolecular forces (strongest to weakest):
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 exists between atoms of H and N, O or F.
Induced dipole – between a polar and a non-polar molecule.
London dispersion forces – instantaneous dipoles created by the constant motion of electrons.
Visual Concept

37. Chapter 10
States of Matter

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

39. 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 very high 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.

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

41. Chapter 10 – Section 2: 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.

42. Chapter 10 – Section 2: Liquids
Liquids (continued)
Vaporization – A liquid or solid changing to a gas.
Evaporation – particles escape from the surface of a liquid and become a gas. This occurs because liquid particles have different 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.

43. Chapter 10 – Section 3: 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.

44. 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.