Chapter Four: Forces Between Particles

2, 12, 14, 20, 22, 26-32, 36, 38, 48-58,

62, 66-74

Chemical Bonding Review• Compounds and Molecules are held

together by chemical bonds• Three types of bonds

– Ionic• Metals and non-metals

– Covalent• Non-metal and Non-metal

– Metallic• Between atoms of metals

Octet Rule• All atoms strive to have electronic

configurations like the Noble Gases• Eight electrons in the outermost shell,

highest principle quantum number (n)• Except H and He follow duet rule

– Want two electrons in outermost shell

• How do the atoms achieve an octet?

Taking or Giving and Sharing Electrons

• Ionic Bonds– Atoms take or give electrons from other atoms

• Covalent Bonds– Atoms share electrons between themselves

• Metallic Bonds– Sea of electrons

n=1

87654321

87654321

87654321

87654321

87654321

87654321

1

n=2

n=3

n=4

n=5

n=6

n=7

Valence Electron Review2

Lewis Dot Structures For Atoms/Ions

• Symbol represents the nucleus and all electrons except for those in the valence shell

• Give the Lewis Dot Structure for:Na F O2-

• Species with the same number of electrons are isoelectricO2- F- Ne Na+

Mg2+

How many electrons does each species have?

Lewis Dot Structures• GN Lewis developed the theory of

covalent bonding• Structures showing covalent bonds are

called Lewis structures• Each line represents a shared pair of

electrons (2 electrons)• Lone pairs of electrons are shown by a

pair of dots

Drawing Lewis Structures• Decide on atom connectivity and placement

– Hydrogen (never in the middle) is frequently bonded to oxygen

– Oxygen is rarely the central atom– Oxygen will not bond to oxygen (except O2 or O3)– Carbon will be the central atom– Least electronegative atom is in the middle

Drawing Lewis Structures

• Count the total number of valence electrons– An atom’s number of valence electrons is equal to its

group number• Determine the total number of shared electrons

electrons needed – valence electrons present• Connect the atoms with single bonds

– A single bond is one shared pair of electrons• Use lone pairs and/or multiple bonds to give

each atom an octet of electrons

Lewis Structure (Single Bonds)

• Draw Lewis Structures for:• H2O

• HCl

• NH3

Lewis Structures (Multiple Bonds)

• CO2

• N2

Ions• Definition: Ions are atoms or groups of

atoms with an electrical charge• Cations: are positively charged, due to

loss of electrons (Metals)• Anions: are negatively charged, due to

gain of electrons (Non-Metals)• Number of electron’s gained or loss is due

to atoms wanting Octet

Examples of Ions

• Na• Ra• Al• Se• O• Cl• F

Ionic Compounds

• Ionic compounds are held together by ionic bonds, or the attraction of oppositely charged ions

• In the solid state, ionic compounds form crystalline lattices– Cations are attracted to

all the neighboring anions, not just one

– Thus, there are no discrete ionic “molecules”

Ball and Stick Model

Transition Metal Cations• Most transition metals form more than 1 cation

+1 only Ag+

+2 only Zn2+, Cd2+

+1 and +2 Hg22+, Cu+

Hg2+, Cu2+

+2 and +3 Cr2+, Fe2+, Co2+

Cr3+, Fe3+, Co3+

+2 and +4 Sn2+, Pb2+

Sn4+, Pb4+

Polyatomic Ions NO3

-

nitrate

SO42-

sulfate

PO43-

phosphate

NO2-

nitrite

SO32-

sulfite

HPO42-

monohydrogenphosphate

CO32-

carbonate

NH4+

ammonium

H2PO4-

dihydrogenphosphate

HCO3-

BicarbonateOr hydrogen carbonate

OH-

hydroxideC2H3O2

-

acetate

Formulas of Ionic Compounds

• The net charge on a formula unit must be zero

S (+) charges = S (-) charges• Since there are no ionic “molecules” the

formula of an ionic compound is the simplest ratio of cation to anion that gives an electrically neutral combination

Al3+ and O2-

Ca2+ and O2-

Writing Ionic Compound Formulas

• Write the formula for each of the following pairs of ions

• Na and Oxygen• Mg and Fluorine• Rb and Iodine

Nomenclature • Rules for naming compounds and molecules• Anions

– Name the element, drop the ending leaving the root and add “ide”• Element – root + “ide”

• Cl• O• N• S• I

Naming Ionic Compounds1. Name the cation by naming the element

• If the cation is a transition metal you need to distinguish the charge using Roman Numerals• Fe2+ is named Iron (II)• Pb4+ is named Lead (IV)

2. Name the anion • Can be an elemental anion or polyatomic

3. Combine them as two words

Naming Ionic Compounds

K2O

Li2CO3

K2SO4

NaHCO3

Cr2O3

Formulas from Names

• What are the formulas of these compounds?

calcium sulfide

iron (III) acetate

Chromium (III) sulfate

Naming Molecular Compounds• Name each element• Indicate how many of each element is

present with a prefix multiplier– Mono =1; di =2; tr i=3; tetra =4; penta =5;

hexa =6; hepta = 7; octa = 8; nona = 9; deca = 10

• Add the suffix “ide” to the last element• The prefix multiplier mono is left off of the

first element in the compound

Naming Molecular Compounds: Examples

• IBr

• NI3

• N2O4

Formulas from Names

• Sulfur dioxide

• Diphosphorous pentoxide

• Carbon tetrachloride

Molecular Compounds: Common Names

• These compounds have common (non-systematic names)– Water (H2O)

• Dihydrogen monoxide– Ammonia (NH3)

• Nitrogen trihydride– Methane (CH4)

• Carbon tetrahydride– Nitrous oxide (N2O)

• Dinitrogen monoxide– Hydrazine (N2H4)

• Dinitrogen tetrahydride

Acids

• Acids are compounds that can donate an hydrogen ion (H+ ion)

• Acids fall into two categories– Binary Acids HX– Oxoacids HXOn

• Polyatomic anions

Binary Acids

• Most binary acids result from dissolving the corresponding molecular compound in water

• Binary acids are named as hydro (stem name of X) ic acid

HCl(g) HCl(aq)

HCN(g) HCN(aq)

Oxoacids

• Oxoacids are named based on the oxoanion

• “Ate” anion => ic acid• “Ite” anion => ous acid

Oxoacids

CO32- (carbonate anion) H2CO3 (carbonic acid)

NO2- (nitrite anion) HNO2 (nitrous acid)

NO3- (nitrate anion) HNO3 (nitric acid)

PO43- (phosphate anion) H3PO4 (phosphoric acid)

SO32- (sulfite anion) H2SO3 (sulfurous acid)

SO42- (sulfate anion) H2SO4 (sulfuric acid)

Polyatomic Anion

Review of What We Know

• We can write formulas• We can name compounds and molecules• We can draw Lewis Structures

– But what do these molecules look like?

VSEPR Theory

• VSEPR: Valence Shell Electron Pair Repulsion

• Like charges repel and want to be as far apart as possible

• Therefore a given combination of electrons will form into a specific shape

VSEPR

1. Draw the Lewis Structure

2. Assign the central atom (A)

3. Determine the number (n) of atoms bonded to (A) designate them (Xn)

4. Determine the number of lone pairs on (A) designate them (Em)

5. Put together the AXnEm notation

X + E = 2

X + E = 3

X + E = 4

X + E = 5

X + E = 6

VSEPR Examples

• What is the geometry of • CO2

• BF3

• H2O

• NH3

Electronegativity• Linus Pauling developed the

electronegativity scale• Electronegativity is a measure of an

atom’s affinity for electrons• Fluorine is the most electronegative

element (EN=4.0)• The closer an atom is to fluorine,

the more electronegative it is

Polar Covalent Bonds• If two atoms of identical electronegativity are

bonded together, the bond is non-polar• If two atoms of different electronegativity are

bonded together, the bond is polar, and the electrons spend more time around the more electronegative atom– This creates partial charges

• The greater the difference in EN between two atoms, the more polar the bond– The limiting example of this is the ionic bond

EN Type of Bonding0.0 Pure covalent bond

(equal sharing of e- ‘s)0.1 – 0.4 Non-polar covalent bond

(almost equal attraction for shared e- pairs)

0.5 – 1.4 Polar covalent bond (unequal sharing of e- ’s)

1.5 – 3.2 Ionic bond (e- transfer)

Example

• The bond in hydrogen is

• The bond in hydrogen chloride is

Molecular Polarity

• Bond dipoles are vectors• The vectoral sum of the bond dipoles gives the

molecular dipole• Based on the shape of the molecules you can

predict if the dipoles will cancel each other or if they will create a dipole moment

• If a dipole moment exists then the molecule is said to be polar

• If no dipole moment exists then the molecule is said to be non-polar

Molecular Polarity Examples

• Is carbon dioxide polar or non-polar?

• Is water polar or non-polar?

• Is boron trifluoride polar or non-polar?

Intermolecular Forces

• These are attractive forces between molecules or atoms or ions

• Immensely important– These forces hold DNA molecules in a helix

and and are the mechanism for DNA transcription

Dipole Dipole Attraction

• This is the attraction between the opposite (partial) charges of polar molecules

H Cld+ d-

H Cld+ d-

Hydrogen Bonding• This is generally stronger than dipolar

attractions• Hydrogen bonding occurs between a

hydrogen atom and O, N or F.• For H-bonding to happen the H must be

directly bonded to a O, N or F.

O—H

H

O—H

H This is an attraction not really a bond

• Also called Van der Waal’s forces, these are created by instantaneous dipoles

• London forces are much weaker than either dipole-dipole or H-bonding

• London forces get stronger with larger atoms/molecules

London Forces

HeHe

London Forces Between Helium Atoms

HeHe-d +d HeHe-d +d -d +dHeHe+d +d-d -d

HeHe

At a given instant, the electrons on an atommay be non-symmetrically distributed.This leads to creation of a temporary dipole.

As the electrons re-distribute, the dipoles andand the attraction vanishes.

This dipole induces temporary dipoles on neighboring atoms.

For the merest fraction of time, there is adipole-dipole attraction between the atoms.

Ion Dipole Attraction

• This is the attraction between an ionic charge and a polar molecule

• This attraction allows ionic solids to dissolve in water

• The strength of this force varies widely and depends on the magnitude of the dipole moment of the polar species and the size of the ion

A Sodium Ion and a Chloride Ion Hydrated by Water Molecules

Effects of Intermolecular Forces

• More intermolecular forces mean:– Higher boiling and melting points– More viscous liquids

• IM Forces also affect solubility– ‘like dissolves like’

Predicting Boiling Points based on IMF’s

SnH4, CH4, GeH4, SiH4 HBr, HI, HCl, HF

Trends in Boiling Point

H2O

H2S

H2Se

H2Te

Example

• Is carbon dioxide soluble in water? Explain

Example

• Are ionic compounds more soluble in water or in gasoline (a non-polar solvent)? Explain