Chapters 4 & 5

Chapters 4 & 5Chemical Bonding

Valence Electrons• Outermost electrons• s and p electrons for main group elements• Responsible for chemical properties of

atoms• Participate in chemical reactions

Core Electrons Valence Electron

Problems

• Write out the electron configurations for the following elements and identify how many core and valence electrons each has.

1) Mg2) S3) Br4) Kr

Lewis Dot Structures

• LDS: a representation of an atom using its chemical symbol surrounded by dots that signify valence electrons

Problems

• Write the Lewis Dot Structures for the following atoms

• Li• Be• Br• C• N• Ne

Li: [He]1s1

Na: [Ne]2s1

K: [Ar]3s1

Octet Rule

• Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells– Natural electron configuration of the Noble

Gases– Done by gaining, losing, or sharing electrons– Increases stability– H and He seek a “Duet”

Ionic Bonding• Ions: atoms that have a charge due to gain or loss of

electrons– Anion: (-) charged atom– Cation: (+) charged atom

• Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

Formula Unit

Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other

• Metal + Non-metal– NaCl

• Metal + Polyatomic Ion– NaNO3

• Polyatomic Ion + Non-metal– NH4Cl

• Polyatomic Ion + Polyatomic Ion– NH4NO3

• Net charge on compound equal to zero

Oxyanions

SO42- Sulfate

SO32- Sulfite

PO43- Phosphate

PO33- Phosphite

NO3- Nitrate

NO2- Nitrite

ClO4- Perchlorate

ClO3- Chlorate

ClO2- Chlorite

ClO- Hypochlorite

Rules For Naming Ionic Compounds

1) Name the cation by its elemental/polyatomic name

2) If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses

3) Next, name the anion and change its ending to “-ide”

4) If the anion is polyatomic, do not change the ending to “-ide”

5) Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present

ProblemsWrite the name for the following compounds:1)KI

2)MgBr2

3)Al2O3

4)FeCl2

5)CaSO4

6)Ba(NO2)2

7)Cu(NO3)2

Write the Formula for the following ionic compounds:8)Sodium Fluoride

9)Calcium Sulfite

10)Calcium Chloride

11)Iron (III) Oxide

12)Cobalt (II) Hydroxide

13)Ammonium Bromide

14)Ammonium Carbonate

15)Aluminum Carbonate

Iron (II) Chloride

Iron (III) Chloride

Covalent Compounds• Covalent Compounds: compounds composed of atoms bonded

to each other through the sharing of electrons• Electrons NOT transferred• No + or – charges on atoms• Non-metal + Non-metal• Also called “molecules”• Examples:

– H2O– CO2

– Cl2

– CH4

or H-H

or

Duet

Naming Covalent Compounds1) Name the first non-metal by its elemental name2) Add a prefix to indicate how many

3) Name the 2nd non-metal and change its ending to “-ide”4) Add a prefix to indicate how many

1 62 73 84 95 10

ProblemsWrite the name of the following compounds:1)CO

2)NI3

3)N2O

4)SF6

5)B2O3

Write the formula for the following compounds:

6)Phosphorous Pentachloride

7)Nitrogen Monoxide

8)Dinitrogen Tetroxide

9)Tetraphosphorous Decoxide

Problems1) KCl2) Na2S

3) H2O

4) SO2

5) K3PO4

6) FeCl3

7) (NH4)2SO4

8) SCl2

9) Cu(OH)2

10) P2O5

8) Sodium Iodide

9) Aluminum Sulfate

10) Phosphorous Pentabromide

11) Magnesium Nitride

Naming Acids

• Acids that do not contain oxygen1) Begin the name with “hydro”2) Name the anion, but change the ending to “-ic”3) Add “acid” on the end

• HCl• HF

• Acids that contain oxygen1) Do not put “hydro” at the beginning2) Begin the name with the anion3) If the anion has the ending “-ate,” change this

to “-ic acid”4) If the anion has the ending “-ite,” change this

to “-ous acid”

1) HClO4

2) HClO3

3) HClO2

4) HClO

Problems• Name the following

1) HBr(g)2) HBr(aq)3) HNO2(aq)4) HNO3(aq)5) HI (aq)6) HI (g)7) H2CO3 (aq)8) H3PO4 (aq)9) H3PO3 (aq)10) HCN (aq)

Molecular Structures

Water

Ball & Stick Models Space-Filling Models

Methane

Ethanol

Lewis Dot Structures1) Count the total number of valence electrons in the

molecule. Ex: PCl3

2) Use atomic symbols to draw a proposed structure with shared pairs of electrons.• Atoms don’t tend to bond to other atoms of the same

element when they can avoid it• Exception: Carbon

3) Place lone pair electrons around each (except H) to satisfy the octet rule, beginning with the terminal atoms

4) Place any leftover electrons on the central atom5) If the number of electrons around the central atom is less

than 8, change single bonds to the central atom to multiple bonds (double or triple).• Ex: CH2O

ProblemsDraw the LDS’s for the following molecules:

1)Cl2O

2)C2H4

3)C2H6O

What Things Like To Do• Halogens

– Like to be terminal– Like to have one single bond and 3 lone pairs

(non-bonding electrons)

• Carbon– Likes to have 4 single bonds and no lone pairs

• A double bond counts as two singles• A triple bond counts as three singles

– Likes to be central– Likes to bond to other carbons

• Silicon– Likes to do what carbon does

• Oxygen– Likes to have two single bonds and 2 lone pairs

• Sulfur– Likes to do what oxygen does

• Nitrogen– Likes to have 3 single bonds and one lone pair

• Phosphorous– Likes to do what nitrogen does

• Hydrogen– Likes to be terminal with only one single bond– No lone pairs!

Problems1) SH2

2) C3H8

3) Si2H6

4) PI3

5) CH3OH

6) C2H2

7) CCl2O

8) N2H4

9) CH2OS

10) C2H6O

11) CO

12) BrHO

Electronegativity• The measure of the ability of an atom to

attract electrons to itself– Increases across period (left to right) and– Decreases down group (top to bottom)– fluorine is the most electronegative element– francium is the least electronegative element

Electronegativity Scale

Types of Bonding

1) Non-Polar Covalent Bond:• Difference in electronegativity

values of atoms is 0.0 – 0.4• Electrons in molecule are

equally shared• Examples: Cl2, H2, CH4

ENCl = 3.03.0 - 3.0 = 0

Pure Covalent

2) Polar Covalent Bond:• Difference in

electronegativity values of atoms is 0.4 – 2.0

• Electrons in the molecule are not equally shared

• The atom with the higher EN value pulls the electron cloud towards itself

• Partial charges• Examples: HCl, ClF, NO

ENCl = 3.0ENH = 2.1

3.0 – 2.1 = 0.9Polar Covalent

3) Ionic Bond: • Difference in EN

above 2.0• Complete transfer of

electron(s)• Whole charges

ENCl = 3.0ENNa = 1.0

3.0 – 0.9 = 2.1Ionic

Problems

• Predict the type of bonding in the following compounds using differences in EN values of the atoms. Indicate the direction of the dipole moment if applicable

1) KBr2) HF3) BrI4) FI

Valence Shell Electron Pair Repulsion Theory

• VSEPR theory:– Electrons repel each other– Electrons arrange in a

molecule themselves so as to be as far apart as possible

• Minimize repulsion• Determines molecular

geometry

Defining Molecular Shape• Electron pair geometry: the geometrical

arrangement of electron groups around a central atom– Look at all bonding and non-bonding e-’s

• Molecular Geometry: the geometrical arrangement of atoms around a central atom– Ignore lone pair electrons

• 2 e- groups surrounding the central atom– e- pair geometry: linear– MG: linear– AXE designation: AX2E0

• A: Central Atom• X: Bonding pairs• E: Non-bonding pairs

– Example: BeCl2

3 e- groups• 3 Bonds, 0 Lone Pairs

– e- PG: Trigonal Planar (Triangular planar)

– MG: Trigonal Planar– AX3E0

– BF3

• 2 Bonds, 1 Lone Pair– e- PG: Trigonal Planar

(Triangular planar)– MG: Bent/angular– AX2E1

– GeCl2

4 e- groups• 4 bonds, 0 Lone Pairs

– e- PG: Tetrahedral– MG: Tetrahedral– AX4E0

– CH4

• 3 bonds, 1 Lone Pair– e- PG: Tetrahedral– MG: Triangular Pyramidal– AX3E1

– NH3

• 2 bonds, 2 Lone Pairs– e- PG: Tetrahedral– MG: Bent/Angular– AX2E2

– H2O

Drawing LDS With Correct Geometry

Molecular Polarity

Problems

1) NF3

2) CH2O

3) CBr4

4) CHCl3

5) CH2Cl2

Draw the Lewis Dot Structures for the following molecules and then identify the direction of polarity, if any.