Chapter 3-Chemical Bonds

Chemical Bonds

Chapter 3

Lewis Model of Bonding In 1916, Gilbert N. Lewis pointed out that the

lack of chemical reactivity of the noble gases indicates a high degree of stability of their electron configurations.

He 1s2

Ne

Ar

Kr

Xe

[He]2s2

2p6

[Ne]3s2 3p

6

[Ar]4s2 4p

6

[Kr]5s2 5p

6

Noblegas

Noble gasnotation

2 Chapter 3- Chemical Bonds

The Octet Rule

Octet rule: The tendency of group 1A-7A elements to react in ways that achieve an electron configuration of eight valence electrons.

– An atom that loses one or more electrons becomes a positively charged ion called an cation.

– An atom that gains one or more electrons becomes a negatively charged ion called a anion.


Na Chlorine atom (17 electrons): 1s2 2s2 2p6 3s2 3p5

The Octet Rule- Ion formation

Cl

Sodium atom (11 electrons): 1s2 2s2 2p6 3s1

Nonmetals Gain electrons to become anions

Metals Lose electrons to become cations


Ionic Bond Formation


+ 2+

3+

Charges of Common Ions

-

2-

2+

+

Label your Periodic Table with these ionic charges!


Naming Cations

+ 2+

3+ -

2-

2+

+


• Elements of Groups 1A, 2A, and 3A form only one type of cation.

• The name of the cation is the name of the metal followed by the word “ion”.

• Most other metals require a roman numeral to indicate charge

Chapter 3- Chemical Bonds 8

Do Not Need to Know Common Names

Naming Cations- Examples Ca2+______________ Na+______________

Al3+_________________________ Potassium ions __________

Magnesium ions ________________ Cesium ions _____________

Fe3+ ________________ Co2+_______________

Mn4+ ___________________ Gold (I) ions ______________


Naming Anions- Monoatomic


add “ide” to the stem name.

Anio n

F-

Cl-

Br-

I-

O2-

S2-

Ste m

name

fluor

chlor

bro m

iod

o x

sulf

A nion

name

fluoride

chloride

bro mide

iodide

oxide

sulfide

Naming Polyatomic Ions

NH4+

OH-

NO2-

NO3-

CH3 COO-

CN-

MnO4-

CO32 -

HCO3-

SO32-

HSO3-

SO42-

PO43 -

HPO42 -

H2PO4-

HSO4-

CrO42 -

Ammonium

Hydroxide

Nitrite

Nitrate

Acetate

Cyanide

Permanganate

Carbonate

Hydrogen carbonate (Bicarbonate)

IonName

Sulfite

Hydrogen sulfite (Bisulfite)

Sulfate

Phosphate

Hydrogen phosphate

Dihydrogen phosphate

Name

Hydrogen sulfate (Bisulfate)

Chromate

Ion

These must be memorized!!! 11 Chapter 3- Chemical Bonds

Predicting formulas of Binary Ionic Compounds

• Metal (+)is always first, nonmetal (-) last

• Must have charge balance- Ex. Aluminum chloride


Al3+ 3Cl--

-

-

-

+

+

+ AlCl3

Try: Sodium Oxide Aluminum Sulfide

• Positive named first – May need a roman

numeral (you will have to figure out what the roman numeral is!!)

• Negative named second – Ending of

monoatomic anions is “-ide”

+ 2+

3+

- 2-

2+

+


CaBr2 Ba3N2

ZnS AgCl

Cd3As2 FeI3

FeF2 Au2Se

Naming Binary Ionic Compounds

• Positive named first – May need a roman

numeral (you will have to figure out what the roman numeral is!!)

– OR Ammonium (NH4

+ )

• Negative named second – Polyatomic ions

from table 3.4

+ 2+

3+

- 2-

2+

+


Naming Compounds Containing Polyatomic Ions

NaNO3 NaH2PO4 NH4OH Fe2(CO3)3 CuSO4

More Examples


NH4I CaSO4

LiCH3COO Na2SO3

K3PO4 Mg(CN)2

Zn3(PO4)2 (NH4)2SO4

Cu(OH)2 Fe2(CO3)3

Lithium acetate Sodium sulfite

Potassium phosphate Cobalt (II) nitrite

Vanadium (IV) chlorate Aluminum chlorite

Manganese (II) acetate Ammonium perchlorate

Ammonium dichromate Lithium permanganate


Ionic Compounds

• Between ions

• Cation Named (Roman Numeral)

• Anion Named (“-ide” or polyatomic

• Charge Balance

Covalent Compounds

• Electrons are shared between atoms

• Pair of shared electrons is a covalent bond

• Prefix system used

Forming a Covalent Bond

A covalent bond is formed by sharing one or more pairs of electrons.

– The pair of electrons is shared by both atoms and, at the same time, fills the valence shell of each atom.

– Example: in forming H2

+H H

the s ingle line represents a shared pair of electrons

.. H H


C:/Users/Marie Dunn/PowerLecture GOB/Multi_media/General_Organic_and_Biochemistry_Power_Lecture/dswmedia/hmswf/covalent_bonding.swf

Polarity of Covalent Bonds

Although all covalent bonds involve sharing of electron pairs, they differ in the equality of the sharing: – Nonpolar covalent bond: Electrons are shared equally. – Polar covalent bond: Electron sharing is not equal. – The equality of the sharing depends on the relative

electronegativities of the bonded atoms. Table 3.6 Classification of Chemical Bonds

Type of Bond

Less than 0.5

0.5 to 1.9

Greater than 1.9

Nonpolar covalent

Polar covalent

Ionic

Two nonmetals or a

nonmetal and a metalloid

Electronegativity

Difference Between

Bonded Atoms

A metal and a nonmetal

Most Likely to

Form Between


Polarity of Covalent Bonds

Examples:

H-Cl

BondDifference in Electronegativity Type of Bond

3.5 - 2.1 = 1.4

3.0 - 2.1 = 0.9

4.0 - 0.9 = 3.1

2.5 - 1.2 = 1.3

polar covalent

polar covalent

ionic

polar covalent

2.5 - 2.5 = 0.0 nonpolar covalent

3.0 - 2.1 = 0.9 polar covalent

O-H

N-H

Na-F

C-Mg

C-S


Polar Covalent Bonds


Polar Covalent Bonds


Using delta notation, label each atom in the

following polar covalent bonds

C-O

H-F

N-O

Non-polar Covalent Bonds

–Elements with the same electronegativity value share

the electrons in a covalent bond equally. If the electrons

are shared equally, the bond is considered nonpolar

–Would the covalent bond that occurs in the diatomic

molecules be considered polar or nonpolar?


Interpreting Lewis Structures

Table 3.7 Lewis Structures for Several Small Molecules. (The number of valence electrons is given in parentheses after the molecular formula.)

Carbonic acidFormaldehydeAcetyleneEthylene

Hydrogen chlorideMethaneAmmoniaWater

H

H N H C H H ClH

H

C C

H

C C HH

H

C

HH

O

H

H2O (8) NH3 (8) CH4 (8) HCl (8)

C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)

H

HH

OH

O OC HH

O


• Determine the total number of valence electrons present in the molecule – Valence electrons for an atom = position with in a row (i.e. Na = 1, C = 4, He = 2)

– If the species is an ion, add electrons if the charge is negative, subtract electrons if the charge is positive

• Write the symbols of the atoms arranged according to what is bonded to what – If only 2 elements are present, the element of which there is only one atom is central (example PF3)

– If more than 2 elements are present, bonding occurs in the same order as the formula is written (example: HCN)

– Hydrogen and fluorine are never central. (these elements can only form one single bond)

– Carbon is often central (it commonly makes 4 bonds)

• Place single bonds (2 electrons) between each bonding pair of atoms

• Complete octets for all atoms EXCEPT the central atom

• Compare number of electrons drawn in structure to total number of valence electrons calculated in step one. – If there are any remaining valence electrons, place them on the central atom

– use pairs when possible-sometimes this will result in MORE than an 8 electrons for the central atom

• If there is LESS than an octet on the central atom, form multiple bonds by shifting nonbonding pairs from outer atoms. – Don’t throw in double bonds prematurely! This is the final step after every other condition has been met!


How To Draw An Electron Dot Picture (Lewis Structure)

Simple Lewis structure examples

CCl4

H2O

HCN

NO3-


Lewis Structures Practice problems:

– Draw a Lewis structure for hydrogen peroxide, H2O2.

– Draw a Lewis structure for methanol, CH3OH.

– Draw a Lewis structure for acetic acid, CH3COOH.


Exceptions to the Octet Rule

PCl5 BH3

XeF4 I3

-

Look at OWL for excellent guided learning. Lesson 3.7 (a-k)


Expanded Valence of Sulfur and Phosphorus


Recognize as

Expanded valence

Resonance

Linus Pauling - 1930s – Many molecules and ions are best described by

writing two or more Lewis structures. These molecules or ions are said to exhibit resonance.

– Individual Lewis structures are called contributing structures.

– Double-headed (resonance) arrows are placed between individual contributing structures.

– The molecule or ion is a hybrid of the various contributing structures.


Resonance Examples

NO2-

CH3 COO-


VSEPR Model

Valence-Shell Electron-Pair Repulsion (VSEPR) – Valence electrons of an atom may be involved in

forming single, double, or triple bonds or they may be unshared.

– Each arrangement of electrons creates a negatively charged region of electron density around a nucleus.

– Because like charges repel each other, the various regions of electron density around an atom spread so that each is as far away as possible from the others.



Illustrations of VSEPR

Four regions of electron density

109.5°

Tetrahedral

Tetrahedral




109.5°

Tetrahedral

Trigonal Pyramidal




109.5°

Tetrahedral

BENT



(Trigonal)

Planar 120°



Regions of electrons density surrounding central atom

Polarity


Remember Bond Polarity

POLAR Molecules

POLAR: Molecule has opposite poles (One + and one -)

LOOK FOR

• Polar Bonds

• Unshared e- pairs on central atom

• Unsymmetrical molecules

• Dipole Moments ADD


Polar Molecules

p. 96 40 Chapter 3- Chemical Bonds

Polar Molecules


NON-POLAR Molecules


Non-polar: (No opposite poles.) •Same on all corners •Di-poles cancel out

Polar or Non-Polar?

• H2S

• HCN

• C2H6


• VSEPR: An acronym that stands for V______ S______ E_______ P____ R_______ • Electron pair geometry indicates the arrangement of the electron pairs around the central atom • Molecular geometry (or molecular shape) indicates the arrangement of atoms around the central atom • The bond angle is the angle formed by any two atoms bonded to the central atom


Formula Lewis Structure Electron Pairs e-pair geometry Bond angle Molecular geometry

Polarity

Total Bonding pairs Non-bonding pairs

CO2

BCl3

O3


Formula Lewis Structure Electron Pairs e-pair geometry Bond angle Molecular geometry

Polarity

Total Bonding pairs Non-bonding

pairs

CH4

NH3

H2O

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Chapter 3-Chemical Bonds