Chemical Bonds

Bettelheim, Brown, Campbell and Farrell

Chapter 3

Ionization Energy

• Energy required to remove outermost electron (most loosely held)

Ionization energy

Noble Gas Configurations

He 1s2

Ne

Ar

Kr

Xe

[He]2s2 2p6

[Ne]3s2 3p6

[Ar]4s2 4p6

[Kr]5s2 5p6

Noblegas

Noble gasnotation

He

8A

Ne

Ar

Kr

Xe131.3

Rn(222)

BoilingPoint(°C)

MeltingPoint(°C)Element

HeliumNeonArgonKryptonXenon

-272-249-189-157-112

-269-246-186-152-107

Radon -71 -62

2

4.00310

20.1818

39.9536

83.8054

86

Noble gas configuration s2p6 very stable

The Octet Rule

• Octet rule:Octet rule: Group 1A-7A elements to achieve an outer shell of eight valence electrons

• Anion: Anion: Negative ion formed when an atom gains electrons

• Cation:Cation: Positive ion formed when an atom loses electrons

The Octet Rule—Cations

Cation: Sodium atom loses an electron to form a sodium ion, which has the same electron configuration as neon

Na (11 electrons): 1s2 2s2 2p6 3s1

Na+ (10 electrons): 1s2 2s2 2p6

Na Na+ + e-

A sodiumatom

A sodiumion

Anelectron

The Octet Rule—Anions

Anion: Chlorine atom gains an electron to form a chloride ion, which has the same electron configuration as argon

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

chloride ion (18 electrons): 1s2 2s2 2p6 3s2 3p6

Anelectron

Cl Cl-+ e-

A chlorineatom

A chlorideion

Cation Names

• Groups 1A, 2A, and 3A – The name of the element followed by the word

“ion”

Li+H+ Hydrogen ion

Lithium ion

Sodium ionPotassium ion

Ion

K+

Na+

Mg2+ Magnesium ion

Calcium ion

Strontium ionBarium ion

Ca2+

Sr2+

Ba2+

Al3+ Aluminum ion

Name Ion Name Ion Name

Group 1A Group 2A Group 3A

Transition Metal Cations• Cations derived from transition and inner

transition elements more than one type of cation

• Stock System (IUPAC): – Use Roman numerals to show charge: – Fe2+ is Iron (II) Fe3+ is Iron (III)– Cu+ is Copper (I) Cu2+ is Copper (II)

• Old System:– Use the suffix -ous-ous to show the lower positive charge

and the suffix -ic-ic to show the higher positive charge– Fe2+ is Ferrous Fe3+ is Ferric– Cu+ is Cuprous Cu2+ is Cupric

Transition Metal Ion Names

Fe3+

Fe2+

Hg+

Hg2+

Cu2+

Cu+ Copper(I) ionCopper(II) ion

Iron(II) ionIron(III) ion

Mercury(I) ionMercury(II) ion

Cuprous ionCupric ion

Ferrous ionFerric ion

Mercurous ionMercuric ion

Cupr- from cuprum, the Latinname for copper

Hg from hydrargyrum, theLatin name for mercury

IonSystematic name

Common name

Origin of the symbol of theelement or the common name of the ion

Ferr- from ferrum, the Latinname for iron

Sn2+

Sn4+

Tin(II) ionTin(IV) ion

Stannous ionStannic ion

Sn from stannum, theLatin name for tin

Anion Names

• Add “ide” to the root name of the element

Anion

F-

Cl-

Br-

I-

O2-

S2-

Stemname

fluorchlorbromiodoxsulf

Anionname

fluoridechloridebromideiodideoxidesulfide

Polyatomic Ions– Contain two or more atoms– Common names often used (in parentheses)

NH4+

OH-

NO2-

NO3-

CH3COO-

CN-

MnO4-

CO32-

HCO3-

SO32-

HSO3-

SO42-

PO43-

HPO32-

H2PO4-

HSO4-

CrO42-

Ammonium

HydroxideNitrite

NitrateAcetate

Cyanide

Permanganate

Carbonate

Hydrogen carbonate (Bicarbonate)

IonName

SulfiteHydrogen sulfite (Bisulfite)

Sulfate

Phosphate

Hydrogen phosphate

Dihydrogen phosphate

Name

Hydrogen sulfate (Bisulfate)

Chromate

Ion

Naming Ionic Compounds

• Name the positive ion first, then the negative ion

• Number of each ion not included

NaBr Al2O3

MgSO4 K2S

(NH4)3PO4

Naming Ionic Compounds

NaBr Sodium bromide

Al2O3 Aluminum oxide

MgSO4 Magnesium sulfate

K2S Potassium sulfide

(NH4)3PO4 Ammonium phosphate

Formulas of Ionic Compounds

• The total number of positive charges must equal the total number of negative charges

– Li+ and Br- form LiBr (+1) + (-1)

– Ba2+ and I- form BaI2 (+2) + 2(-1)

– Al3+ and S2- form Al2S3 2(+3) + 3(-2)

Forming Chemical Bonds

• Ionic bond:Ionic bond: the force of electrostatic attraction between a cation and an anion– Atom loses or gains electrons to make a filled

valence shell (octet) and become an ion.

• Covalent bond:Covalent bond: a pair of electrons that are shared by two atoms– Atom shares electrons to make a filled

valence shell (octet)

Forming an Ionic Bond--NaCl

• Formation of sodium chloride, NaCl

Single-headed curved arrow used to show the transfer of the electron

Na + Cl Na+ Cl -

+Na+(1s22s22p6)

Cl(1s22s22p63s23p5)+Na(1s22s22p63s1)

Cl-(1s22s22p63s23p6)

Sodium atom Chlorine atom

Sodium ion Chloride ion

NaClformation.mov

Ionic Bonds• Force of attraction between a cation and an

anion.

• Depends on electronegativity– measure of an atom’s attraction for shared pair

of electrons in chemical bond with another atom)

- -

Electronegativity increases

Covalent Bonds

• Result of one or more pairs of electrons that are shared by two atoms– Each atom has full valence shell (octet)

• In H2, each hydrogen contributes one electron to the single bond

+H H

the single line represents a shared pair of electrons

.. H H

Molecular Compounds

• Molecular compound:Molecular compound: only covalent bonds

• Naming molecular compounds– the less electronegative element is named first (it

is generally written first in the formula)– prefixes “di-”, tri-”, etc. are used to show the

number of atoms of each element; the prefix “mono-” is generally omitted

• Exception: carbon monoxide• NO is nitrogen oxide (nitric oxide)

• SF2 is sulfur difluoride

• N2O is dinitrogen oxide (laughing gas)

Electronegativity

F has highest value Noble gases have 0 value

Ionic bonds form when electronegativity difference ≥ 1.9

Polarity of Bonds

• Nonpolar:Nonpolar: Electrons are shared equally

• Polar:Polar: Electrons are NOT shared equally

Type of Bond

less than 0.5

0.5 to 1.9

greater than 1.9

nonpolar covalent

polar covalent

ionic

two nonmetals or anonmetal and a metalloid

ElectronegativityDifference Between

Bonded Atoms

a metal and a nonmetal

Most likely to be Formed Between

Polarity of Covalent Bonds

H-Cl

BondDifference in Electronegativity Type of Bond

3.5 - 2.1 = 1.43.0 - 2.1 = 0.94.0 - 0.9 = 3.12.5 - 1.2 = 1.3

polar covalentpolar covalentionicpolar covalent

2.5 - 2.5 = 0.0 nonpolar covalent

3.0 - 2.1 = 0.9 polar covalentO-HN-HNa-FC-MgC-S

Polarity of Covalent Bonds

More electron density shown by δδ-- or the head of a crossed arrow

Less electron density shown by δδ++ or the tail of a crossed arrow

Polarity of Molecules• Polar molecule

– has polar bonds, and– Has partial positive and partial negative

charges in different parts of molecule, i.e., is a dipole (has two poles)

• Carbon dioxide, CO2, has two polar bonds but, because of its geometry, the pulls balance out so it is a nonpolar molecule

O C O-- +

Carbon dioxide(a nonpolar molecule)

::

::

Polarity of Molecules

• Water, H2O, has two polar bonds and, because of its geometry, is a polar molecule

OH H

-

+center of partial positivecharge is midway betweenthe two hydrogen atoms

Water(a polar molecule)

Polarity of Molecules

• Both dichloromethane, CH2Cl2, and formaldehyde, CH2O, have polar bonds and are polar molecules

C

H HFormaldehyde

O+-

Dichloromethane

C

H H

Cl Cl+--

Lewis Structures

• Used to decide on the arrangement of atoms in the molecule

• Bonding (shared) electrons are shown as bonds (lines)

• Nonbonding electrons are represented as a pair of Lewis dots

Drawing Lewis Structures

1. Determine the number of valence electrons in the molecule

2. Decide on the arrangement of atoms in the molecule

3. Connect the atoms by single bonds

4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots

5. In a single bondsingle bond, atoms share one pair of electrons; in a double bonddouble bond, they share two pairs, and in a triple bondtriple bond they share three pairs.

Exceptions to the Octet Rule

• H and He have a maximum of 2 electrons (duet)

• Period 2 elements have a maximum of 8 electrons (use 2s and 2p orbitals)

• Atoms of period 3 elements may have more than 8 electrons

Lewis Structures• Examples:Examples: (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

HC C

HC C HH

HC

HHO

H

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

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

H

HHO

H

O OC HH

O

Lewis Structures

• ExamplesNH3

CH3OH

CH3COOH

.

Valence-Shell Electron-Pair (VSEPR) Model

• Because like charges repel each other, the various regions of electron density around an atom spread so that they are as far away from each other as possible– CH4: measured H-C-H bond angles

are 109.5°

H C HH

H

Summary of Molecular Geometry

# Electron Lone Bond Angle Shape regions pairs

2 0 180o Linear

3 0 120o Planar 2 1 Angular (bent)

4 0 109.5o Tetrahedral 3 1 Trigonal pyramidal 2 2 Angular (bent)

HH

NH

CH H

HC C

H

O C

HC

HHO

H

CH C HO

HH

O109.5°4

PredictedBond

AnglesExamples

(Shape of the molecule)

2 180°

120°3

PredictedDistributionof ElectronDensity

tetrahedral

trigonal planar

linear

Methane(Tetrahedral)

Ammonia(Pyramidal)

Water(Bent)

Ethylene(Planar)

Formaldehyde(Planar)

Carbon dioxide(Linear)

Acetylene(Linear)

H

H

: :

:

: :

: :

:

:

Regions ofElectron DensityAround Central

Atom

VSEPR Model

– the measured H-N-H bond angles are 107.3°– the unshared pair is not shown on this model

– the measured H-O-H bond angle is 104.5°

H N HH

H O Htetrahedron.mov