CHAPTER 12 – CHEMICAL BONDING

CHEMICAL BOND – A force that holds two or more atoms together as a unit

Individual atoms will naturally bond together to achieve a lower energy state (to be more stable)

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TYPES OF BONDS

Metal atoms easily lose electrons forming positive ions, and nonmetal atoms easily gain electrons forming negative ions

1) METAL ATOMS AND NONMETAL ATOMS

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IONIC BOND – The electrostatic attraction between positive and negative ions

Ionic bonding forms giant crystalline networks containing billions of positive and negative ions that are strongly attracted together

Ionic bonding exists between metal and nonmetal ions

+

Fe atoms O atoms (molecules) Fe ions and O ions

Elemental Iron Elemental Oxygen Rust

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Nonmetal atoms attract each other’s valence electrons, and share the valence electrons between pairs of atoms

Covalent Bond – The electrostatic attraction of shared electrons to the nuclei of bonding nonmetal atoms

Covalent bonding forms individual units called molecules, and while the atoms that covalently bond together strongly attract each other, the molecules that are created weakly attracted each other

Covalent bonding exists between nonmetal atoms

2) NONMETAL ATOMS

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+

C atoms Cl atoms (molecules) CCl4 molecules

Elemental Carbon Elemental Chlorine Carbon Tetrachloride

NONPOLAR COVALENT BOND – A bond in which 2 atoms are sharing electrons equally

POLAR COVALENT BOND – A bond in which 2 atoms are sharing electrons unequally

IONIC BOND – A bond in which two atoms have transferred electrons

Picture

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ELECTRONEGATIVITY – The attraction of an atom for shared electrons

The difference in the EN’s of 2 atoms tells the type of bond they make

Atom with the highest EN? Atom with the lowest EN?

EN Difference Bond

0.00.1 to 1.61.7 to 3.3

Nonpolar CovalentPolar CovalentIonic

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N-N Bond

EN for N = 3.0

3.0 – 3.0 = 0.0

Nonpolar Covalent Bond

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Polar covalent bonds have partially positive and a partially negative ends

DIPOLE MOMENT – The amount of separation of the positive and negative charges in a bond

C-O Bond

EN for C = 2.5, O = 3.5

3.5 – 2.5 = 1.0

Polar Covalent Bond

H-S Bond

EN for H = 2.1, S = 2.5

2.5 – 2.1 = 0.4

Polar Covalent Bond

C –– O+ d-d

H –– S+ d-d

DIPOLE MOMENT ARROW – Shows the direction of the dipole moment, pointing toward the negative end of the bond

C –– O H –– S

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Na-Cl Bond

EN for Na = 0.9, Cl = 3.0

3.0 – 0.9 = 2.1

Ionic Bond

Na+ Cl-

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BONDING IN IONIC COMPOUNDS

Atoms form ions to obtain a stable, octet electron arrangement

Sodium chloride

Na . .. Cl :

. .

A sodium chloride crystal is a symmetrical array of sodium and chloride ions in a 1:1 ratio

EMPIRICAL FORMULA – The simplest whole number ratio of atoms of different elements in a compound

Empirical Formula: NaCl

+ -

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Magnesium fluoride

Mg . . . F :

. .

. . . F :

. .

+ -2 -

Empirical Formula: MgF2

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Potassium nitride

Empirical Formula:

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SIZES OF ATOMS AND IONS

Positive ions are smaller than their neutral atoms and negative ions are bigger than their neutral atoms

Na atom Cl atom

Na+ ion Cl- ion

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1) The more energy levels an atom or ion has the larger it will be

2) With the same number of energy levels, the more protons an atom or ion has the smaller it will be

Li F Na Cl

ElectronsEnergy LevelsProtonsBig to Small

3233rd

9294th

113

111st

173

172nd

21

4th

1029

2nd

102

113rd

183

1st

Li+ F- Na+ Cl-

ISOELECTRONIC – Ions or atoms with the same number of electrons

Sizes of atoms or ions are determined by

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BONDING IN COVALENT MOLECULES

Atoms share electrons to obtain a stable, octet (or duet) arrangements

Water (H2O)

H . . . O :

.

H

. .H – O :

H

← LONE PAIR← BONDING PAIR

LEWIS STRUCTURE – A diagram using electron dot notation to show how the valence electrons are arranged among bonded atoms

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To draw a proper Lewis Structure:

1 – Add up the valence e-s for all of the atoms in the molecule or ion

2 – Draw a skeletal structure by using pairs of electrons to make bonds

4 – If octets are not produced, make the atoms that have octets share more e- pairs with atoms that do not have octets

3 – Complete octets (or duets for H) for all atoms, outer atoms first, using the remaining valence e-s

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Sulfur dichloride, SCl2

6 + 7 + 7 = 20 valence e-s

Cl S Cl

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Phosphorus tribromide, PBr3

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Ammonia, NH3

5 + 1 + 1 + 1 = 8 valence e-s

H N H

H

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Methane, CH4

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Fluorine, F2

7 + 7 = 14 valence e-s

F F

SINGLE BOND – One shared pair of e-s between two atoms

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Oxygen, O2

6 + 6 = 12 valence e-s

O O

DOUBLE BOND – Two shared pairs of e-s between two atoms

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Nitrogen, N2

5 + 5 = 10 valence e-s

N N

TRIPLE BOND – Three shared pairs of e-s between two atoms

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Hydrogen cyanide, HCN

1 + 4 + 5 = 10 valence e-s

H C N

Carbon disulfide, CS2

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Sulfate, SO42-

6 + 4(6)

O

O S O

O

+ 2 = 32 valence e-s

2-

Ammonium, NH4+

H

H N H

H

5 + 4(1) - 1 = 8 valence e-s

+

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Ozone, O3

6 + 6 + 6 = 18 valence e-s

O O O O O Oor

O O OO O O ↔

RESONANCE – When more than one Lewis structure can be drawn for a molecule or ion

RESONANCE STRUCTURES – The Lewis structures that can be drawn for the molecule or ion

The real ozone molecule is a average of its resonance structures

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O O OO O O ↔

O O O

2 “1½” bonds

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MOLECULAR SHAPE

VSEPR THEORY (Valence Shell Electron Pair Repulsion) – All atoms and lone pairs attached to a central atom will spread out as far as possible to minimize repulsion

A Lewis structure must be drawn to use the VSEPR Theory

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H

H C H

H

STERIC NUMBER (SN) – The sum of the bonded atoms and lone pairs on a central atom

The steric number of carbon is 4 (SN = 4): 4 bonded atoms and no lone pairs

Tetrahedral

Bond angle is 109.5°

H

C

HH H

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H N H

H

The steric number of nitrogen is 4 (SN = 4): 3 bonded atoms and 1 lone pairs

Trigonal Pyramidal

Bond angle is 108° N

HH H

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The steric number of oxygen is 4 (SN = 4): 2 bonded atoms and 2 lone pairs

Bent

Bond angle is 105° O

HH

. .H – O :

H

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Formaldehyde, H2CO

1 + 1 + 4 + 6

H

H C O

= 12 valence e-s

The steric number of carbon is 3 (SN = 3): 2 bonded atoms and 1 lone pairs

Trigonal Planar

Bond angle is 120°

O

H HC

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SiS2

4 + 6 + 6

S Si S

= 16 valence e-s

The steric number of silicon is 2 (SN = 2): 2 bonded atoms and 0 lone pairs

Linear

Bond angle is 180° S Si S

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SN

4

4

4

3

3

2

Atoms

4

3

2

3

2

2

Shape

Tetrahedral

Trigonal Pyramidal

Bent (109.5°)

Trigonal Planar

Bent (120°)

Linear

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Lone Pairs

0

1

2

0

1

0

MOLECULAR POLARITY

A BOND is polar if it has a positive end and a negative end

A MOLECULE is polar if it has a positive end and a negative end

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To determine if a molecule is polar or nonpolar:

1) Draw the correct Lewis structure

2) Draw its correct shape

3) Use EN’s to determine if the BONDS in the molecule are polar or nonpolar

4) For the polar bonds, label the positive and negative ends with δ+ and δ-

5) If a line can be drawn separating all δ+’s from all δ-’s, the molecule is polar, it not its nonpolar

. .H – O :

H

O

HH

δ+

δ-

δ+δ-

EN’s: O = 3.5, H = 2.1

3.5 – 2.1 = 1.4 the O-H BONDS are polar

All of the δ+’s can be separated from all of the δ-’s, the H2O MOLECULE is polar

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δ+

δ-

δ+

δ-

EN’s: N = 3.0, H = 2.1

3.0 – 2.1 = 0.9 the N-H BONDS are polar

All of the δ+’s can be separated from all of the δ-’s, the NH3 MOLECULE is polar

H N H

H

N

HH H δ+

δ-

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F

F C F

F

Carbon tetrafluoride, CF4 4 + 4(7) = 32 valence e-s

F

C

FF F

EN’s: C = 2.5, F = 4.0

4.0 – 2.5 = 1.5 the C-F BONDS are polar

All of the δ+’s cannot be separated from all of the δ-’s, the CF4 MOLECULE is nonpolar

δ+

δ-

δ-δ+

δ-

δ+

δ-δ+

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REVIEW FOR TEST

Electromagnetic Radiation, Photons

Ground State, Excited State

Orbital

Energy Levels

Sublevels

Orbital Notation

Electron Configuration Notation

Electron Dot Notation

Valence Electrons

Octet

Electron Pair

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REVIEW FOR TEST

Periodic Trends in

Metal, Nonmetal Activity

Atomic Radii

Ionization Energy

Electron Affinity

Ionic Bonds, Covalent Bonds

Electronegativity

Bond Polarity from Electronegativities

Ion Sizes

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REVIEW FOR TEST

Lewis Structures for

Ionic Compounds

Covalent Compounds

Resonance

Molecular Shapes

Molecular Polarity

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