CHEMICAL BONDING NOMENCLATURE€¦ · I. CHEMICAL BONDING: IONIC AND COVALENT BONDS A. Types of Chemical Bonds II. MOLECULES A. Electron Configuration and the Octet Rule B. Orbital

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CHEMICAL BONDING

& NOMENCLATURE

"What's in a name? That which we call a rose

By any other name would smell as sweet." (William Shakespeare. Romeo and Juliet, II, ii, 1-2)

CHEMISTRY SG BONDING & NOMENCLATURE p. 1

TABLE OF CONTENTS

STUDY GUIDE 6: CHEMICAL BONDING & NOMENCLATURE

I. CHEMICAL BONDING: IONIC AND COVALENT BONDS A. Types of Chemical Bonds

II. MOLECULES A. Electron Configuration and the Octet Rule

B. Orbital Notation (electron configuration)

C. Oxidation Numbers

D. Lewis Dot Structures (Electron-Dot Notation)

3. Notes:

E. Expanded Octet (One exception to the octet rule.)

F. Formal Charges and Chemical Structures

G. Resonance and Chemical Structure

H. Hybridization

I. Sigma () and Pi () bonds

J. VSEPR (Valence-Shell Electron-Pair Repulsion Theory)

K. Intermolecular Bonding

III. IONIC BONDING AND IONIC COMPOUNDS A. Formation of Ionic Compounds

B. Comparison of Ionic and Molecular Compounds

C. Polyatomic Ions

IV. METALLIC BONDING

V. COVALENT-NETWORK

VI. NOMENCLATURE - NAMING COMPOUNDS A. Ionic Compounds

B. Molecules

C. Acids


I. CHEMICAL BONDING: IONIC AND COVALENT BONDS

A. Types of Chemical Bonds

If atoms only existed as individual units, by themselves, it would be a very limited and boring universe.

However, atoms seldom exist as independent particles. From the water that makes up most of our bodies

and planet, to the rocks and almost everything you can see, substances are composed of combinations of

atoms held together by chemical bonds. In other words, the chemical bond is the mutual attraction

between valence electron(s) on one atom and the nucleus (protons) on an adjacent atom.

Atoms often gain, lose or share electrons to satisfy the octet rule (achieve the same number of electrons

as the noble gas closest to them).

chemical bond a mutual attraction between nuclei and valence electrons of different atoms

that hold the atoms together.

Bonds can be divided into to general types:

Three types of bonds:

Bond Type Found In: Comments

ionic ionic compounds gain/loss electrons

covalent molecules share electrons

metallic metals ‘sea’ of electrons

The two most important types of compounds for our purposes are ionic (giving rise to ionic compounds)

and molecular compounds (or simply molecules). Ionic compounds are formed by ionic bonds, molecular

compounds by covalent bonds.

ionic bonds results from the attraction between large numbers of cations and anions;

produces ionic compounds.

covalent bonds results from the sharing of pairs of electrons between two atoms; produces

molecules.

A cation is formed when the atom loses one or more electrons. An anion is formed when the atom gains

one or more electrons (Figure 1). A multitude of cations combine with a multitude of anions to form an

ionic compound (Figure 2). The formula of the ionic compound (e.g., NaCl) is an empirical formula

representing the ratio of the enormous number of bonded atoms. When the electrons are shared between

atoms the bond is a covalent bond. For molecules, the chemical formula represents a single entity.

Figure 1. Formation of ions. (A.) The neutral sodium ion (Nao) loses an electron to form a sodium ion

(Na+), thereby having a complete octet of valence electrons. The sodium cation is isoelectric with the

noble gas, neon. (B). A neutral chlorine atom completes its outer valence shell by gaining an electron,

resulting in chlorine ion having the same electron configuration as the noble gas, argon. When the

electron lost by sodium is gained by chlorine, the atoms can combine with other similar ions to form an

ionic compound.


Figure 2. Formation of Ionic Compounds and Molecules.

What determines whether a chemical compound is ionic or molecular? In fact, bonding between two

different elements is rarely purely ionic or purely covalent. The degree to which two elements are bonded

is reflected in the polarity of the compound. If the valence electrons spend the vast majority of the time

equally divided between the two elements, the bond is said to be nonpolar. If, on the other hand, the

valence electrons spend most of the time around only one of the two atoms, the bond is polar.

polar an unequal (unsymmetrical) distribution of charge

nonpolar an equal (symmetrical) distribution of charge

All ionic bonds are, by definition, polar: ionic bonds are always polar. Covalent bonds can be divided

into polar covalent and nonpolar covalent bonds (Figure 3).

nonpolar covalent bond a covalent bond in which the bonding electrons are shared equally by the

bonded atoms, resulting in a balanced distribution of charge (electrons)

polar covalent bond a covalent bond in which the bonding electrons are not shared equally by

the bonded atoms, resulting in an unbalanced distribution of charge

(electrons)

A.

B.

C.

Figure 3. Comparison of the distribution of charge (electron density) for nonpolar and polar bonds. Delta

() represents a partial positive (+) or partial negative (-) charge resulting from electron distribution

(which, in turn, results from the atom’s relative electronegativity). A: nonpolar covalent bond; B: polar

covalent bond; C: ionic bond.

One can predict the type of bond – ionic, nonpolar covalent, or polar covalent – between two elements

based on the difference in electronegativity (Table 1 and Table 2). Looking at the periodic table

displaying the electronegativities of the elements, it should be apparent that the only anions capable of

forming ionic bonds are derived from nitrogen, oxygen, fluorine, and chlorine. The electronegativity of

the least electronegative atoms is 0.7. For a bond to be ionic, the difference between the atoms must be at

least 1.7. This means that the more electronegative atom in a bond must be at least 1.7 + 0.7,

or 2.4, for the bond to be ionic. Because so few of the atoms we encounter have an

electronegativity of less than 0.9, the only anions we will generally encounter that form ionic

bonds are N, O, and the halogens (F, Cl, Br, and I). When calculating bond type, only use two

atoms at a time to determine what bonds are present (e.g., in CH3OH, there’s C-H, C-O, and O-H bonds

(see small figure above). Then calculate the difference in electronegativities between each pair of two

atoms in those bonds – e.g., 2.5(C) – 2.1(H) = 0.4(polar covalent).


Table 1. Determination of the type of bond between two elements.

Dif

fere

nce

in

Ele

ctro

neg

ativ

ity 3.3

IONIC 100%

Percen

tage o

f

Ion

ic Ch

aracter

The type of bond character between two elements is

determined by the difference in their electronegativity values.

For example, the bond between sodium and chlorine (see

Table 2) is: 1.7 50%

POLAR-

COVALENT 0.3 5%

Element Electronegativity Difference

NON-POLAR

COVALENT

chlorine: 3.0

0.0 0% sodium: 0.9 = 2.1 ionic bond

Table 2. Electronegativities of the Elements.

1 2

H He 2.1 - -

3 4 5 6 7 8 9 10

Li Be B C N O F Ne 1.0 1.5 2.0 2.5 3.0 3.5 4.0 - -

11 12 13 14 15 16 17 18

Na Mg Al Si P S Cl Ar 0.9 1.2 1.5 1.8 2.1 2.5 3.0 - -

19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 0.8 1.0 1.3 1.5 1.6 1.6 1.5 1.8 1.9 1.9 1.9 1.6 1.6 1.8 2.0 2.4 2.8 - -

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 0.8 1.0 1.2 1.4 1.6 1.8 1.9 2.2 2.2 2.2 1.9 1.7 1.7 1.8 1.9 2.1 2.5 - -

55 56 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86

Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 0.7 0.9 1.3 1.3 1.5 1.7 1.9 2.2 2.2 2.2 2.4 1.9 1.8 1.9 1.9 2.0 2.2 - -

87 88 89 103 104 105 106 107 108 109

Fr Ra Ac Lr Rf Db Sg Bh Hs Mt Key: 1

0.7 0.9 1.1 - - - - - - - - - - - - - - H

2.1

atomic number

electronegativity


II. MOLECULES

Molecule A neutral group of atoms that are held together by covalent bonds (shared

valence electrons)

Molecular Compound a chemical compound whose simplest units are molecules

Diatomic Molecular a molecule containing two identical atoms. mnemonic: HONClBrIF

(hydrogen, oxygen, nitrogen, chlorine, bromine, iodine, and fluorine)

Molecular Formula the arrangement of symbols and numbers that represent how the atoms in a

molecule are combined.

C8H18

Another example: The formula unit for the ionic compound, Al2(SO4)3 (aluminum sulfate) contains

at total of 2 aluminum atoms, 3 sulfur atoms and 12 oxygen atoms.

How are molecules formed? The reason that two atoms will combine in a covalent bond is because they

are at a lower potential energy in the bonded state than when they are independent particles (Figure 4).

The atoms are brought together by the attractive force between the valence electrons on each atom and the

oppositely charged nucleus of the other atom.

Figure 4. Formation of a covalent bond.

Formation of a covalent bond results from a decrease in

potential energy. (a) Two independent atoms do not

affect each other. (b) As the two atoms approach, there

is a decrease in their potential energies as a result of the

attraction between the valence electrons on each atom

and the nucleus of the other atoms. (c) As the potential

energy reaches a minimum, the attractive forces are

balanced with the repulsion forces (i.e., electrons and

electrons; nucleus and nucleus). (d) If the nuclei

approach each other too closely, the repulsion between

like charges outweighs the attractive forces and the

potential energy increases.

Bond length the average distance between nuclei at their minimum potential energy (i.e., two

bonded atoms)

Bond energy the energy required to break a chemical bond and form two neutral isolated atoms;

expressed in kilojoules per mole (kJ/mol)

There is a relationship between bond length (the distance between two bonded atoms) and bond energy:

generally, the shorter the bond, the greater the bond energy (Table 3).

Table 3. An Example of the Relationship Between Bond Length and Bond Energy.

Bond Bond Length (pm) Bond Energy (kJ/mol)

N-N (single bond) 145 180

N=N (double bond) 125 418

NN (triple bond) 110 942

subscript representing eight carbon atoms in one molecule of octane

subscript representing 18 hydrogen atoms in one molecule of octane

symbols for the elements (carbon and hydrogen) in one molecule of octane


A. Electron Configuration and the Octet Rule

The octet rule is one of the most important concepts in chemistry, explaining how and why chemicals

react the way they do.

OCTET RULE CHEMICAL COMPOUNDS TEND TO FORM SO THAT EACH ATOM – BY GAINING,

LOSING, OR SHARING ELECTRONS – HAS AN OCTET OF ELECTRONS IN ITS

HIGHEST OCCUPIED ENERGY LEVEL. Table 4. Exceptions to the Octet Rule.

Element Comment

o hydrogen forms bonds to be surrounded by 2 electrons

o boron, aluminum typically forms bonds to be surrounded by 6 electrons (e.g., BF3)

o expanded octets elements often bonded to the most electronegative elements (O, F and Cl)

frequently involve the electrons in the d-orbitals as well as s- and p-orbitals

o In addition, the octet rule fails when molecules and polyatomic ions:

(1) have an odd number of valence electrons (e.g., ClO2, NO, NO2, and O2–)

(2) in which an atom has fewer than an octet of valence electrons (e.g., q.v., boron above)

(3) in which an atom has more than an octet of valence electrons (e.g., PCl5, SF4, AsF6–, and ICl4

–).

Let’s illustrate the octet rule in two ways – first using orbital notation and then using the Lewis or

electron-dot structures.

B. Orbital Notation (electron configuration)

The tendency of atoms to achieve an electron configuration equivalent to a noble gas can be

illustrated by the bonding of two hydrogen atoms:

individual atoms hydrogen molecule

(bonding electron pair in overlapping orbitals)

H: H:

1s 1s

H: H:

1s 1s

Figure 5. Bond formation by overlapping electron orbitals.

By sharing their 1s1 electrons, each hydrogen atom gains a 1s

2, or noble-gas electron configuration.

Another example of two atoms forming a bond to achieve noble gas electron configuration is HCl.

(Do these atoms form an ionic, polar covalent, or nonpolar covalent bond? ____________________ )

Hydrogen and chlorine atoms Hydrogen chloride molecule

Bonding electron pair in overlapping orbitals

H: H:

1s 1s

Cl: Cl:

1s 2s 2p 3s 3p 1s 2s 2p 3s 3p


C. Oxidation Numbers

When elements react to form compounds, the oxidation numbers reflect the number of electrons each

atom tries to lose or gain to become more stable.

Groups 1, 2 and 13 (Groups 1A, 2A & 3) will lose 1, 2, or 3 electrons, respectively, and

are assigned oxidation numbers of 1+, 2+, and 3+.

Typically, transition metals usually have a valance shell with 2 electrons in the s-

sublevel. They will lose these s-electrons first, giving them a typical oxidation number

of 2+. In a strongly oxidizing environment, some lower energy level electrons in the d-

sublevel can be lost, one at a time. For example, Manganese is classified as a Group 7B

element because it has a maximum of seven electrons to lose in a reaction. (It has an

electron configuration of [Ar]4s23d

5 and, thus, will have oxidation numbers ranging from 2+ to 7+.)

The other B groups follow a similar pattern. In contrast to Mn, iron ([Ar]4s23d

6) has only 2+ and 3+

oxidation numbers. For Fe to have an oxidation number above 3+, electrons would have to be removed

from the relatively stable half-filled d-sublevel (which would not normally occur).

The heavier metals in Groups 13, 14 & 15 (Groups 3A, 4A, and 5A) follow a similar

change in oxidation numbers. For example, tin ([Kr]4d10

5s25p

2) can have an oxidation

number of (a) 2+ by losing the two valence electrons from the 5p sublevel or (b) 4+ by

losing 2 additional electrons from the 5s sublevel. Electrons in the filled 4d sublevel are

too stable to be lost.

Nonmetals in Groups 15, 16 and 17 (Groups 5A, 6A, 7A) will try to gain 3, 2, or 1

electrons, respectively, and have oxidation numbers of 3–, 2–, and 1–. When two of

these elements are bonded together (e.g., IF), the nonmetal with the lower

electronegativity will be assigned a positive oxidation number and the value depends on

the number of shared valence electrons. When combined with oxygen, chlorine (having a lower

electronegativity) will usually share 1, 3, 5, or all 7 of its valence electrons – depending on how many

oxygen atoms chlorine combines with (e.g., ClO4-

, ClO3-

).

Not all inner transition metals behavior similarly. The lanthanides (left) generally have oxidation

numbers of 3+; the actinides (right), 2+.

Noble gases are remarkably unreactive (which is why they were unknown at the time

Mendeleev created his first periodic table) and almost always have an oxidation number

of zero. They do occur in some compounds – e.g., in 1962, Neil Bartlett synthesized the

first noble gas compound (XePtF6).


D. Lewis Dot Structures (Electron-Dot Notation)

1. Overview

Lewis dot structures are extremely helpful for visualizing the loss, gain or sharing of valence

electrons. To do so, determine how many valence electrons (typically electrons in the s- and p-

orbitals). Then place a corresponding number of dots around the element’s symbol (Table 5). Each

dot represents a single electron. Sometimes, ‘x’s are used instead of dots, especially when two atoms

are shown in a chemical bond (q.v., below).

Table 5. Lewis Dot Structures

Number of Valence Electrons Electron-Dot Structure Example

1

2

3

4

5

6

7

8

A. B. C.

Figure 6. Lewis Dot Structures for Selected Molecules. A. Hydrogen molecule showing two shared

electrons between the two hydrogen atoms, each donating one electron so that each, sharing their two

electrons, has an electronic configuration isoelectric with helium. B. Fluorine molecule showing

each fluorine atom’s three unshared (or lone) pair of electrons and the one pair of shared electrons.

In this figure, the electrons of one fluorine atom are written as dots, the other as x’s. C. Each pair of

shared electrons in a covalent bond can be represented by a single line. Note, that the unshared pairs

of electrons are still displayed.

unshared (or lone) pair of electrons a pair of electrons that is not involved in bonding and

belongs exclusively to only one atom in a bond.

When a single pair of electrons is shared, a single covalent bond is formed. When two pairs of

electrons are shared, a double covalent bond is formed. When three pairs of electrons are shared, a

triple covalent bond is formed. There is no quadruple bond (four electron pairs) bond.

A. B. C. D. E.

Figure 7. Multiple Covalent Bond Formation. A. Draw the Lewis dot structure for each of the two

elements in the covalent bond. B. Each atom donates a single electron for form the first, a single,

covalent bond. C/D. Each atom is still deficit for a complete octet. However, each can still donate

one more electron, forming a second – or double – covalent bond. E. The double covalent bond is

represented by a double line, and the lone pairs of electrons are arranged to be as far apart from each

other as they can be.


2. How To Draw Lewis Structures (for molecules): Steps a. Example #1: NH3

1. Predict the location of central atoms: a. central atom: select the least electronegative

atom (order generally: C, N, O) b. terminal ends: H, halogens. Leave for the final

attachments.

central terminal terminal terminal

2. Total number of valence electrons:

electrons valence8

)atom H

electrons valence1*

atom H 3()

atom N

electrons valence5*

atom N 1(

3. Draw single bonds from central to terminal atoms:

4. Assigning electrons:

a. calculate pairs of electrons remaining:

(total pairs e’s) – (pair e’s used) = pairs

available

(4 pairs) – (3 pairs used) = 1 pair available

b. remaining electron pairs include double, triple

bonds and lone pairs.

(1) place lone pairs around terminal atom

(based on atom’s Lewis dot structure)

(2) place remaining pairs around central atom (above)

c. If central atoms does not have octet, convert

terminal’s lone pairs into a double/triple bond to

central atom

(n/a)

Steps b. Example #2: CO2


atom (order generally: C, N, O) b. terminal ends: H, halogens. Leave for the final

attachments.

central terminal terminal


electrons valence61

)atom O

electrons valence6*

atom O 2()

atom C

electrons valence4*

atom C 1(

3. Draw single bonds from central to terminal atoms: 4. Assigning electrons:



available


b. remaining electron pairs include double, triple

bonds and lone pairs.

(1) place lone pairs around terminal atoms

(2) place remaining pairs around central atom


terminal’s lone pairs into a double/triple bond to

central atom


Steps c. Example #3: PO43–


atom (order generally: C, N, O) b. terminal ends: H, halogens. Leave for the

final attachments.

central terminal terminal terminal


evalenceechnegativethefromelectrons

atomO

evalenceatomO

atomP

evalenceatomP

__32)arg_____3(

)_

__6*

__4()

_

__5*

__1(

3. Draw single bonds from central to

terminal atoms:

4. Assigning electrons:



available


b. remaining electron pairs include double,

triple bonds and lone pairs.

(1) place lone pairs around terminal atoms

(2) place remaining pairs around central

atom

Octet is satisfied around P. So charge is distributed

around PO43-

.


terminal’s lone pairs into a double/triple

bond to central atom

3. Notes:

a. Triple bonds are stronger than double bonds (more electron pairs are shared), and double

bonds are stronger than single bonds.

b. Elements that are commonly encountered in multiple bonds are as follows:

Nitrogen: single, double, and triple. (The triple nitrogen bond, NN, is a very

strong and difficult bond to break.)

Carbon: single, double, and triple.

Oxygen: single and double

other elements in the same groups as N, C and O can form similar multiple bonds.

c. A single covalent bond is composed of two electrons (i.e., a pair of electrons). Most often, a

covalent bond is not shown as pair of electrons (Table 6).

Table 6. Examples of Lewis Structures with bonds written as lines.

(a) CO2 with dots showing electrons.

(b) CO2 with double-bond lines. Note lone pair

of electrons is still shown.

(c) single covalent bonds (CH4)

(d) triple covalent bonds (hydrogen cyanide)

(d) ethene (ethylene) (C2H4)

(e) ethyne (acetylene) (C2H2)


E. Expanded Octet (One exception to the octet rule.)

Atoms in the second period (e.g., Li through Ne) cannot have more than eight valence electrons.

However, atoms in the third period and above can have more than eight valence electrons

around the central atom. The d-orbitals can be used to form an expanded octet. One example

is the very stable sulfur hexafluoride (SF6). The electron configuration of the sulfur is

[Ne]3s23p

4. In SF6, each of sulfur’s six valence electrons forms a covalent bond with one

fluorine atom, resulting in twelve valence electrons around the central (sulfur) atom.

F. Formal Charges and Chemical Structures

Some molecules and ions can be drawn with more than one Lewis structure.

For example, ozone (O3) has two possible structures:

To determine which structure is most realistic, we use the formal charge for each atom. The formal

charge of any atom in a molecule is the charge the atom would have if all the atoms in the molecule had

the same electronegativity.

1. All nonbonding (lone) pairs of electrons are assigned to the atom on which they are located.

2. All bonds (single, double, or triple) are assigned a value of ½ the total number of electrons.

Table 7. Determining Formal Charge.

Two structures are possible for CO2. Although we may intuitively (or aesthetically) believe one

structure to be superior to another, we need to supply empirical evidence. For that, we use formal

charges:

A

B

O C O O C O

Number of valence electrons:

Number of valence electrons: 6 4 6 6 4 6

- Number of nonbonded electrons: 4 0 4 6 0 2

- ½ Number of bonded electrons: ½(4)=2 ½(8)=4 ½(4)=2 ½(2)=1 ½(8)=4 ½(6)=3

= Formal Charge 0 0 0 –1 0 +1

In both of the above cases, the total formal charge is zero [A: 0+0+0=0; B: (-1)+0+(+1)=0]. This is

consistent with the total charge on the CO2 molecule being zero (or neutral). However, deciding which

structure is most probably that of the actual CO2 molecule, we use the following two rules:

1. Choose the Lewis dot structure in which the atoms have a total formal charge of zero.

2. Choose the Lewis dot structure in which any negative charge resides on the most electronegative

atoms.

Thus, in the CO2 example above, structure A is the more likely to represent the ‘real’ structure.


G. Resonance and Chemical Structure

Sometimes a molecule can be drawn equally well with more than one chemical structure. In this case, the

molecule is said to have resonance.

Resonance bonding that can be represented by more than one Lewis structure.

The structure is ‘resonating’ from one of the structures to the other that the actual structure is a

combination of the structure. For example, the nitrate polyatomic ion (see below) (NO3–) can be drawn as

three different structures, each equally correct (Figure 8):

+

+

Figure 8. Resonance Structures for Nitrate Ion. The first three structures are summarized by the

furthest right structure.

The bond angle between each of the N-O bonds is 120o. As it turns out, the nitrogen-oxygen bond is an

average of the three possibilities: 2/3rd

single N-O bond and 1/3rd

double N=O bond.

H. Hybridization

Bonds are formed by the overlap of electron orbitals. For example, H2 is formed from two hydrogen

atoms by the overlapping of each electron in the adjacent 1s orbital (Figure 5 above). Thus, we would

first believe that can be three identical bonds in a molecule, at the most (from the three p-orbitals).

However, theory must agree with experimental evidence: the bond lengths and angles are identical for

each of the four C-H bonds in methane (CH4). This is explained by bond hybridization orbitals of

equal energy produced by combining two or more orbitals on the same atom. Illustrating with the

methane molecule, each bond between the hydrogen and carbon is composed of 25% character s-orbital

and 75% p-orbital. It is as if all four orbitals (one s- and three p-orbitals) are combined and then

redistributed into four equal parts (Figure 9). This is sp3 hybridization: the orbitals have 1 part s and three

parts p.

H:

1s 1s

H:

1s 1s

H:

1s 1s

H:

1s 1s

C: C:

1s 2s 2p 1s 2sp3

Figure 9. sp3 Hybridization Orbitals in CH4


I. Sigma () and Pi () bonds

Molecular bonds are formed when atomic bonds overlap. Principally, covalent bonds are formed by the

overlap of electrons in the s- and p-orbitals (Figure 10 and Figure 11).

A single covalent bond is a sigma () bond. Sigma

bonds are formed when (a) two s-orbitals overlap, (b)

an s-orbital and a p-orbital overlap, or (c) two p-orbitals

overlap end-to-end (e.g., px + px).

a.

b.

c.

A double covalent bond is comprised of a sigma and a

pi () bond. -Bonds are made when (D) two p-

orbitals overlap laterally (side-by-side). A triple

covalent bond is composed of a -bond and two p-

orbitals overlap, a triple covalent bond is formed.

d.

Figure 10. Schematic representations of molecular bond formation (sigma and pi).

One -bond holds together

the two hydrogen atoms in a

hydrogen molecule.

One -bond and one -bond

holds together the two carbon

atoms in an ethene (ethylene)

molecule.

One -bond and two -bonds

hold together the two nitrogen

atoms in a nitrogen molecule.

Figure 11. Sigma- and pi-bonds form single, double and triple bonds.


J. VSEPR (Valence-Shell Electron-Pair Repulsion Theory)

VSEPR is used to determine the geometry of the molecule. Its assumption is that the geometry of a

molecule results from the pairs of electrons (because they are like charges and repel each other) move as

far apart from each other as they can (Table 8). To apply VSEPR, “A” = the central atom, “B” = number

of atoms that are bonded to that central atom, and “E” = number of lone pairs electrons.

Table 8. Valence-Shell Electron-Pair Repulsion Theory and Molecular Geometry.

Sh

ape

Nam

e

Mo

lecu

lar

Sh

ape

Ato

ms

Bo

nd

ed t

o

Cen

tral

Ato

m (

B)

Lo

ne

Pai

rs

of

Ele

ctro

ns

(E)

Ty

pe

of

Mo

lecu

le

Ex

ample

Lew

is

Str

uct

ure

of

Ex

ample

Co

mm

ents

Linear 2 0 AB2 BeF2

Bent

2 1 AB2E SnCl2

Bent

2 2 AB2E2 H2O

104.5o

between

H-O & H-O

Trigonal-

planar

3 0 AB3 BF3

120o

between

B-F & B-F

Tetrahedral

4 0 AB4 CH4

109.5o

between C-H’s

Trigonal-

pyramidal

3 1 AB3E NH3

107o between

N-H’s

Trigonal-

bipyramidal

5 0 AB5 PCl5

expanded octet

Octahedral

6 0 AB6 SF6

expanded octet


AX AX2 AX3 AX4 AX5 AX6

HCl

Polar

sp3

1/0

Linear

180o

BeF2, CO2

Nonpolar

sp

2/0

Linear

180o

BF3, SO3

Nonpolar

sp2

3/0

Trigonal planar

120o

CH4

Nonpolar

sp3

4/0

Tetrahedral

109.5o

PF5

Nonpolar

sp3d

5/0

Trigonal

bipyramidal

90o, 120

o, 180

o

SF6

Nonpolar

sp3d

2

6/0

Octahedral

90o,180

o

E1

GeF2, SO2

Polar

sp2

2/1

Bent

<120o

NH3

Polar

sp3

3/1

Trigonal

pyramidal

107o

SF4

Polar

sp3d

4/1

See-saw 90o,

120o, 180

o

ClF5

Polar

sp3d

2

5/1

Square

Pyramidal 90o,

180o

Molecular Formula

Polarity Hybridization

[diagram]

Shared/Unshared

(pairs of electrons

around the central

atom)

shape

bond angle(s)

E2

H2O

Polar

sp3

2/2

Bent

104.5o

ClF3

Polar

sp3d

3/2

T-shaped 90o,

180o

XeF4

Nonpolar

sp3d

2

4/2

Square planar

90o, 180

o

Note:

For molecules

with 5 pairs

around the

central atom,

the unshared

pairs are found

on the

equatorial

positions.

E3

Note: Left of

this line has up to an

octet; Right

of this line has an

expanded

octet.

XeF2

Nonpolar

sp3d

2/3

Linear

180o

For molecules

with 6 pairs

around the

central atom,

the unshared

pairs are found

on the axial

positions.

E. Spinelli, RHS


K. Intermolecular Bonding

Group of H-bonds, dipole-dipole interactions, and London forces are called van der Waals forces.

(After Johannes van der Waals’s gas equation).

1. H-Bonding (H and N, O, F) – strongest of intermolecular forces.

2. Dipole-Dipole Interactions (polar molecule – polar molecule)

3. Ion-Dipole Interactions (polar molecules – ions)

4. London Forces (1928; induced dipoles cause nonpolar molecules to be mutually attracted)

III. IONIC BONDING AND IONIC COMPOUNDS

ionic compound compound composed of positive and negative ions that are combined so that the

compound’s total charge is zero

formula unit simplest ratio of atoms in an ionic compound.

Unlike a molecular formula – for which there is a single entity, such as a single H2O – ionic compounds

exist as a large array of ions. There is no single entity of NaCl but a large crystal lattice composed of

many sodium ions and many chloride ions that are present in a 1:1 ratio (e.g., Figure 2).

A. Formation of Ionic Compounds

Molecules are formed by atoms sharing electrons to satisfy the octet rule. Ionic compounds are formed

by losing or gaining electrons, to satisfy the octet rule, and the resulting cations and anions interacting

(Figure 12).

a. b.

Figure 12. Ionic Compound Formation by the Loss/Gain of Electrons. Formation of an ionic compound

results from the loss/gain of electrons and the resulting attraction of the oppositely-charged particles.

Shown above are formation of (a) sodium chloride and (b) calcium chloride.

Ionic compounds are generally organized into a three-dimensional crystal lattice structure. The energy

holding together varies by charge and size of ions: the smaller, more highly charged ions form stronger

lattices.

Lattice energy energy released when one mole of an ionic compound is formed from gaseous ions.

For example, NaCl has more lattice energy than KCl because sodium is a smaller ion than potassium, and

CaF2 has more lattice energy than NaCl because the calcium ion has a much larger charge than sodium

(2+ compared with 1+).

B. Comparison of Ionic and Molecular Compounds

Each ion in an ionic compound is held in place by a large number of strong ionic bonds with other ions

(e.g., each Na+ is bonded to six Cl

– ions). In a molecule, each atom is held in a molecule by strong

intramolecular, covalent bonds. However, the attraction between molecules (intermolecular forces) is

weak compared with either ionic or covalent bonds. This difference in strength – between ionic bonds and

intermolecular forces, accounts for much of the differences between ionic and covalent compounds (Table

9).


Table 9. Comparison of Ionic Compounds and Molecules.

Physical Property Ionic Compounds Molecules

Melting Point High Low

Boiling Point High Low

Brittle Yes No

Hard Yes No

Conducts Electricity molten state or dissolved in water No

C. Monoatomic Ions Monoatomic ions are made up of from a single atom. Metals form cations; nonmetals,

anions. For example, sodium cation, Na+, is formed to in order to be isoelectronic with neon.

Chlorine forms chloride anion for the same reason. NB: the suffix changes only in the anion;

transition metals are assumed to be +2, but have different charges (e.g., iron has a +2 and a

+3). Examples of monatomic ions include Mg2+

, N3–

, and O2–

.

D. Polyatomic Ions

Besides the monoatomic ions (e.g., Na+, Cl

–, O

2–), polyatomic ions are also very common (Table

10). All are very stable – many have resonance structures. When forming or breaking apart

compounds, consider the polyatomic ions as remaining as a single unit. For example,

ammonium nitrate breaks apart in water to form ammonium ion and acetate ion. These ions can

decompose, but strong conditions are required such as strong acid or base hydrolysis,

combustion, or electrolysis.

(NH4)2NO3 OH2

2 NH4+

+ NO32–


Table 10. Common Polyatomic Ions.

CATIONS

1+ 2+

ammonium

4NH dimercury

hydronium1

OH 3

ANIONS

1– 2– 3–

acetate CH3COOH

or

232 OHC carbonate

2

3CO phosphate 3

4PO

sulfate 2

4SO

chlorate

3ClO sulfite 2

3SO

chlorite

2ClO

bicarbonate or

hydrogen carbonate

3HCO thiosulfate

hydrogen sulfate

4HSO

hydroxide OH

hypochlorite ClO

nitrate

3NO

nitrite

2NO

perchlorate

4ClO

permanganate

4MnO

IV. METALLIC BONDING “Sea of Electrons” Model

Electrons are not associated with specific nucleus as are electrons in ionic or covalent bonds. Instead

they act like a mesh surrounding the particles.

Conducts Electricity: Electrons’s freedom of motion results for high electrical conduction

Luster: Electrons contain may orbitals separated by extremely small energy differences

Malleability/Ductility: Bonding is same in all directions so when metal is stressed one plane of

atoms can slide across another without encountering resistance

( break or brittle)

1 only occurs in aqueous solutions


V. NOMENCLATURE - NAMING COMPOUNDS There are three systems of naming chemical compounds:

1. Ionic compounds

2. Molecules

3. Acids

A. Ionic Compounds 1. General

a. Ionic compound composition is typically:

b. a metal and a nonmetal

OR c. a metal and a polyatomic anion

d. one common exception: the cation is the polyatomic ion ammonium (NH4+)

2. Naming: Going from Formula to Name:

a. There are two basic types of ionic compounds:

Binary Compounds = contain metal + nonmetal (e.g., NaCl, MgCl2)

Ternary Compounds = contain metal + polyatomic ion (e.g., Na2SO4)

In either case, the first name is the name of the cation, which is the same name as the atom.

i. Binary Compounds (contain one-atom cation + one-atom anion)

(1) First name is the cation. Second name is the anion but the suffix is changed to –

ide.

e.g., NaCl = sodium chloride

(chlorine, the atom, is changed to the anion, chloride)

e.g., KF = potassium fluoride

(potassium, the atom, is changed to the anion, fluorine)

ii. Ternary Compounds (contain one-atom cation + polyatomic anion)

(1) First name is the cation. Second name is the name of the polyatomic ion (Table

10. Common Polyatomic Ions.)

e.g., MgO = magnesium oxide

e.g., Na2SO4 = sodium sulfate (sulfate, SO42–

, is the polyatomic ion)

e.g., exception:

cation is ammonium (NH4+), but the first name is ammonium.

1. e.g., NH4Cl is ammonium chloride

2. e.g., NH4NO3 is ammonium nitrate

3. Naming: Going from Name to Formula:

a. cation: always goes 1st; anion 2

nd (e.g., Na

+ and Cl– NaCl; oxygen + sodium Na2O)

b. formula:

1. write symbols for ions side by side: Mg2+

O2-

2. cross over charges (not sign) for subscript of other: Mg2 O2

3. write formula: Mg2O2

4. check subscripts for lowest common denominator: Mg2O2 MgO


4. Common Monoatomic Anions:

carbon carbide nitrogen nitride oxygen oxide fluorine fluoride

phosphorous phosphide sulfur sulfide chlorine chloride

arsenic arsenide selenium selenide bromine bromide

iodine iodide

5. Stock System - Used for Transition Metals

1. cations – Some elements have more than one possible charge (e.g., Fe2+

and Fe3+

). All

transition metals can be written in the stock system. Commonly encountered non-

transition metals that have more than one oxidation number are lead (2+ and 4+) and tin

(2+ and 4+).

2. formula: same as for binary ionic compound (q.v., above)

3. naming: include Roman numeral for charge in parentheses

e.g., Fe3+

& O2–

formula: Fe2O3, name: iron(III) oxide

Fe2+

& O2–

formula: FeO, name: iron(II) oxide

6. Ionic Ternary Compounds

a. like ionic binary compounds, cation always goes 1st; anion 2

nd

b. formula:

(1) write symbols for ions side by side: NH4+ SO4

2–

(2) cross over charges (not sign) for subscript of other, with

parenthesis holding together polyatomic ion and subscript going

outside of a parenthesis: (NH4)2(SO4)1

(3) write formula without ‘1’ in subscript (or, then, parenthesis): (NH4)2SO4

(4) (as with binary ionic compounds, check subscripts for lowest common denominator)

c. naming:

(1) cation – name remains the same

(2) anion – name remains the same

(3) e.g., (NH4)2SO4 = ammonium sulfate

B. Molecules

o Molecule two or more atoms, usually nonmetals, covalently bonded without having a net

charge (e.g., N2O5, CO2; NO3 is the molecule nitrogen trioxide but NO3– is the

nitrate ion)

1. Whereas ionic compounds use the stock system, molecules use the prefix naming system (Table

11). The first element, if there’s only one atom in the formula, can have the ‘mono’ dropped.

Otherwise, the prefix remains – even if the second element is only one atom.

Table 11. Molecular Prefix Naming System

# Prefix # Prefix

1 mono- 6 hexa-

2 di- 7 hepta-

3 tri- 8 octa-

4 tetra- 9 nona-

5 penta- 10 deca-

e.g., NO = nitrogen monoxide N2O = dinitrogen monoxide

NO2 = nitrogen dioxide N2O5 = dinitrogen pentoxide


C. Acids

1. Binary Acids = hydrogen + monoatomic anion

a. formula: e.g., HCl, HF, HBr

b. naming:

(1) hydrogen + (base of anion)-ic

e.g., HCl = hydrochloric acid; HI = hydroiodic acid; HF = hydrofluoric acid

2. Ternary Acids = hydrogen + polyatomic anion

a. formula: e.g., H2SO4 (hydrogen & sulfate)

b. naming: use the name of the polyatomic anion EXCEPT that the:

–ate changes to –ic e.g., H2SO4 = sulfuric acid; HClO4 = perchloric acid

–ite changes to –ous e.g., H2SO3 = sulfurous acid; HClO = hypochlorous acid

3. Some Common Acids (Table 12.)

Name Formula Common Use Anion

hydrochloric acid HCl stomach acid chloride (Cl–)

nitric acid HNO3 (industrial uses) nitrate (NO3–)

sulfuric acid H2SO4 car battery acid sulfate (SO42–

)

phosphoric acid H3PO4 colas (e.g., Pepsi, Coke) phosphate (PO43–

)


NOTES:

Documents

CHEMICAL BONDING NOMENCLATURE€¦ · I. CHEMICAL BONDING: IONIC AND COVALENT BONDS A. Types of Chemical Bonds II. MOLECULES A. Electron Configuration and the Octet Rule B. Orbital