chapters 7 8: chemical bonding nomenclature€¦ · a. electron configuration and the octet rule...
TRANSCRIPT
CHAPTERS 7 & 8:
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 07-08: CHEMICAL BONDING 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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 2
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.
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 3
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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 2
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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 3
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).
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 4
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.
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 5
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
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
2. Total number of valence electrons:
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:
a. calculate pairs of electrons remaining:
(total pairs e’s) – (pair e’s used) = pairs
available
(8 pairs) – (2 pairs used) = 6 pair 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
c. If central atoms does not have octet, convert
terminal’s lone pairs into a double/triple bond to
central atom
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 6
Steps c. Example #3: PO43–
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:
evalenceechnegativethefromelectrons
atomO
evalenceatomO
atomP
evalenceatomP
__32)arg_____3(
)_
__6*
__4()
_
__5*
__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
(16 pairs) – (4 pairs used) = 12 pair 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-.
c. If central atoms does not have octet, convert
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)
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 7
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.
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 8
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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 9
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.
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 10
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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 11
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, 120o, 180o
SF6
Nonpolar
sp3d2
6/0
Octahedral
90o,180o
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, 180o
ClF5
Polar
sp3d2
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
sp3d2
4/2
Square planar
90o, 180o
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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 12
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).
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 13
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. 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 2
2Hg
hydronium1 OH 3
ANIONS
1– 2– 3–
acetate2
232 OHC carbonate 2
3CO phosphate 3
4PO
bromate
3BrO chromate 2
4CrO arsenate 3
4AsO
chlorate
3ClO dichromate 2
72OCr
chlorite
2ClO hydrogen phosphate 2
4HPO
cyanide CN oxalate 2
42OC
dihydrogen phosphate
42 POH peroxide 2
2O
hydrogen carbonate3
3HCO sulfate 2
4SO
hydrogen sulfate
4HSO sulfite 2
3SO
hydroxide OH thiosulfate
2
32OS
hypochlorite ClO
iodate
3IO
nitrate
3NO
nitrite
2NO
perchlorate
4ClO
permanganate
4MnO
1 only occurs in aqueous solutions 2 acetate, in the old system, is CH3COO– 3 also called ‘bicarbonate’ in the older naming system
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 14
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)
V. COVALENT-NETWORK Many atoms bonded by covalent bonds are not molecules but continuous three-dimensional networks
(like ionic compounds, sort of).
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 15
VI. 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 2nd (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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 16
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 2nd
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
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 17
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–
)
CHEMISTRY SG 07-08: CHEMICAL BONDING p. 18
NOTES: