2009 Chemistry I: Modern Chemistry Holt, Rinehart, & Winston
Modified Learning: Chapter 6 – Chemical Bonding, pp 175 – 207
Vocabulary
1 bond energy; 2 chemical bond; 3 chemical formula; 4 covalent bonding; 5 dipole; 6 ductility;
7 electron-dot notation; 8 formula unit; 9 hybridization; 10 hybrid orbitals; 11 hydrogen bonding;
12 ionic bonding; 13 ionic compound; 14 lattice energy; 15 Lewis structure; 16 London dispersion
forces; 17 malleability; 18 metallic bonding; 19 molecular compound; 20 molecular formula;
21 molecule; 22 multiple bonds; 23 nonpolar-covalent bond; 24 polar; 25 polar-covalent bond;
26 polyatomic ion; 27 resonance; 28 single bond; 29 structural formula; 30 VSEPR theory
1 energy required to break a chemical bond & form neutral isolated atoms
2 mutual electrical attraction between nuclei & valence electrons of different atoms binding them
together
3 atomic symbols & numeric subscripts indicating relative numbers of atoms of each kind in compounds
4 chemical bonding resulting from the sharing of electron pairs between two atoms
5 force created by equal but opposite charges that are separated by a short distance
6 ability of substance to be drawn, pulled, or extruded through a small opening to produce a wire
7 electron configuration notation in which only valence electrons of a particular element are shown,
indicated by dots placed around the element’s symbol
8 simplest collection of atoms from which an ionic compound’s formula can be established___________
9 the mixing of > 2 atomic orbitals of similar energies on the same atom to produce new hybrid atomic
orbitals of equal energies
10 orbitals of equal energy produced by the combination of > 2 orbitals on the same atom
11 intermolecular forces in which a H atom bonded to a highly electronegative atom is attracted to an
unshared pair of electrons of an electronegative atom in a nearby molecule
12 chemical bonding resulting from the electrical attraction between anions & cations
13 composition of positive & negative ions that are combined so that numbers of (+) & (-) charges equal
14 energy released when 1 mole of an ionic crystalline compound is formed from gaseous ions
15 formulas in which atomic symbols represent nuclei & inner-shell electrons & dot-pairs or dashes
between atomic symbols represent electron pairs in covalent bonds, whereas lone ones represent
unshared electrons
16 intermolecular attractions from constant motion of electrons & creation of instantaneous dipoles____
17 ability of substance to be hammered or beaten into thin sheets
18 chemical bonding resulting from attraction between metal atoms & surrounding sea of electrons
19 chemical compound whose simplest units are molecules
20 indication of the types & numbers of atoms combined in a single molecule of a molecular cpd
21 neutral group of atoms held together by covalent bonds
22 double &/ or triple bonds
23 covalent bond in which bonding electrons are shared equally by the bonded atoms_____________
24 bonds or molecules in which there is an unequal distribution of electrons/ charge
25 covalent bond in which bonded atoms have an unequal attraction for the shared electrons
26 a charged group of covalently bonded atoms
27 bonding in molecules or ions that cannot be correctly represented by a single Lewis structure
28 covalent bond in which 1 pair of electrons is shared between 2 atoms
29 representation indicating the kind, number, arrangement & bonds but not unshared pairs of atoms
in a molecule
30 the repulsion between sets of valence-level electrons surrounding an atom causes these sets to be
oriented as far apart as possible
Main Ideas
I. Types of Chemical Bonding
A. Ionic
B. Covalent
C. Metallic
II. Degree of Types
A. Difference in electronegativities of bonded atoms (Fig 2, p176 Modern Chemistry)
a. 0 - 0.3 = nonpolar covalent = 0% - 5% ionic character
b. 0.3 – 1.0 = moderately polar covalent
c. 1.0 – 1.7 = very polar covalent=>5% -50% ionic character of 0.3-1.7 difference
d. > 1.7 = ionic
*** Figure 20 on p161 of Modern Chemistry lists electronegativities for all elements.
III. Covalent Bonds
A. Single covalent bonds
1. A bond in which 2 atoms share a pair of electrons
2. Is represented by a dash between the atoms in structural formulas [H—H ]
3. Forms “molecules” rather than “compounds”
a. Molecular formulas reflect actual numbers of atoms in each molecule
b. Attractive forces dominate, and then balance repulsive forces
4. Forms between > 2 nonmetals to achieve noble gas configuration
a. Attraction corresponds to decrease in potential energy of the atoms
b. Diatomic halogens [F—F, Cl—Cl, Br – Br, etc… ]
c. Water, ammonia, carbon tetrachloride, etc…
5. Forms in carbon-based molecules
a. Carbon has 4 valence electrons, 1s2 2s2 2px1 2py1 2pz0  1s2 2s1 2px1 2py1 2pz1
b. Can combine with cations, such as H+, or anions, such as halogens, oxides, etc..
B. Lewis Dot Structures
1. Add up valence electrons for each atom in the molecule
2. Write the symbols for the atoms in the molecule
a. The central atom is often the 1st atom in the formula
3. Draw a dash between each pair of electrons covalently bonded
4. For each dash subtract 2 electrons from the total # of electrons in step 1
5. Draw remaining electrons as dots around the atom (most have 8 valence electrons)
6. If there are not enough electrons for full octet, shift unbonded electrons as needed or
convert single bonds to double or triple to give each atom its noble gas config.
C. Double & Triple Covalent Bonds
1. Double covalent bonds share 4 electrons and triple covalent bonds share 6 electrons.
2. Exceptions to the octet rule
a. O2 (with 6 valence electrons each) predominantly has single bond
b. NO2 will contain an unshared pair of electrons O—N – O & O – N – O
c. BF3 has B deficient by 2 e—
d. PCl5 has P with 2 extra e—
e. SF6 has S with 4 extra e—
3. Diamagnetic molecules have all e-- paired & are weakly repelled by magnetic fields
4. Paramagnetic ones have > 1unpaired e-- & have strong attraction to magnetic fields
Paramagnetic compounds appear to have a greater mass in magnetic fields
D. Resonance Structures
1. Single structure that is the average of two structures; cannot be represented well by a
single Lewis structure
a. Ozone (O3.) – has 2 types of O – O bonds, one single & one double
O=O–OO–O=O
E. Covalent-Network Bonding
1. Many covalently bonded compounds to not contain individual molecules
2. Exist as continuous, 3-D networks of bonded atoms
IV. Ionic Bonding & Ionic Compounds
A. Electron configuration
1. Valence electrons = the e- in the highest occupied energy level of an element’s atoms
a. Typically the only electrons involved in formation of chemical bonds.
b. Only these e- are shown in electron-dot structures
2. Group Number is equivalent to the number of valence electrons
a. All in a given group then will have the same # of e- dots in their structures.
3. Octet rule
a. Developed by Gilbert Lewis, 1916
b. Atoms tend to achieve e- configuration of noble gases, i.e., set of eight (octet).
c. Of course, there are some exceptions.
+
Normal: Na  Na + e-; Mg  Mg2+ + 2e-; Not: Fe  Fe2+ + 2e- (ferrous); Fe  Fe3+ + 3e- (ferric)
EXCEPTIONS: Ag (1s22s22p63s23p63d104s24p64d105s1) would have to lose 11 e- to be like Kr and have to
gain 7 e- to be like Xe. Losing only the 5s1 e- gives a full outer energy level with 18 e-, and is thus
energetically stable/ favorable. This gives a “pseudo noble-gas configuration”.
Other elements that do this too are Cu+, Au+, Cd2+, & Hg2+.
Anions:
Atoms of nonmetallic elements GAIN electrons to complete the octet rather than lose any. Cl + e-  ClHalide ions are those ions formed when any of the halogens (Group VIIA) gain 1 electron: fluoride F -,
chloride Cl -, bromide Br -, iodide I -: Other common anions: hydroxide OH-, hypochlorite ClO-,
nitrate NO3- , acetate C2H3O2-, bicarbonate HCO3-, oxide O2-, sulfide S2-, sulfate SO42-, phosphate PO43nitride N3-, phosphide P3Ionic Compounds
Electrically neutral groups of ions joined by electrostatic forces
Ionic bonds = electrostatic forces of attraction that bind oppositely charged ions
Na+ + Cl -  NaCl
3 Br - + Al 3+  AlBr3
0
Properties of ionic compounds: At Rm T, most are crystalline solids & have high melting temperatures.
This is due to the arrangement of atoms in the crystal. Na+ is surrounded by 6 Cl- & each Cl- has 6 Na+
***The melting point, boiling point, & hardness of a compound depend on how strongly basic units are
attracted to each other. Although covalent bonds are very strong, there are relatively weak forces
between molecules. Ionic crystals are much more brittle than molecular compounds because the balance
of attractive & repulsive forces in a crystal are so critically arranged that even the slightest shift of one
row or layer of ions disrupts the entire structure. Crystals are therefore lousy electrical conductors when
in their solid form and must be dissolved in water before any conductance occurs. Some are not even
soluble because of their rigid/ complete attraction between ions of the crystal.
The coordination number = # of ions of opposite charge that surround the ion in a crystal
NaCl has a coordination # of 6. CsCl & TiO2 have a coordination # of 8. [TiO2 = rutile]
Internal structures of crystals are determined by x-ray diffraction crystallography.
Metallic Bonds & metallic properties
Metals are composed of tightly packed cations (+ charge) rather than neutral atoms.
These cations are surrounded by mobile valence electrons that can freely drift within the metal.
Metallic bonds = attraction of mobile valence e- for the positively charged metal ions.
Metals are good conductors of electricity. They are ductile (drawn into wires) & malleable (capable of
being hammered or forced into shapes). Under pressure, the metal cations easily slide past one another
like ball bearings immersed in oil. This is in contrast to ionic crystals which would shatter if struck
because of ions of like charge repelling themselves.
Metals also exhibit luster, or shininess, which is due to their mobile sea of electrons. Metals have
many orbitals separated by extremely small energy differences and therefore, can absorb a wide range of
light frequencies. This allows sufficient excitation of electrons to higher levels. As the electrons return
to their native energy state, the absorbed energy is re-radiated (or reflected)… and we call it ‘shiny’.
Metallic bond strength is measured as a function of the metals’ enthalpy of vaporization, or the amount
of energy / heat required to vaporize the metal.
Crystalline structure of metals
Metals that contain just one kind of atom are among the simplest forms of all crystalline solids.
Arrangement: very compact & orderly patterns
body-centered cubic: every internal atom has 8 neighbors <Na, K, Fe, Cr, W-tungsten>
face-centered cubic: every internal atom has 12 neighbors <Cu, Ag, Au, Al, Pb-lead>
hexagonal close-packed: every atom also has 12 neighbors < Mg, Zn, Cd>
Alloys = mixtures composed of > 2 elements, at least one of which is a metal. Product is superior to its
components. Sterling silver = 92.5% Ag + 7.5% Cu; Bronze = 7:1 of Cu:Sn; Cast iron = 96% Fe + 4%C
Substitutional alloys = atoms of components are ~ same size, so they can replace one another
Interstitial alloys = atoms of smaller atoms fit in the interstitial spaces between larger atoms
V. II. Bonding Theories
A. Molecular Orbitals
1. Based on the quantum mechanical model of bonding
2. When 2 atoms combine, their atomic orbitals overlap
a. This creates molecular orbitals that apply to the entire molecule
b. Each molecular orbital is filled if it contains 2 e—
c. The # of molecular orbitals = # of overlapping atomic orbitals
3. When 2 atomic orbitals overlap, 2 molecular orbitals are created.
a. One is the bonding orbital with an energy that is < the parent atomic orbital.
The e— density between the nuclei is high.
Sigma bond forms when mol. orbital is symmetrical along axis between atoms
b. One is the antibonding orbital with energy > the parent atomic orbital.
The e— density between the nuclei is low.
Balance between attraction & repulsion favors repulsion, e.g. He2 won’t work
4. Overlapping orbitals can be s or p or combinations of s & p
a. H2 is formed with overlapping s orbitals
b. F2 is formed with overlapping p orbitals
5. Overlapping orbitals can be end-to-end in sigma () bonds or side-by-side pi ( bonds
a.  bonds are weaker than  bonds
B. VSEPR Theory
1. Valence Shell Electron-Pair Repulsion theory
2. Describes molecules’ molecular shape
a. Molecular shape adjusts so valence e— pairs are maximally separated
CH4 (or C with any 4 atoms) forms a tetrahedron with 109.5o angles
NH3 has an unshared pair of e—, so it forms pyramid with 107o angles
H2O has 2 unshared e— pairs bending down the O-H bonds to give 105o angles
CO2 has 2 double bonds & 2 unshared e— pairs on each O, so it’s linear (180o)
C. Hybrid Orbitals
1. Describes molecules’ bonding and shape
2. Multiple atomic orbitals overlap to form same # of equivalent hybrid orbitals
a. Single bonds (e.g. CH4)
1 s orbital & 3 p orbitals of C overlap the 1 s orbital of H (sp3)- tetrahedral
b. Double bonds (e.g. C2H4)
One 2s & two 2p orbitals of C overlap (sp2) – trigonal planar
2 sp2 orbitals of each C overlap the 4 H’s 1 s orbitals in a total of 4  bonds
The 3rd sp2 orbital of each C overlap to form a C—C -bonding orbital
The non-hybridized 2p C orbitals overlap side-by-side in a -bonding orbital
c. Triple bonds (e.g. C2H2)
2s orbital of C mixes with 1 of the 3 2p orbitals to give 1 sp orbital for each C
A C—C bonding molecular orbital forms.
The other 2sp C orbital overlaps with the 1s H orbitals.
2 C—H bonding molecular orbital forms.
The remaining pair of C’s p atomic orbitals overlaps side-by-side.
2 C—C -bonding molecular orbital forms.
VI. Polar Bonds and Molecules
A. Bond Polarity
1. The shared electrons in covalent bonds are pulled between the nuclei of the atoms.
a. If atoms pull equally on the electrons, the bond is nonpolar.
This is true for all elemental diatomic molecules (H2, O2, N2, F2, Cl2…)
b. If one atom is >> electronegative, it will pull more e— toward it, & bond is polar.
Electronegativity tends to increase across the period to the right & up the group.
B. Polar Molecules
1. Presence of polar bonds in a molecule often (but not always) makes the molecule polar.
2. Dipoles are molecules with polar ends, one slightly negative, the other slightly positive.
3. The orientation of the polar bonds & shape of the molecule determine the polarity.
a. H2O is bent planar with O having (-) end and H having slightly (+) ends.
b. CO2 is linear & bond polarities cancel because they go in opposite directions.
C. Attractions between Molecules
1. van der Waals forces = Weakest attractions
a. Dispersion forces
Caused by the motion of electrons
Typically increase with increasing # of electrons
Major attraction for diatomic halogen molecules
Accounts for different states of matter for halogens
b. Dipole interactions
Attraction between polar molecules (e.g. H2O to H2O)
2. Hydrogen bonds
a. Occurs when a H covalently bonded to a very electronegative atom is also
weakly bonded to another electronegative atom’s unshared pair of e—
b. Often seen with H bonded to F, O, or N, leaving H very electron-deficient
H-bonds are ~ 5% of the strength of average covalent bond.
c. May occur within the same molecule or between nearby molecule
3. London Dispersion Forces
a. Even noble-gas atoms and nonpolar molecules have weak intermolecular
attraction. These are the ONLY forces between such molecules/atoms.
b. Caused by motion of electrons and creation of instantaneous dipoles
D. Intermolecular Attractions & Molecular Properties
1. Physical properties
a. Depends on the type of bonding displayed, i.e. ionic vs. covalent
b. Ionic: mostly solid; good conductors; soluble in H2O; melting points > 300 oC
c. Covalent: solid, liquid or gas; poor conductors; variable solubility; low m.p.
Some have high m.p. esp. “network solids/ crystals” like diamonds & SiC.