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1.1
Covalent bonding
A covalent bond is formed when the electron clouds of two nonmetal atoms overlap (1). Where the
clouds overlap they are thicker, and their electric charge is stronger. Both nuclei (2) feel a force of
attraction towards this thick electron cloud, and so the two atoms are held together. A group of atoms
joined in this way is known as a molecule (3). The most common type of covalent bond is the single
bond (4), the sharing of only one pair of electrons between two atoms. Sharing two pairs is called a
double bond (5). Sharing three pairs is called a triple bond (6). Non-bonded electron pairs are also
called lone pairs (7).
+
F
O= O
NN
+
F
F
F
Nonpolar covalent bond
In nonpolar covalent bonding the electronegativity difference between the bonded atoms is small or
non-existent. These covalent bonds are called nonpolar covalent bonds (8) because the electrons
shared by the adjacent atoms in the bonds are shared equally. The consequence of this equal sharing
of electrons is that there is no charge separation (dipole moment).
Polar covalent bond
In a water molecule the atoms are linked by polar covalent bonds (9). In a polar covalent bond, the
electrons shared by the atoms spend a greater amount of time, on the average, closer to the Oxygen
nucleus than the Hydrogen nucleus. This is because of the great electronegativity difference between
the Hydrogen atom and the Oxygen atom. The result of this pattern of unequal electron association is a
charge separation in the molecule, where one part of the molecule has a partial negative charge (10)
and the other part has a partial positive charge (11). The dipole moment (12) of the molecule is > 0.
H
O
C
H
H
H
H
H
NONPOLAR
1.2
Ionic bonding
An ionic bond is a type of chemical bond that can often form between metal and nonmetal ions. In short, it is
a bond formed by the attraction between two oppositely charged ions.
The metal donates one or more electrons (13), forming a positively charged cation with a stable electron
configuration. These electrons then enter the nonmetal (14), causing it to form a negatively charged anion
which also has a stable electron configuration (15). The electrostatic attraction between the oppositely
charged ions causes them to come together and form a bond. An ionic bond (16) is not uni-directional (17); it
exists in all directions. In a crystal lattice (18), ionic bonding extends throughout the entire structure.
Na
Cl
+ e


Na+
Cl
+ e
d+
crystall lattice
1.3
Metallic bonding
In metals, the atoms lose their outer electrons (valence electrons) to form metal cations. The electrons from
all the metal atoms form a "sea" of electrons (19) that can flow around these metal cations (20). These
electrons are often described as delocalized electrons. As the metal cations and the electrons are oppositely
charged, they will be attracted to each other, and also to other metal cations. These electrostatic forces are
called metallic bonds (21). Metallic bonds are not uni-directional (22).
2 Intermolecular forces
Two factors determine whether a substance is a solid, a liquid, or a gas:
- The kinetic energies of the particles (atoms, molecules) that make up a substance.
Kinetic energy tends to keep the particles moving apart.
- The attractive intermolecular forces between particles that tend to draw the particles together.
If the average kinetic energy is greater than the attractive forces between the particles, a substance will not
condense to form a liquid or a solid. If the kinetic energy is less than the attractive forces, a liquid or solid will form.
(23)
solid substance
Ekin <<
attractive forces
liquid substance
Ekin <
attractive forces
gaseous substance
Ekin >
attractive forces
2.1
London dispersion forces (Van der Waals forces)
The London dispersion force (24) is the weakest intermolecular force. The London dispersion force is a
temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make
the atoms form temporary dipoles. This force is sometimes called an induced dipole-induced dipole attraction.
Because of the constant motion of the electrons, an atom or molecule can develop a temporary dipole (25)
when its electrons are distributed unsymmetrically about the nucleus. A second atom or molecule, in turn, can
be distorted by the appearance of the dipole in the first atom or molecule (because electrons repel one
another) which leads to an electrostatic attraction between the two atoms or molecules.
Dispersion forces are present between all molecules, whether they are polar or nonpolar.
Br Br
Br Br
2.2
Dipole dipole forces
Dipole-dipole forces (26) are attractive forces between the positive end of one polar molecule and the
negative end of another polar molecule (permanent dipole).
d+ dH Br
2.3
d+ dH Br
Hydrogen bonds
Hydrogen bonding (27) is a special type of dipole-dipole attraction between molecules. It results from
the attractive force between a hydrogen atom covalently bonded to a very electronegative atom such as
a N, O, or F atom and another very electronegative atom. Because of the difference in electronegativity,
the H atom bears a large partial positive charge and the N, O or F atom bears a large partial negative
charge. A H atom in one molecule is electrostatically attracted to the N, O, or F atom in another
molecule.
d-
H
O
H
d+
d-
H
O
H
2.4
Ion dipole forces
An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion
and a neutral molecule that has a dipole.
A positive ion (cation) attracts the partially negative end of a neutral polar molecule (28).
A negative ion (anion) attracts the partially positive end of a neutral polar molecule (29).
Ion-dipole attractions become stronger as either the charge on the ion increases, or as the radius of
the ion decreases.
d+
d–
d–
d+
d+
d+
d–
d+
d+
d–
2.5
Solubility rules
Solubility of molecular compounds in water
Molecular compounds are highly soluble in water, if hydrogen bonds are formed between their molecules
and H2O and if their nonpolar part does not contain more than 3 C-atoms.
Highly soluble molecular compound (30):
CH3OH
Poorly soluble molecular compound (31):
CH3CH2CH3
Solubility of salts in water
Two forces determine the extent to which solution will occur:
-
Force of attraction between H2O molecules and the ions of the solid
If this force is the predominant factor, then the compound may be highly soluble in water.
-
Force of attraction between oppositely charged ions
If this force is a major factor, then water solubility may be very low.
Salts are hardly soluble in water, if the charge of the cation and the anion is 2 or higher and Q1  Q2  4 ,
respectively.
NaCl
Highly soluble salt (32):…………………………….
Al2O3
Hardly soluble salt (33):…………………………….
1. Which of the following pairs of atoms are least likely to form an ionic compound?
a) Ni, O
b) Na, F
c) Cu, Cl
d) Li, Mg
e) Li, F
2. The bond in nitrogen (N2) is a:
a) double bond
b) single bond
c) triple bond
d) lone pair
e) none of the above
3. Which of the following formulas are incorrect?
a) LiCl
b) MgO
c) Na2O
d) CO2
e) none of these
4. Which of the following bonds is most polar?
a) H-F
b) H-Cl
c) H-H
d) F-F
5. Which of the following molecules requires two resonance hybrid Lewis structures (mesomere Grenzformeln) to
account for the bonding?
a) SO2
b) CO2
c) O2
d) CH4
e) C2H4
6. Which of the following pairs of atoms are most likely to form a covalent compound?
a) Ni, O
b) Na, F
c) Cu, Cl
d) C, O
e) Li, F
7. The number of valence shell electrons in chlorine is:
a) 17
b) 2
c) 5
d) 7
e) 4
8. The bond in oxygen (O2) is a:
a) double bond
b) single bond
c) triple bond
d) lone pair
e) none of the above
9. Given the following table of bond enthalpies:
average bond
bond enthalpy (kJ/mol)
C-H
413
C-C
348
C-N
293
N-H
391
H-H
436
N-N
163
C=C
614
C N
891
Calculate the H for the following reaction in the gas phase: HCN(g) + 2 H2(g)
a) -736 kJ
b) -551 kJ
c) -138 kJ
d) +230 kJ
e)+275 kJ
H
H
C
N
+
2H
H
H
C
H
413 + 891 +
2·436
H
N
H
- 3·413 -293 - 2·391= -138 kJ
CH3NH2(g)
12 Paare = 24 Elektronen
24 = 5 + 7+2x
x=6
10. Which of the following could be the atom X
in the following neutral (uncharged) molecule?
a) H
b) F
c) C
d) N
e) O
19 Paare = 38 Elektronen
38 = 4·6 + 2x + 2
x=6
11. Atom X in the structure below is most likely:
a) Cl
b) S
c) P
d) Si
e) Al
12. Which molecule has one pair of nonbonding electrons on the central atom?
a) PCl3
1
b) CO2
0
c) SO3
d) BF3
0
0
e) SF2
2
lone pairs
13. Which of the following compounds has the largest lattice energy in the crystal state?
a) LiCl
Li+ Cl-
b) NaBr
Na+ Br-
c) MgI2
Mg2+ 2I-
d) CaO
Ca2+ O2-
14. CO2 is a nonpolar molecule (µ=0) whereas SO2 is polar (µ > 0). This difference is due to the fact that:
a) C and O have approximately the same electronegativity while S and O have different
electronegativities.
b) CO2 has an even number of double bonds whereas SO2 has an odd number of double bonds.
c) C and O are in different groups whereas S and O are in the same group.
d) the C-O bond is nonpolar while the S-O bond is polar.
e) CO2 is linear whereas SO2 is not linear.
15. Which of the following is a non-polar covalent compound?
a) CH4
b) HCN
c) CH3CN
d) HCl
e) NaCl
16. The type(s) of bonding present in a sample of sodium nitrate, NaNO 3, are:
a) covalent bonds only
b) ionic bonds only
d) covalent and ionic bonds
e) ionic and metallic bonds
Na+,
NO3
O
-
+
N
O
O
c) metallic bonds only
17. Which of the following species is best described by a Lewis structure involving resonance?
2-
3-
-
a) CO3
b) PO4
d) All of the above are best described by resonance
c ) HSO4 (H bonded to O)
e) None of (a)-(c)
18. Which of the following molecules has a non-zero dipole moment?
a) Cl2O
b) I2
c) BF3
d) CO2
e) CF4
19. Which of the following compounds (or ions) contains an atom that does NOT follow the octet rule?
a) BrF3
b) CBr4
c) NO
+
-
d) BF4
e) NCl3
F
3·7 + 7
= 28El.
= 14 P.
Br
F
F
20. What is atom X in the Lewis structure of the following dianion?
a )Xe
b) Cl c) S
d) N
e) C
16 Paare = 32 Elektronen
32 = 4·6 + x + 2
x=6
21. Based on intermolecular interactions, which of the following should have the highest boiling point?
a) CH4
b) CHCl3
c) H2S
d) CH3OH
22. Hydrogen bonding is unimportant in
a) ice formation
b) the liquid properties of water
d) liquid CH4
e) liquid HF
c) DNA structure
23. Which one of the following statements is incorrect?
a) London dispersion forces are the weakest type of intermolecular interactions.
b) The strong intermolecular attractions in H2O result from hydrogen bonding.
c) The boiling point of H2S is lower than H2O.
d) The boiling point of non-polar substances tends to decrease with increasing molecular weight.
24. Which of the species below does not exhibit hydrogen bonding?
a) C2H6
b) NH3
c) HF
d) H2O
e) C2H5OH
25. What type of attractive forces are being overcome when liquid oxygen boils at 90 K?
a) ionic bonds
b) dipole-dipole forces
d) London dispersion forces
c) covalent bonds
e) hydrogen bonds
26. Covalent bonds are the only type of bond or intermolecular forces in:
a) KF(s)
b) CO2(s)
c) H2O(s)
d) NH4NO2(s)
o
e) C (diamond)
27. A white substance melts with some decomposition at 730 C. As a solid, it is a nonconductor of electricity but it
dissolves in water to form a conducting solution. The white substance is:
a) a covalent network solid
d) a metallic solid
b) an ionic solid
e) solid Ar
c) molecular solid
The next two questions refer to 4 liquids, all of which have the molecular formula C 4H10O and which
o
o
o
o
have the following boiling points: A (117.5 C); B (82.9 C); C (34.6 C); D (99.5 C).
28. The intermolecular attractive forces of these are likely to be ordered:
a) A > D > B > C
b) B > D > A > C
c) C > B > D > A
d) D > C > B > A
29. Given the structural isomers below, the one most likely to be liquid C is:
30. Which of the following compounds feature intermolecular bonding in the liquid or solid state?
a) argon
atoms
b) diamond
c) water
d) sodium chloride
covalent network molecules
e) none of these
salt
31. Which force makes the most important contribution to the lattice energy of solid argon?
a) metallic bonding b) hydrogen bonding
c) ionic bonding
d)covalent bonding
e) van der Waals forces
32. What happens to molecules in a liquid when the liquid is heated and vapor starts to form?
a) the intramolecular forces between liquid molecules are disrupted
b) the intermolecular forces between liquid molecules are disrupted
c) kinetic energy is removed from the system
Chemical bonding and intermolecular forces
d) London forces become stronger
Page 1
33. Which of the following statements best explains why diamond (carbon) has a higher melting point
than dry ice (CO2)?
a) because carbon atoms in diamond are held together by a network of covalent bonds
b) because CO2 molecules are held together by dipole - dipole forces
c) the electrostatic attractions in CO2 are not as strong as those in diamond (carbon)
c) the dispersion forces in diamond are greater than in CO2
d) all of the above
Hauptvalenzen
Ionenbindung
(allseitig)
Atombindung
(gerichtet)
O
H
Metallbindung
(allseitig)
H
Molekül
Anion
Elektronengas
Atomrumpf
positiv geladen
Kation
Atomgitter
Hauptvalenzen halten das Gitter zusammen
 keine Nebenvalenzen
Hauptvalenzen und Nebenvalenzen
Ionenbindung
(allseitig)
Atombindung
Metallbindung
(allseitig)
(gerichtet)
H
O
H
Molekül
H
Anion
O
H
Elektronengas
Atomrumpf
positiv geladen
Kation
Atomgitter
Hauptvalenzen halten die Moleküle zusammen
Nebenvalenzen zwischen den Molekülen halten das Gitter zusammen
Nebenvalenzen
Van der WaalsKräfte
(allseitig)
Br Br
Wasserstoffbrücken
(gerichtet)
Dipol-DipolKräfte
d+ dH Br
d+ dH Br
Ion-DipolKräfte
d-
H
O
H
H
d+
Br Br
H
H
O
Ca
H
H
d+
O
2+
O
d-
H
H
O
H
O
H
H