CHEMISTRY FOUNDATIONS
Chemistry is the study of matter and its transformations. More specifically, it looks at how some
materials interact with other materials.
Matter can be defined as a substance that has:
•
Mass
•
Occupies space (i.e., it has volume)
Mass is a measure of the physical amount of a substance. Mass is not to be confused with weight
although the two terms are sometimes used interchangeably in everyday life. Weight is the action of the
acceleration due to gravity on mass; as such, it is a force. For example, a man with a mass of 50 kg on
earth will still have a mass of 50 kg on the moon, but will weigh about 1/6 less on the moon than on the
earth. This is because the moon’s gravity is about one-sixth that of the earth.
Matter can be found in three basic forms (referred to as the states of matter):
1. Solid – the material has a definite size and shape
2. Liquid – the material has a definite size, but takes on the shape of the container
3. Gas – the material has no definite size or shape; it expands (or contracts) to fill the container
1. PHYSICAL AND CHEMICAL PROPERTIES AND CHANGES
The physical properties of a substance are the physical characteristics by which a substance may be
recognised. These properties include for example, colour, odour, state (solid, liquid, gas etc.), volume,
density, melting point and boiling point (substances have many other physical characteristics).
The chemical properties of a substance tell us how the substance combines with other substances to
form new substances (having different physical properties). For example, a chemical property of
hydrogen is that under certain well-defined conditions, it combines with oxygen to from water. Water
has completely different properties from the substances of which it is formed (hydrogen and oxygen).
Exercise:
Classify the following properties as physical or chemical.
1. Gold is yellow.
2. Gold resists a change in identity.
3. Iron rusts when exposed to air and water.
4. A 1 kilogram block of lead occupies less space than 1 kilogram of feathers.
5. Ammonia smells differently from hydrogen sulphide.
Physical changes occur when an object is subjected to a relatively mild physical stress, e.g., heating a
block of ice (solid water) may result in the formation of liquid water. Physical changes may be
recognised because they are often reversible. For example, when ice is warmed, it eventually melts to
form liquid water. If the liquid water is sufficiently cooled, it becomes ice once more.
Chemical changes also occur when an object is subjected to physical stress. The difference between
physical changes and chemical changes is that in chemical changes, new substances are formed; the new
substances do not have the same properties as the original substances. Unlike physical changes,
chemical changes are not often easily reversible. For example, if a mixture of hydrogen gas and oxygen
1
gas is heated strongly, water will be produced. However cooling back the water will not regenerate the
original hydrogen gas and oxygen gas. When two substances are mixed, one or more of the following
observations may signal that a chemical change has occurred:
1. A solid is formed on mixing two liquids.
2. A gas is formed on mixing two substances.
3. A colour change has occurred on mixing two substances.
4. The temperature of the mixture changes (the mixture gets hotter or colder).
2. ATOMS
All matter is composed of atoms. Atoms may vary in mass, size and in the way they interact with other
atoms but they all share some common characteristics.
For the purposes of chemistry, all atoms are made up of three fundamental particles.
1. The electron (e–) is a subatomic particle (i.e., smaller than an atom) that has a negative charge.
2. The proton (p+) is a subatomic particle that has a positive charge equal in magnitude but
opposite in sign to the electron.
3. The neutron (n0) is a subatomic particle that has no charge.
Each atom has a specific number of protons, electrons and neutrons.
The structure of all atoms is as follows.
All of the protons and neutrons exist in a tightly packed region at the centre of the atom called
the nucleus. The electrons are in constant motion outside the nucleus; they move about the
nucleus at random but do occupy preferred regions in space.
The atom is mostly empty space. For example, the nucleus of the smallest atom, hydrogen, is
about 10–15 m. To understand the scale of an atom, suppose the hydrogen atom nucleus is
expanded to the size of a tennis ball and centered at Dawson College; then the electron of the
hydrogen atom will be, on average, about half a kilometre further west than Decarie Boulevard.
The point is that the electron is rather far from its nucleus.
The nucleus consists
of all of the protons
and neutrons
The electrons are
moving around in
the space outside of
the nucleus.
In a free atom the number of protons equals the number of electrons.
The diverse character of different atoms is due to the number of protons that are packed into the small
dense nucleus together with the number and organization of the electrons in the volume outside the
nucleus.
2
2.1. Atomic Mass
Electrons are so tiny that they weigh almost nothing (9.109×10–31 kg); protons and neutrons are nearly
2000-times heavier. Almost all of the mass of an atom therefor comes from the protons and neutrons,
which reside in the nucleus. However, even protons and neutrons are too small to be conveniently
weighed in kilograms. Therefore scientists have defined a unit called the unified atomic mass unit
(symbol, u), also known as the Dalton (symbol, Da). This 1 u = 1 Da. The unified atomic mass unit is
approximately the mass of one proton (as well as one neutron). Thus , if we know the total number of
protons and neutrons, we know the mass of the atom in units of Da.
Example:
Calculate the mass of an atom that has 6 protons and 6 neutrons.
Since the mass of each proton is 1 Da and the mass of each neutron is 1 Da, the total mass
 1 Da 
 1 Da 
mtotal =
6 protons 
12 Da
 + 6 neutrons 
=
proton
 neutrons 


2.2. Isotopes
We know that the chemical properties and many of the physical properties of an atom depend on the
energy, number, and arrangement of its electrons. We also know that the number of protons must
equal the number of electrons in a free atom. Thus the character of an atom is defined through its
protons and electrons.
Unlike electrons and protons, neutrons do not contribute to the chemical or most of the physical
properties of the atoms. Atoms of a given type can have variable number of neutrons. Thus, for
example, a sample of pure carbon atoms can contain carbon atoms having 6, 7, 8 or more neutrons but
must always contain only 6 protons and 6 electrons; if one of the atoms has a different number of
protons, it cannot be a carbon atom.
When atoms of a given element have the same number of protons but have different numbers of
neutrons they are known as isotopes.
The conventional notation for writing isotopes is as follows:
A
Z
Sy
Here,
Sy = the one to two letter symbol for the atom;
Z = the atomic number of the atom (the number of protons, also equalling the number of
electrons in the free atom); and
A = the mass number of the isotope.
For example, the element hydrogen has three isotopes.
Name
Hydrogen-1
(Protium)
Hydrogen-2
(Deuterium)
Hydrogen-3
(Tritium)
Symbol
Atomic
Number
Number of
Protons
Number of
electrons
Number of
neutrons
Mass
Number
1
1
H
1
1
1
0
1
2
1
H
1
1
1
1
2
3
1
H
1
1
1
2
3
3
Note that all of these atoms are equally “hydrogen” (although the hydrogen atom having 0 neutrons is
by far the most abundant).
3. ELEMENTS
Elements are composed of chemical combinations of isotopes. Thus, for example, the element oxygen is
commonly found as a chemical combination of two oxygen atoms (two identical isotopes of oxygen or
two different isotopes of oxygen). The element oxygen is written as O2 to indicate this.
There are more than a hundred known elements of which about 90 occur naturally on earth. The
remainder of the elements have been produced artificially.
The number of protons present in the nucleus identifies the
atom uniquely as being the atom of a specific element.
Atoms of the same element:
•
Always have exactly the same number of protons and electrons. This number of protons
defines the element uniquely and is known as the atomic number (symbol: Z = the number
of protons in the atom).
•
May have different numbers of neutrons (i.e., may have different isotopes). Each element is
made up of different isotopes in a fixed relative abundance.
Atoms of different elements always have a different number of protons (hence different values of Z).
Compounds (molecules) are chemical combination of atoms of different elements. A compound always
consists of a chemical association of atoms of different elements. The physical and chemical properties
of a compound are different from those of its building block elements. A compound can be broken down
into its constituent elements through chemical changes, although this will often require a lot of energy.
A compound always has the same chemical combination of atoms of its building block elements. For
example, the compound carbon dioxide always consists of one carbon isotope chemically combined with
two oxygen isotopes. It is written as CO2. All samples of carbon dioxide, regardless of their origin, consist
of exactly one atom of carbon and two atoms of oxygen.
4. PURE SUBSTANCES AND MIXTURES
4.1. Pure Substances
A pure substance contains only one type of material. There are only two classes of pure substances:
• elements
• compounds
Pure substances are rarely found in nature.
4.2. Mixtures
Mixtures are by far the most common way that chemicals are found. They are a physical combination of
different compounds or elements (i.e., physically combining different pure substances). Mixtures may
be either homogeneous or heterogeneous.
Homogeneous mixtures (more commonly referred to as solutions, see section 9) are uniform in
composition throughout. For example, when table salt (sodium chloride) is dissolved in pure water, the
4
resulting salt water has a uniform appearance. No part of this mixture differs from any other part, so salt
water is a homogeneous mixture.
Heterogeneous mixtures do not have a uniform appearance. Parts of such mixtures physically appear
different. For example, if sand is poured into water, it settles to the bottom. The mixture of sand and
water is therefore a heterogeneous mixture.
Schematic Description of Substances
Schematically, the relationship between these forms of matter is depicted in the figure below.
Fundamental
Particles
e–, p+, n0
Pure
Substances
Atoms
Elements
Mixtures
Compounds
Heterogeneous
mixtures
Homogeneous
mixtures
Exercise: Classify the following substances as a pure substance, a homogeneous mixture or a
heterogeneous mixture:
1. Wine
2. Air
3. A salt-sand mixture
4. Zinc
5. Oxygen
6. Tap water
7. Steel
8. Wood
9. Blood
10. A page in the textbook
5. THE PERIODIC TABLE OF ELEMENTS
The Periodic Table of Elements is an organisation of atoms according to their properties. It was devised
in 1867 by the Russian chemist, Dimitry Mendeleev and independently that same year by the German,
Lothar Meyer. Mendeleev is considered to be the “Father of Modern Chemistry” because Meyer did not
make any further use of his table.
Mendeleev (and Meyer) noted that many of the known elements have very similar physical and
chemical properties. Those atoms with such similar properties were put in columns. Thus, for example,
atoms such as Na (sodium) and K (potassium) were soft, silvery, greyish metals which reacted
5
explosively with water to form extremely strong basic solutions (see Section 9.2). Elements with such
similar properties were organized into columns, called groups in the Periodic Table.
For example, the group consisting of Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), etc., is
known as the alkali metals. Other metals such as Mg and Ca found in the earth also formed alkaline
solutions with water, but not in nearly the same explosive fashion as Na and K. These metals were
known as the alkaline-earth metals.
This arrangement of atoms according to their properties soon proved its value. Gaps in Mendeelev’s
Table soon began to appear, and he was able to predict the existence of as yet unknown elements.
Chemists began to search for these “missing” elements and found several of them.
The modern Periodic Table has three basic characteristics that provide information about the atom. The
example of the carbon atom is shown below.
Atomic Number, Z, (number
of protons and electrons
present in the free atom)
Mass Number, u, (the average
mass of a carbon atom in
atomic mass units)
6
C
Symbol of the atom
12.011
The elements may be classified according to whether they are metals, nonmetals or metalloids (or
semimetals).
A Periodic Table of Elements is provided on the following page. You are expected to know the names
and symbols of some common elements (these have been highlighted in the Periodic Table on the
following page).
6
PERIODIC TABLE OF THE ELEMENTS
1
1.0079
2
4.003
H
He
hydrogen
hydrogène
3
6.941
4
9.012
Li
lithium
lithium
11
22.99
Na
sodium
sodium
19
39.098
K
potassium
potassium
37
85.468
Rb
rubidium
rubidium
55
132.91
Cs
cesium
césium
87
223
Fr
francium
francium
atomic number
atomic mass
Be
5
10.811
Symbol
beryllium
béryllium
12
24.31
B
English name
French name*
*all are masculine
boron
bore
13
26.98
Mg
magnesium
magnésium
20
40.08
Ca
calcium
calcium
38
87.62
Sr
strontium
strontium
56
137.33
Ba
barium
baryum
88
226.03
Ra
radium
radium
Al
21
44.96
22
47.87
Sc
scandium
scandium
39
88.91
Ti
titanium
titane
40
91.22
Y
yttrium
yttrium
57
138.91
Zr
zirconium
zirconium
72
178.49
La
lanthanum
lanthane
89
227.03
Hf
hafnium
hafnium
104
261
Ac
actinium
actinium
23
50.94
Rf
24
51.996
V
vanadium
vanadium
41
92.91
Nb
niobium
niobium
73
180.95
Ta
tantalum
tantale
105
[268]
Db
rutherfordium dubnium
rutherfordium
58
140.12
Ce
cerium
cérium
90
232.038
Th
thorium
thorium
59
140.91
Pr
praseodymium
praséodyme
91
231.036
Pa
protactinium
protactinium
25
54.94
Cr
Mn
chromium
chrome
42
95.94
manganese
manganèse
43
98
Mo
molybdenum
molybdène
74
183.84
Tc
technetium
technétium
75
186.21
W
tungsten
tungstène
106
[271]
Re
rhenium
rhénium
107
[270]
Sg
seaborgium
60
144.24
Nd
neodymium
néodyme
92
238.029
U
uranium
uranium
Bh
bohrium
61
145
Pm
promethium
prométhium
93
237.048
26
55.85
Fe
iron
fer
44
101.07
Ru
ruthenium
ruthénium
76
190.21
Os
osmium
osmium
108
[277]
Hs
27
58.93
28
58.69
Co
cobalt
cobalt
45
102.91
Ni
nickel
nickel
46
106.42
Rh
rhodium
rhodium
77
192.22
Pd
palladium
palladium
78
195.08
Ir
iridium
iridium
109
[276]
Mt
Pt
platinum
platine
110
[281]
Ds
29
63.546
Cu
copper
cuivre
47
107.87
Ag
silver
argent
79
196.97
Au
gold
or
111
[280]
Rg
hassium
meitnerium
darmstsdtium Roentgenium
62
150.36
63
151.96
64
157.25
Sm
samarium
samarium
94
244
Eu
europium
europium
95
243
Np
Pu
Am
neptunium
neptunium
plutonium
plutonium
americium
américium
Gd
gadolinium
gadolinium
96
247
Cm
curium
curium
7
65
158.93
Tb
terbium
terbium
97
247
Bk
berkelium
berkélium
aluminum
aluminium
31
69.72
30
65.39
Zn
zinc
zinc
48
112.411
Cd
cadmium
cadmium
80
200.59
Hg
mercury
mercure
112
[285]
Cn
copernicium
66
162.50
Dy
dysprosium
dysprosium
98
251
Cf
californium
californium
Ga
gallium
gallium
49
114.82
In
indium
indium
81
204.38
Tl
thallium
thallium
113
[284]
Uut
Ununtrium
67
164.93
Ho
holmium
holmium
99
252
Es
einsteinium
einsteinium
6
12.011
C
carbon
carbone
14
28.086
Si
silicon
silicium
32
72.61
Ge
germanium
germanium
50
118.71
Sn
tin
étain
82
207.2
7
14.007
N
nitrogen
azote
15
30.974
P
phosphorus
phosphore
33
74.922
As
arsenic
arsenic
51
121.76
Sb
antimony
antimoine
83
208.98
Pb
lead
plomb
114
[289]
Bi
bismuth
bismuth
115
[288]
Uuq
Uup
8
15.9994
O
oxygen
oxygène
16
32.07
S
sulfur
soufre
34
78.96
Er
erbium
erbium
100
257
Fm
fermium
fermium
69
168.93
Tm
thulium
thulium
101
258
Md
mendelevium
mendélévium
9
18.998
F
fluorine
fluor
17
35.453
Se
selenium
sélénium
52
127.60
Cl
Br
bromine
brome
53
126.904
Te
tellurium
tellure
84
209
Po
polonium
polonium
116
[293]
Uuh
70
173.04
Yb
ytterbium
ytterbium
102
259
No
nobelium
nobélium
Ne
neon
néon
18
39.95
chlorine
chlore
35
79.904
Ununquadium Ununpentium Ununhexium
68
167.26
helium
hélium
10
20.18
I
iodine
iode
85
210
Ar
argon
argon
36
83.80
Kr
krypton
krypton
54
131.29
Xe
xenon
xénon
86
222
At
astatine
astate
117
[294]
Uus
Ununseptium
71
174.97
Lu
lutetium
lutécium
103
262
Lr
lawrencium
lawrencium
Rn
radon
radon
118
[294]
Uuo
Ununoctium
1.1 Classification of the Elements
Atoms of the elements are organized in the Periodic Table according to property in the sequence of their
atomic numbers, Z (the number of protons in the atom). There are several ways in which the elements
are classified and it is necessary to know them.
1.1.1 Metals and Nonmetals
Halogens
Chalcogens
Pnictogens
Alkaline-Earth metals
Alkali metals
The vast majority of elements are metals.
B
Transition metals
Aℓ
Si
Ge
As
Sb
Te
Po
At
Non-metals are to the right of the stepladder structure (bold lines).
Metalloids (semi-metals) are at the boundary of the stepladder and are indicated on the Periodic Table
above by the lighter shade of grey. They have many properties typical of the metals but are generally
semi-conductors rather than the excellent conductors of electricity and heat that metals are.
1.1.2 MONATOMIC AND DIATOMIC ELEMENTS
The elements written in outline font in the Periodic Table above are diatomic elements. That is, they are
found in nature under normal conditions as a chemically bonded pair of atoms. Thus, for example, the
elements hydrogen, nitrogen, oxygen and iodine (as well as fluorine and bromine) are found respectively
as:
H2, N2, O2, Cℓ2 and I2.
The first four diatomic elements in the list above (H2, N2, O2and Cℓ2) are gases; of the last element I2, is a
solid.
Apart from this group of diatomic elements, all other elements are treated as if they are monatomic;
that is, the atom is the same as the element. This is strictly speaking not true but is useful for the
purposes of simplicity.
8
1.1.3 Atoms in Living Organisms
Although there reportedly 115 known elements as of 2013, only about 90 are found in nature; the other
elements have been synthesized (created) by scientists in the laboratory. Of the 90 elements found
naturally, living organisms require only about 20 of them, and of these 20, the 95% of the mass of the
human body is made up of just four:
Carbon
C
Hydrogen
H
Oxygen
O
Nitrogen
N
Some of the important elements in living organisms have been highlighted in the Periodic Table on page
7 which you have been asked to learn. They are listed (again) below.
Element Name
Atomic Symbol
Hydrogen
H
Carbon
C
Nitrogen
N
Oxygen
O
Sodium
Na
Magnesium
Mg
Phosphorus
P
Sulfur
S
Chlorine
Cℓ
Potassium
K
Calcium
Ca
Iron
Fe
9
6. FORMATION OF COMPOUNDS
In nature, all systems tend to move to the state of greatest stability. Thus, for example, atoms combine
with each other to form compounds only so that they can become more stable.
Recall that an atom consists of protons and neutrons densely packed into a nucleus at the centre of the
atom. The electrons are moving rapidly outside of, and far away, from the nucleus. Because of this
separation between the electron and its nucleus, the electron can “forget” to which nucleus it belongs
and attach itself to the nucleus of another atom that may be closer. Therefore the electrons of an atom
can become “lost”.
Thus, atoms are capable of losing and gaining electrons. This movement of electrons between nuclei
results in the formation of chemical bonds – either ionic bonds or covalent bonds.
1.2 Ionic Bonds
If an atom gains an electron, it acquires an overall negative charge because each electron carries a
negative charge.
e.g., Cℓ + e– → Cℓ–
Similarly, if an atom loses an electron, it acquires a positive charge (i.e., it loses a negative charge).
e.g., Na → Na+ + e–
+
Atoms acquire a charge only by gaining or losing electrons NOT by gaining or losing protons. Note that
the total number of electrons is conserved in the compound. The electrons are simply jumping from one
nucleus to another.
Substances may lose or gain more than one electron.
e.g.
Fe → Fe3+ + 3 e–
O + 2 e– → O2–
Anions are chemical substances which have negative charges (e.g., N3–, O2–, I–)
Cations are chemical substances which have positive charges (e.g., H+, Na+, Ca2+)
10
Typically:
Metals, to the left of the Periodic Table, tend to lose electrons thereby forming cations.
Non-metals tend to gain electrons thereby forming anions.
In general, substances will not give up electrons unless there are other atoms ready to receive those
electrons. Likewise, a substance cannot accept electrons without other substances ready to supply those
electrons. Thus, for example, sodium can lose an electron:
Na → Na+ + e–
but will not do so unless it is in the presence of another atom willing to receive that electron, such as
chlorine:
Cℓ + e– → Cℓ–
Combining these equations
Na + Cℓ + e– → Na+ + e– + Cℓ–
Cancelling the electrons on either side of the equations yields the net result:
Na + Cℓ → Na+ + Cℓ–
Note that there are no net free electrons remaining at the end. All of the electrons lost by Na atoms
have been absorbed by the Cℓ atoms. At this point the system consists of a cation (Na +) and an anion
(Cℓ–). These two oppositely charged ions now attract each other to form an ionic bond.
Na+ + Cℓ– → Na+
Cℓ– = NaCℓ
The attraction of positive to negative results in the
formation of an ionic bond. The resulting neutrally
charged compound is known as an ionic compound.
Ionic bonds are typically very, very strong. However, when ionic compounds are dissolved in various
solutions, the chemical environment can serve to weaken the ionic compound and even break the ionic
compound apart. This is because individual atoms of the ionic compound can interact with the
compounds in the surroundings.
Ionic compounds are often referred to as salts. Salts release ions when dissolved in water (see Section
9); thus mixtures of salts and water have the ability to conduct electricity. In this context, salts are
therefore sometimes called electrolytes: electrolyte = salt = ionic compound.
1.3 Covalent Bonds
As mentioned, nonmetals, to the right of the step-ladder in the Periodic Table tend to gain electrons. In
ionic compounds, a metal atom is present to supply the needed electron(s). If, however, no such
supplier is present (i.e., there is no metal), another method of acquiring the needed electrons must be
found.
Consider the case of a sample consisting only of 17Cℓ atoms. Each chlorine atom in the sample needs one
electron to become more stable, which it can only obtain by stealing it from a nearby chlorine atom.
However none of the chlorine atoms want to give up their electrons. Each chlorine atom seeks to seize
one electron from the other chlorine atom while retaining all of its own electrons.
11
Cl
Cl
The result is that the two atoms are drawn together with the two electrons, one from each chlorine
atom, located between the two nuclei.
Cℓ
Cℓ
This type of bond is known as a covalent bond. In covalent bonds, the electrons are shared between two
atoms.
Covalent bonds are formed between two nonmetal atoms.
7. BOND POLARITY
Covalent bonds are formed when two nonmetal atoms combine chemically by sharing electrons.
However the electrons are not always shared equally. Consider the bond between a carbon and oxygen
atom. Both atoms are nonmetals so the two atoms share a pair of electrons. However the carbon atom
has 6 protons, whereas the oxygen atom has 8 protons. The shared electrons (negative) will therefore
be pulled more strongly toward the more positive oxygen nucleus:
6C
8O
Electron pair is
pulled toward
the oxygen
This half of the bond
appears to be positive
because the electrons
are on the other side
6C
Midpoint
between the
two atoms
δ+
Partial positive charge
This half of the bond
appears to be negative
because the electrons
are on this side
8O
δ–
Partial negative charge
The result is that one half of the bond (in this case on the oxygen side) appears to have a negative
charge while the other half of the bond appears to be positive (the carbon side). The negative side is
referred to as having a partial negative charge (δ–); remember this is a covalent bond. In an ionic bond
this would be a full negative charge (–). Similarly, the positive side is referred to as having a partial
positive charge (δ+).
A polar bond is sometimes simply represented by an arrow pointing from the positive end to the
negative end of the bond.
The unequal sharing of electrons in a covalent bond results in one end of the bond being partially
negative while the other end is partially positive. This bond is said to be a polar covalent bond (or just
simply a polar bond). Most covalent bonds are polar. The common bond types in biology and their
relative polarity are provided below.
A
B
δ+
δ–
12
H—C
C—N
C—Cℓ
N—O
H—N
Equal bond polarities
Least Polar
H—Cℓ
C—O
H—O
Most Polar
Equal bond polarities
The ability of an atom to attract electrons to itself compared with other atoms is termed its
electronegativity. All of the atoms in the periodic table have assigned values of electronegativity. By
comparing the electronegativities of the atoms in a covalent bond, we can predict which of the atoms
will have the partial negative charge and which will have the partial positive charge. The
electronegativity difference between atoms will give the strength of the bond polarity. The
electronegativity trends among the atoms is shown in the schematic Periodic Table below.
F
Largest EN
Larger EN
Smaller EN
Smallest EN
Larger EN
Cs
Smaller EN
As we go down the Periodic Table, atoms get larger and larger; consequently this reduces their ability to
attract electrons, so their electronegativity (EN) decreases. As we cross a given row in the Periodic
Table, the number of protons in the atoms increases, thereby increasing their ability to attract electrons,
so we observe an increase in electronegativities.
8. Hydrogen Bonding and Water
There are several very important consequences of bond polarity. These relate to the shape of molecules,
an overall polarity of the molecule and some physical properties of the molecule.
Consider the case of the molecule carbon dioxide, CO2. Carbon dioxide is a linear molecule with carbon
in between two oxygen atoms.
180°
O
C
O
(The double lines represent double bonds which are two electron pairs linking the atoms together.)
From the discussion in Section 7, the C—O is quite polar.
C—O
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There are two polar bonds present in the structure of CO2, one for each (C—O) bond.
O
C
O
However, since the O —C—O angle is 180°, the “pull” of electrons to the left is precisely balanced by the
“pull” of the electrons to the right. The net effect is that there is no net “pull” so the molecule overall is
nonpolar.
Contrast the CO2 molecule with water, H2O. Water has a bent structure with oxygen in between two
hydrogen atoms.
O
H
104.5°
H
Oxygen attracts electrons much better than hydrogen, so each H—O is highly polar. However, the fact
that the water molecule is bent means that all of the “pull” of electrons do not cancel off as in CO2.
O
H
H
In fact, water has a net overall polarity, that is, it has an overall positive end and an overall negative end.
Therefore the whole water molecule can be represented by an arrow as for individual polar bonds.
δ–
δ+
If two water molecules are brought in contact with each other, since they each have partial positive and
partial negative ends, the positive end of one water molecule will naturally align with the negative end
of the other water molecule. This attraction polar between molecules has consequences such as high
boiling temperatures. The data for carbon dioxide and water presented below illustrate the significance
of these intermolecular forces.
Molecule
O
C
O
O
H
H
Polarity
Boiling Temperature at normal pressure
Nonpolar
–57°C
Highly Polar
+100°C
The interaction between different molecules is known as intermolecular forces (i.e., forces between
molecules).
The bond polarities of the N—H and O—H bonds are particularly strong in comparison with other types
of intermolecular forces (although still weak compared with a “real” bond). To indicate this,
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intermolecular forces between compounds that have N—H and/or O—H bonds are given the name
hydrogen-bond.
Schematic Illustration of Hydrogen Bonding
δ–
δ–
O
H
δ+
O
H
δ+
H
δ+
H
Hydrogen Bond
Water Molecule
δ+
Water Molecule
Hydrogen bonding is not necessarily limited to pairs of molecules. It is possible for many water
molecules to form clusters:
Hydrogen bonding is extremely important in nature. The high boiling point of water has been noted.
Another example is density: most materials are denser in the solid (frozen) state than in the liquid form
resulting in the solid being “heavier” than the liquid. The solid therefore sinks to the bottom in
substance that is becoming frozen. Water is the opposite. Ice actually floats on top of liquid water
because water molecules in their solid form arrange themselves so as to maximize the number of
hydrogen bonds; this causes the molecules to spread out, reducing the density compared with the liquid
water form. Other properties of water that can be attributed to hydrogen bonding are the shapes of
molecules such as proteins and DNA. The two strands of DNA, for example, are held together through
hydrogen bonding.
9. SOLUTIONS, ELECTROLYTES AND pH
A solution is a homogeneous mixture, which means that it is uniform throughout in both physical
appearance and at the molecular level. There are two general classes of materials that make up a
solution:
1. The solute: the substance (or substances) in lower quantities;
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2. The solvent: the substance in the highest quantity. In the human body, and in all living things,
the solvent is water.
A solution is most likely to be formed when solute molecules are attracted to solvent molecules. For
example, when polar covalent compounds dissolve in water, one pole of the compound is attracted to
the polar water molecule (as discussed in section 8) and are soluble.
Polar molecules such as vitamin C can dissolve in water. Thus even an excess of vitamin C can be
eliminated from the body in the urine.
Conversely, since nonpolar molecules do not form polar interactions with water, they are not soluble.
For example, bacon eaten is large quantities is unhealthy. The fat in bacon contains nonpolar covalent
bonds so it does not dissolve in water. Therefore bacon fat will aggregate with other insoluble fats.
These fat aggregates can form deposits in blood vessels, termed plaque, leading to cardiovascular
problems.
When an ionic material is dissolved in water, something similar to the interaction between two water
molecules occurs. The positive end of a water molecule will align with the anion (negative ion) of the
ionic compound. Similarly, the negative end of a water molecule will align with the cation (positive ion)
of the ionic compound. These attractions can pull the ionic compound apart resulting in the freed ions;
the ionic compound is said to have dissociated (dis + associated) in water. The freed ions can conduct
electricity, thus any substance that releases ions in water will be an electrical conductor and is therefore
called an electrolyte. Strong electrolytes dissociate completely or almost completely in water. For weak
electrolytes, only some of the ionic compound will dissociate so there will not be many ions in solution.
Having a sufficient intake of electrolytes is absolutely necessary to proper body functions. For example,
electrolytes are involved in maintaining water balance (which, for example, has an effect on blood
pressure) and on normal cell function.
Of particular interest in chemical and biological systems are acids and bases, so the following definitions
are important.
9.1.
Acids
Acids can be defined as substances that release H+ in water. Acids can be strong (strong electrolytes) or
weak (weak electrolytes).
Examples of common acids are:
Hydrochloric acid (HCℓ)
Strong
Carbonic acid (H2CO3)
Weak
Citric acid (C6H8O7)
Weak
Phosphoric acid (H3PO4)
Weak
Sulfuric acid (H2SO3)
Strong
Acetic acid (CH3COOH)
Weak
Do not confuse the terms strong and weak with more dangerous or less dangerous. The terms “strong”
and “weak” in chemistry simply refer to the degree of dissociation. For example, one of the most
dangerous acids is hydrofluoric acid which is classified as a weak acid.
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Examples of common acids in the human body include hydrochloric acid which is produced in the
stomach. One of many reasons for the production of this acid is to help destroy bacteria that have
entered the stomach with ingested food.
You may have assumed that the amount of carbon dioxide (CO2) you exhale is linked to the amount of
oxygen (O2) that you inhale. This is not so. CO2 is a by-product of your cells’ work and is transported
from your cells to your lungs as carbonic acid regardless of how much O2 you breathe in.
9.2.
Bases
Bases can be defined as substances that can bind H+. As with acids, bases can be either strong or weak.
Examples of common bases are:
Ammonia (NH3)
Weak
Sodium bicarbonate (NaHCO3)
Weak
Sodium carbonate(Na2CO3)
Weak
Sodium hydroxide (NaOH)
Strong
Sodium hypochlorite (NaOCℓ)
Weak
*Sodium hypochlorite is the important component of bleach.
Examples of common bases in the human body include ammonia which is produced in the body through
the breakdown of proteins, for example during exercise. The ammonia dissolves in water and is
eliminated in urine.
Sodium bicarbonate absorbs excess hydrogen ions in the stomach contents as the hydrogen ions move
into the small intestine. If this absorption does not occur, the hydrogen ions (from the hydrochloric acid
in the stomach) would burn holes in the intestinal tract.
9.3.
pH
The pH scale is used as a measure of acid concentration or base concentration in water solutions. Acids
release H+ in water, so the greater the number of H+ ions present in the solution the more acidic the
solution will be. Bases bind H+, therefore the smaller the number of H+ ions present in the solution the
more basic will be the solution. The pH scale has a range of 0–14 (although for concentrated solutions it
is possible for the range to extend outside of these limits).
Pure water has equal amounts of H+ and OH– ions (the OH– ion is an acceptor of H+ so it is a base):
H+ + OH– → H2O
The middle of the pH range is 7. Therefore a neutral solution has a pH = 7. Acidic solutions will have
pH < 7 while basic solutions will have pH > 7.
Stronger Acid
pH < 7 Acid
pH > 7 Base
High H+
Low H+
pH = 7
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Stronger Base