CHEMISTRY DEPARTMENT, FUHSO
CHM 101 (General Chemistry)
Lecture Note Prepared by: Rotella Idongesit A.
Reference Texts: Ebbing Fourth Edition General chemistry (Pg.444-); Chemistry in context
TOPIC: Gross features of atomic structure; Periodic Table.
Introduction: In order to consider the gross features of atomic structure, we will review the
following;
The concept of existence of atom
It was Greek philosophers Leucippus and Democritus who first proposed the existence of an
atom around 5th century. “That matter cannot be endless divided and a continued sub-division
of matter ultimately yield an atom” which cannot be split further. For about 2000 years, the
Greek philosophers’ concept of existence of an atom were mere speculation as they did not carry
out any planed experiment to convinced scientists to believe their views.
Dalton’s atomic theory
Between 1803-1808, John Dalton proposed atomic theory, “that atom is the smallest indivisible
particle of an element”. It was Dalton’s atomic theory that convinced scientist beyond doubt to
believe on the existence of an atom since his theory was proceeded based on series of
experimentations. Dalton’s atomic theory is summarized into the following postulates;
(i) Matter is made up of small indivisible particles called atoms.
(ii) Atoms can neither be created nor destroyed
(iii) Atoms of the same element are identical or alike in every aspect but different from atoms
other elements.
(iv) Atoms combine to form compounds and they do so in simple whole number ratio.
(v) All chemical changes occur as a result of combination or separation of atoms.
Limitation and modification of Dalton’s atomic theory
(i)
Dalton’s atomic theory did not explain or make provision for the possibility of
existence of isotopes (i.e atoms of the same element which have different masses) and
a reaction that could lead to creation or destruction of atoms and a possibility where
some atoms of elements can combine in non-integral ratios. These limitations
therefore led to the modification of some of the postulates of Dalton’s atomic theory
which include the following;
Modification of Dalton’s atomic theory
Four of the five postulates of Dalton’s atomic theory have been modified as given below;
(i) Atoms are divisible as it contains other smaller particles such as protons, neutrons and
electrons. (i.e atom is no longer an indivisible particle)
(ii) Atoms can be created and destroyed in nuclear reactions
(iii) Atoms of the same element are not identical or alike in every aspect due to existence of
isotopes
(iv) In some larger or macro organic molecules compounds like polymers, atoms combine in
non-integral ratios.
Simple atomic structure
In 1900, the alpha particle scattering experiment of Ernest Rutherford led to the discovery of the
nucleus of an atom. The conclusions drawn from of the experiment included that, an atom has a
dense of about surrounded by empty spaces. In addition, in 1913, Niel Bohr proposed that “an
atom has imaginary lines orbiting the nucleus and the lines or orbits (or shells) are situated at
various distances away from the nucleus”. Therefore, in consideration of Rutherford and Niel Bohr
experimental conclusions, and for simplicity “electron-dot method is used to depict a simple
Shell or orbit
atomic model or structure as thus;
e
Electron revolving round the nucleus
Nucleus containing protons and neutrons
(i)
Arrangement of atomic shells
O-shell
N-shell
M-shell
K-shell
L-shell
Shells or orbits are imaginary lines orbiting or surrounding the nucleus of an atom. The shells
are assigned letters K, L, M, N, O, …….; with respect to the principal quantum number n = 1, 2,
3 ,4, ……. For n = 1, it is the K-shell and it is the shell closest to the nucleus, followed by n =2,
which is the L-shell.
(ii) Nucleus: Nucleus is the atomic central core which is positively charged and contains protons
and neutrons and most of the atomic mass. The nucleus has a diameter of about 10-15 m (10-5 Å)
and which is a hundred thousand times smaller than the atom.
The fundamental or basic sub-atomic particles
The discovery of electron by a British physicist in 1897 through his cathode ray experiment
followed by the discovery of the nucleus, and coupled with other developments led to the
understanding that an atom is not indivisible particle as it contains much smaller particles called
sub-atomic particles and which among others include the basic ones known as electron, proton,
and neutron.
(i) Electrons: Electron is a very light, negatively charged particle the exist in the region around
or outside the positively charged nucleus. Electron has a mass of 9.10939 x 10-31 kg, a charge of
-1.60218 x 10-19 C, atomic mass unit (amu)* of 0.00055 and approximate charge (e) of -1. Electron
is the smallest of the three fundamental particles of an atom.
(ii) Protons: A proton is a nuclear particle having a positive charge equal to (+e) and a mass
more than 1800 times that of electron. Proton was discovered in 1911 by Ernest Rutherford and
it has a mass of 1.67262 x 10-27 kg, a charge of +1.60218 x 10-19 C, atomic mass unit (amu)* of
1.00728 and approximate charge (e) of +1. It is bigger in mass than electron but smaller than
that of neutrons.
(iii) Neutrons: Neutron is a nuclear particle having a mass almost identical to that of the proton
but has no electrical charge. Neutron was discovered by James Chadwick in 1932 and has of
1.67483 x 10-27 kg, a charge of 0C, atomic mass unit (amu)* of 1.00866 and approximate charge
(e) of 0. Neutron has the biggest in mass than proton and electron.
Relative or approximated masses and charges of sub-atomic particles
Particle
Electron
Proton
Neutron
Position
Mass
Approx. Charge
in atom
mass
-31
Shell/orbit 9.11×10 Kg
1/1840 - 1.602 x 10-19C
Nucleus 1.67265×10-27Kg
1
+1.602 x 10-19C
Nucleus 1.67500×10-27Kg
1
0
Approx.charge
-1
+1
0
Mass number and Atonic number
Atomic symbol which is a one- or two- or three-lettered notation used to represent an atom
corresponding to a particular element carries mass number and atomic number. While mass
number is denoted with a letter (A), atomic number is denoted with a letter (Z). Every symbol of
an atom of an element carries two inscriptions or numbers. While the superscript is the mass
number, the subscript is the atomic number as shown;
Mass Number
Atomic Number
Mass Number
𝐴
𝑧𝑀
E.g
214
90𝑃𝑏
Atomic Number
Mass number or nucleon is defined as the number of protons and neutrons in the nucleus of
atom of an element. It is expressed as; mass number = number of proton(s) + number of
neutron(s). Atomic number is defined the number of protons in the nucleus of atom of an
element. Or it is the number of electrons of an atom outside the nucleus. Since atom of an element
has equal number of protons and electrons which cancel out each other, an atom of an element is
electrical neutral.
Isotopy
From the modification of Dalton’s atomic theory “atoms of the same element are not identical or
alike in every aspect due to existence of isotopes’ and the production of mass spectrometer, it
was known that most elements occur naturally in isotopic mixture of atoms with slight different
masses. These atoms of an element with slight different mass are called isotopes. The phenomenon
or occurrence or existence of two or more atoms of an element having the same atomic number
but different mass number is known as isotopy. Isotopes therefore are defined as atoms of a
particular element whose nuclei have the same atomic number but different mass number; that is
the nuclei have the same number of protons but different number of neutrons.
Mass spectrometric analysis is a technique commonly used to determine the exact number of
isotopes of an element, the masses of the isotopes, percentage and fractional abundance of
each isotope using an instrument called mass spectrometer.
Examples of elements the exhibit isotopy
(i) Oxygen: Oxygen exists naturally in three isotopic forms as
(a) oxygen-16 ( 𝟏𝟔𝟖𝑶), with fractional abundance of 99.759%,
(b) oxygen-17 ( 𝟏𝟕𝟖𝑶), with fractional abundance of 0.037% and
(c) oxygen-18 ( 𝟏𝟖𝟖𝑶) ,with fractional abundance of 0.204% .
(ii) Neon: The fractional abundance in of neon in naturally occurring neon are;
(a) Neon-20 ( 20
10𝑁𝑒 ), with fractional abundance of0.9051%,
(b) Neon-21 ( 21
10𝑁𝑒 ), with fractional abundance of 0.0027% and
(c) Neon-22 ( 22
10𝑁𝑒 ) ,with fractional abundance of 0.0922% .
NOTE: Fraction abundance of an isotope is the fraction of the total number of atoms that is
composed of a particular isotope whereas percentage abundance of an isotope is the fraction of
the total number of atoms that a particular isotope is composed of expressed in percentage.
Relative atomic mass of atoms of elements
Definition of relative atomic mass: Relative atomic mass of elements is defined as the mass of
an atom of an element compared to 1/12 the mass of an atom of carbon-12 isotope. Or It is defined
as the average mass of one atom of an element compared to 1/12 mass of one atom of 12C. Or it is
defined as the number of times that an atom of an element is heavier than 1/12 of the mass of
carbon-12 isotope.
Calculation of relative atomic mass of elements from percentage abundance of isotopes
The relationship or formula given below is used;
R.A.M of element = (% abund. of 1st isotope x mass No) + (% abund. of the 2nd isotope x mass No) + ……
100
100
Example 1: Mass spectrometric analysis shows that, chlorine occurs in two isotopic forms as
37
chlorine-35 ( 35
17𝐶𝑙 ) and chlorine-37 ( 17𝐶𝑙 ) with percentage abundance of 75.05% and 24.95% ≈
Solution
R.A.M of element = ( % abundance of
100
35
17𝐶𝑙
x 35) + (% abundance of
100
37
17𝐶𝑙
x 37 )
= (75.05 x 35) + (24.95 x 37) = 26.2675 + 9.2315 = 3.499 ≈ 35.5.
100
100
Calculation of relative atomic mass of elements from Ratio of abundance of isotopes
R.A.M of element = (Ratio of 1st isotope x mass No) + (Ratio of the 2nd isotope x mass No)
The sum of the ratios
The sum of the ratios
Example 1: Two isotopes of an element Y with mass 18 and 20 occur in the ratio of 1:2, calculate
the relative atomic mass of the element Y.
Solution
Relative atomic mass of Y = (1 x 18) + (2 x 20) = 6 + 13.33 = 19.33.
1+2
3
OR
Relative atomic mass of Y = (33.33 x 18) + (66.66 x 20) = 5.9994 + 13.332 = 19.33.
100
100
Exercise
(1) Magnesium occurs in three isotopic forms as 2412Mg, 2512Mg and 2612Mg with percentage
abundance of 78%, 10% and 12% respectively. Calculate the relative atomic mass (R.A.M) of
Magnesium;
(2) Silver exists in nature as 10747Ag and 10947Ag in the ratio of 3:1. What is the relative atomic
mass of silver?
(3) An element X consists of two isotopes; 95X and 115X in the ratio of 9:1. Calculate the mean
relative atomic mass of X.
(4) Copper exists naturally in two isotopic forms as 6329Cu and 6529Cu. If the relative atomic
mass of copper is 63.546, calculate the percentage abundance of each of the two isotopes.
(5) Copper exists naturally in two isotopic forms as 6329Cu and 6529Cu. If the relative atomic
mass of copper is 63.546, calculate the percentage abundance of each of the two isotopes.
PERIODIC TABLE
By 1900 over fifty (50) elements were discovered by different scientists including the first ever
discovered metals sodium and potassium which were isolated by Humphry Davis. By that time,
chemists began to experience difficulty of memorizing a long list of properties of elements and
this therefore became a necessity for chemists to arrange into groups with similar properties.
Among many scientists who attempted arranging elements into groups are, Johannes Dobereiner
who arranged elements into what he called “triads” and John Newlands who grouped elements
into what he called “octave” (i.e musical notes). However, in 1869 a Russian chemist Dmitri
Mendeleev and the German Chemist J. Lother Meyer, working independently made similar
discoveries and they published that when elements are arranged in order of their increasing
atomic weight or masses, their properties reoccur at certain intervals or are periodic in
nature.
Because of other elements discovered years later and these elements fit-in into the gaps that
Mendeleev created in his periodic table, his work was adjudged or considered as the origin of the
modern periodic table.
Mendeleev postulates on Periodic Table: Mendeleev postulated that; (i) Chemical properties of
elements are the periodic function of their atomic masses or weight. Or (ii) Elements are arranged
in the periodic table according to the increasing order of their atomic masses or weight.
Definition of periodic Table: A periodic table is a tabular arrangement of elements in rows
and columns highlighting or showing the regular repetition of properties of elements.
Differences between Mendeleev periodic table and the modern periodic table
Mendeleev Periodic Table
Modern Periodic Table
Elements were arranged in order of their Elements are arranged in order of their
increasing atomic weight or masses.
increasing atomic number or protons.
Gaps were created for future elements to No gaps are left for future elements
be discovered
Periodic law of elements: The law states that the properties of elements (physical and chemical)
are the periodic function of their atomic number. Or the law states that when are arranged in order
of their increasing atomic number, their properties repeat at certain intervals.
Features/components of a typical modern periodic table
A typical modern periodic table is a chart that is made up of eight (8) vertical columns called
the groups or families and seven (7) horizontal rows called the periods.
Groups of elements: A group consists of the elements in any one vertical column in the periodic
table. Or a group is any of the eight vertical columns of elements in the periodic table. All elements
in the same group have the same number of valence electron(s) (i.e the same number of
electron(s) in the outer most shell of their atoms). The elements in the periodic table are further
grouped into two sub-groups called the A and B. Elements of the B sub-group are the transition
elements while those in the A sub-group are called the representative elements.
Periods of elements: A period consists of the elements in any one horizontal row in the periodic
table. Or a period is any of the seven horizontal rows of elements in the periodic table. All
elements in the same period have the same number of electron shell(s). For example, all
elements in period one (I), have one electron shell, those in period two (II) have two (2) electron
shells and so on.
Period one (1) is the shortest period of the periodic table containing only two elements running
from hydrogen to helium. Periods two (2) and three (3) contain eight elements in each, running
from lithium to neon in period (2) and from sodium to argon in period (3). Periods four (4), five
(5) and six (6) contain eighteen (18) elements in each running from potassium to krypton in period
(4), rubidium to xenon in period (5) and from cesium to radon in period (6). Period seven (7)
contains thirty-two (32) including the inner transition elements; lanthanide and actinide metals.
Determination of position of position of atoms of elements in the periodic table
Position of element in the group
The number of valence electron(s) or
electron(s) in the outer most shell or proton(s)
of atom of an element determines the group in
which an element belong to in the periodic
table
Position of an element in the period
The number of electron shell(s) or the total
number shells of electron(s) of an atom of
any element determines the period in which an
element belong to in the periodic table
Periodic Properties of elements
Definition of periodicity of properties of elements: Periodicity is the variation of the
properties of elements in a regular pattern both down the groups and across the periods. Or
it is the reoccurrence of properties (physical and chemical) of elements at certain intervals
when elements are arranged in order of their increasing atomic number. The periodic
properties include;
(i) atomic radius/size/volume and ionic radii/size
Definition of atomic radius: Atomic radius is defined as half the distance between the nuclei of
two atoms of the same element that are covalently bonded in a molecule. Or it is defined as half
the distance between the nuclei of any two close atoms of an element in a covalent molecule. Or
atomic radius is also defined as a measure of the size of atoms of an element, the mean distance
from the centre of the nucleus to the boundary of the surrounding shells of electrons.
H
H
Half of this line is atomic radius
Mathematically, atomic radius = Bond distance or bond length.
2
Trends/variations in the periodic table: Atomic radius or size or volume/ionic radius or size of
atoms of elements (i.e the value) decreases a cross any period from left to right and increases down
any group of elements in the periodic table from top to the bottom.
Reasons the trends: Atomic size or radius decreases across any period of elements in the periodic
table from left to right due to increase in effective nuclear charge and no increase in the
screening effect of inner shell(s) across a period of elements. However, it increases down the
groups of elements due to addition of new electron shell(s) and also due to the screening effect
of inner electrons (i.e the inner shell or of electrons help to shield the outer electrons from the
positive charge influence of the nucleus).
Arrangement of ions and atoms in order of ionic or atomic sizes
(i) Cl- ˃ K+ ˃ Ca2+ . Or Ca2+ ˂ K+ ˂ Cl-. The ions or species are iso-electronic or have the electron
configuration and for isoelectronic species or ions, the greater the nuclear charge or the greater
the magnitude of the charge or proton number, the smaller the size of the ion.
(ii) ionization energy(I. E):
Definition of ionization energy: The first ionization energy is the energy required by an isolated
gaseous atom to remove its most loosely heled electron(s) from the outer most shell. Or it is the
energy required by neutral atom in its ground state to lose one or more of its electrons from its
outer most shell. Or the first ionization energy of an atom is the energy required to completely
remove the most loosely bound electron from a neutral gaseous atom. The S.I unit of ionization
energy is KJ/mol.
Consider this ionization of sodium (Na); Na(s) → Na+ + e-. It requires 494KJ/mol. of energy to
remove the one mole of electron
Trends of I.E in the periodic table: Ionization energy or the value of ionization energy increases
across the period of elements from left to right and it decreases down the group of elements from
top to the bottom.
Reasons for the trends: Ionization energy increases across the periods of elements from left to right
due to increasing effective nuclear charge, decreasing atomic radius and decreasing metallic
property across the period. However, it decreases down the groups of elements from top to the
bottom due to increasing atomic radius, screening effect of inner electrons and increasing
metallic property of elements down the groups.
Factors that determine the ionization energy of an atom of elements
(i) Atomic radius or size of an atom of elements (ii) Nuclear charge (iii) The shielding effect or
screening effect of inner electrons (iv) electron configuration (v) the penetration effect
(iii) electron affinity (E.A)
Definition of electron affinity: Electron affinity is the energy required by an isolated gaseous
atom to gain or accept electron(s) to form a negatively charged ion. Or it is the energy associated
with the process of adding electron(s) to neutral atom in its ground state to form anion. Or it is the
energy required by an atom to gain completely one or more electrons in order form a negatively
charged neutral gaseous ion. The S.I unit of electron affinity is KJ/mol.
Trends of electron affinity in the periodic table: Electron affinity or the value of electron affinity
increases across the period of elements from left to right and it decreases down the group of
elements from top to the bottom.
Reasons for the trends: Electron affinity increases across the periods of elements from left to
right due to increasing effective nuclear charge, decreasing atomic radius and decreasing
metallic property across the period. However, it decreases down the groups of elements from top
to the bottom due to increasing atomic radius, screening effect of inner electrons and
increasing metallic property of elements down the groups.
(iv) Electronegativity:
Definition of electronegativity: It is defined as the measure of ability of an atom of an element in
a covalent molecule to attract the bonded pair of electrons to itself. Or it is defined as the tendency
of an atom to pull to itself the shared pair of electrons in a covalent bond. Or it is also defined as
the ability of an atom to attract shared electrons between the atom and another atom in a covalent
molecule or chemical bond. Or it is attractiveness of bonded electrons by an atom of an element
in a covalent molecule or chemical bond.
Variation or trends of electronegativity in the Periodic Table
Trends of electronegativity in the periodic table: Considering the elements in period II and
group VIIA; electronegativity value of elements increases across the periods of elements from
left to right and decreases down the group of elements from top to the bottom
N = 2.5; O = 3.5; F = 4.0;
Cl = 3.0;
Br = 2.8; The halogens or group VIIA elements in the periodic table are the
I = 2.5; electronegative elements and fluorine has the highest electronagetiAt = 2.3; vity value than all other halogens. Electronegativity vales of elements
are obtained using Linus Pauling’s scale formula and values are unit-less or has no unit because
they are obtained as a scale as thus; Electronegativity = Ionization energy + Electron Affinity
2
(v) Metallic and non-metallic properties:
Generally, the metallic property of elements increases down the group of elements from top to the
bottom and decreases across the period from left to right. Whereas the non-metallic property of
elements decreases down the group from top to the bottom and increases across the period from
left to right.
(vi) Boiling and melting points:
For representative metals (i.e group IA-IVA metals), their boiling and melting points decrease
down the group from top to the bottom because metallic property increases down the group and.
and increases across the periods from left to right because decrease in metallic property.
However, for non-metals, their boiling and melting points increase down the group from top to
the bottom and decrease across the periods from left to right.
NOTE: (i) Metals that exist as liquids at room temperature include; (a) Mercury (b) Francium (c)
Gallium (d) Caesium (e) Rubidium etc.
Exercise
(i) Define the tern atomic radius (ii) Explain briefly the meaning of screening effect of inner
electrons (iii) Define ionization energy, electron affinity and electronegativity and explain briefly
with reasons their trends in the periodic table.
2ND TOPIC:
HYDRATION AND HYDROLYSIS
INTRODUCTION: To understand the two chemical terms properly, it is important to note that
the two terms do not mean the same thing or process, however they are properties that explain
the solubility salts or substances. The solubilities of substances in one another vary widely. We
might find a substance being miscible in one solvent but nearly insoluble in another. As a general
rule, one might find that “like dissolves like”. That is, similar substances or substances with
similar forces of attraction dissolve in one another. For example, diesel is miscible in gasoline
(petrol) because both are mixtures of hydrocarbon substances. On the other hand, diesel does not
mix with water because water is a polar substance, whereas diesel and indeed all hydrocarbons are
non-polar
QUESTION: Why do similar substances dissolve in one another to greater extents than do
dissimilar substances?
The two major factors that explain the solubility of one substance in another is the natural
tendency of substances to mix and it is sometimes also be referred to as the natural tendency
toward disorder. If the process of dissolving one substance in another involved nothing more than
simple mixing, we would expect substances to be completely soluble or miscible in one another.
However, in most cases, substances have limited solubility in one another and the factor that limits
solubility of substances is the relative forces of attraction between species (i.e molecules and
ions)
Conditions require to achieve before substances mix together
The solubility of a solute in a solvent (that is, the extent of the mixing of the solute and solvent
specie) depends on a balance between the natural tendency for the solute sand solvent species
to mix and the tendency for a system to have the lowest energy possible.
Solubility of molecular substances
For molecular substances, or substances like one gas dissolving in another or one liquid like
ethanol dissolving in water or kerosene dissolving in petrol, the only factor of importance that
accounts for their solubility is the natural tendency for the molecules to mix as a result of
having similar or nearly equal intermolecular forces of attraction; succinctly expresses that
“like dissolves like”. Substances with similar intermolecular attraction forces usually mix or are
soluble in one another.
Solubility of ionic substances
Ionic substances differ markedly in their solubilities in water. In most cases, these differences in
solubility can be explained in terms of the different energies between the ions in the crystal and
between ions in water (i.e energy of hydration and lattice energy). The energy of attraction
between an ion and a water molecule is due to an ion-dipole force. Water molecules are polar,
and so they tend to orient with respect to nearby ions. When ionic substances are dissolved in
water, the water molecules orient with their oxygen atoms (or hydroxide ions) the negative ends
of the molecular dipoles towards the positively charged ions of the substances while the molecular
orient with their hydrogen atoms (the positive end of the molecular dipole) towards the negatively
charged ions of the ionic substance.
Definition of hydration: Hydration is defined as the attraction of ions for water molecules.
Hydration of ions favours the dissolving of an ionic solid in water as the ions on the surface become
hydrated and then move into the body of the solution as hydrated ions. Again, the solubility of an
ionic solid depends not only on the energy of hydration of ions (i.e energy associated with the
attraction between ions and water molecules) but also on lattice energy; the energy holding ions
together in the crystal lattice.
Lattice energies depend on the charges on the ions (as well as on the distance between the centers
of neighbouring positive and negative ions.). The greater the magnitude of ion charge, the
greater the lattice energy and for this reason, substances with singly charged ions are
comparatively soluble and those with multiply charged ions are less soluble. This explains
why compounds with charged ions like, K+, Na+, NH4+, Cl-, NO3- etc are generally soluble while
those with phosphate ion PO43-, for example are generally insoluble. Lattice energy is also
proportional to the distance between neighbouring ions and the distance depends on the sum of the
radii of the ions. For example, the solubility of these compounds, Mg(OH)2, Ca(OH)2, Sr(OH)2
and Ba(OH)2 increases from Mg(OH)2 to Ba(OH)2 and decreases from MgSO2 to BaSO4.
Hydrates/hydrated substances
Definition: A hydrate is a compound that contains water molecule(s) weakly bound in its
crystals. Hydrated substances are often obtained by evaporation an aqueous solution of the
compound. The formula of hydrates is written with dot at the center to separate the compound and
the water molecules shown below; CuSO4.5H2O. When an aqueous of copper (II)
Tetraoxosulphate (VI) is evaporated, blue crystals form in which each formula unit the CuSO4 is
associated with five molecules of water. However, when the blue crystals are heated, the water
are driven off, and white crystals of CuSO4 called anhydrous are obtained.
Hydrates are named from the anhydrous compound followed by the word hydrate with a prefix to
indicate the number of water molecules (i.e of crystallization) per formula unit of the compound.
Examples are;
(i) Magnesium tetraoxosulphate (VI) heptahydrate - MgSO4.7H2O.
ii) Sodium Trioxocarbonate (IV) decahydrate – Na2CO3.10H2O.
(i) Calcium tetraoxosulphate (VI) dihydrate - CaSO4.2H2O
Exercise
(i) Give the IUPAC name of the following salts; K2CO3.10H2O; Na2SO4.10H2O;
HYDROLYSIS
Many salt solutions are not neutral or do not have pH’s of exactly 7. These solutions are either
basic (alkaline) or acidic because the ions of the solute react with water molecules. Anions
derived from weak acids (conjugate base) react with water to increase the concentration of
hydroxide ion (OH-) in the solution. For example; acetate ion reacts with water as;
C2H3O2- + H2O → HC2H3O2 + OH-. This reaction of acetate ion is similar to that of any other
weak baes or weak acids anions such as NH3, CO32-, SO32-, PO43-, CN- etc.
(i) NH3(l) + H2O(l) → NH4+(aq) + OH-(aq);
(ii) CO32- + H2O → HCO3+ + OH-;
(ii) SO32- + H2O → HSO3+ + OH(iv) CN-(aq) + H2O(l) → HCN(aq) + OH-(aq).
According to Bronsted-Lowry point of view, these ions NH3, CO32-, SO32-, PO43-, CN- act as a
base because they accept a proton from water. The reaction of each of the ions with water as shown
is called hydrolysis of each of the ions. These ions, CO32-, SO32-, and, CN- react with water or are
hydrolyzed to form conjugate acids and OH-.
Definition of hydrolysis: The hydrolysis of an ion is defined as the reaction of an ion with
water to produce the conjugate acid and hydroxide ion or the conjugate base and hydrogen
ion.
On, hydrolysis of ions, two (2) important considerations are; (i) Prediction of whether a salt
solution is acidic, basic or neutral and (ii) Calculations of the concentration of H+ and OH- ion the
salt solution (or equivalent).
Prediction of whether a salt solution is acidic, basic or neutral
Using CN- as an example, when the ion reacts with water or hydrolyzes, it produces the conjugate
acid HCN, hydrogen cyanide which is a weak acid as shown above. It means that the HCN tends
to hold on to the proton strongly (i.e it does not ionize readily). In order words, the cyanide CN-,
tends to pick up a proton easily, so it acts as a base. This explanation can be generalized: the
anions of weak acids are basic whereas the anions of strong acids have hardly any basic
character; that is, these ions do not hydrolyze. For example, the Cl- which is a conjugate base
to the strong acid HCl, shows no appreciable reaction with water as shown thus;
Cl-(aq) + H2O(l) → No reaction.
Meanwhile, the cation conjugate to a weak base such as NH4+ behaves like an acid. The cations
of weak bases are acidic. On the other hand, the cations of strong bases (metal ions of group
IA and IIA elements except Be) have hardly any acidic character; that, is these ions do not
hydrolyze as shown;
Na+(aq) + H2O(l) → No reaction
From the above observations; the following rules can be used in deciding whether a salt solution
will be neutral, acidic, or basic. It should also be noted that, these are only applied to normal salts
(i.e those salts in which the anion has no basic hydroxide ions or acidic hydrogen ions/atoms).
(i) A salt of a strong acid and strong base has no hydrolysable ions and so give a neutral
aqueous solution. Examples are; NaCl, KCl, LiCl, CsCl, etc.
(ii) A salt of a strong base and a weak acid, the anion of the salt is the conjugate of the weak
acid and it hydrolyzes to give a basic solution. Examples are; Na2CO3, K2CO3, NaCN, etc.
(iii) A salt of weak base and a strong acid, the cation of the salt is the conjugate of the weak
base and it hydrolyzes to give an acidic solution. Examples are; NH4Cl, (NH4)2SO4, MgCl2, AlCl3
etc.
(iv) A salt of a weak acid and a weak base, both ions hydrolyze and whether the solution is
acidic or basic depends on the relative acid-base strength of the two ions. Examples are;
ammonium formate-NH4CHO2 is slightly acidic,
Exercise
Copy and complete the table by stating the acid and base and whether the salt solution is acidic,
basic or neutral.
Salts
NaF
Zn(NO3)2
NH4CN
Al(NO3)2.
Na3PO4.
Acid and base that the salt
Nature of solution formed
Hydrolysis equilibrium of ions (Ka and Kb)
Hydrolysis equilibrium occurs only when an ion can form a molecule of a weak electrolyte
in the reaction with water. Strong acids and bases do not exist as molecules in aqueous solution
and so their cations and anions do not hydrolyze. It is only anions derived from weak acids and
cations derived from with bases that hydrolyze.
Problems involving hydrolysis equilibrium of ions are solved by considering the equilibrium
constant Ka and Kb of the conjugate acid-base pairs where Ka is the ionization constant of acid
and the Kb is equilibrium constant for the hydrolysis of an ion. Considering the acid ionization
of HCN and the base ionization of CN-. When the two reactions are added, we obtained the
ionization of water as shown
H+(aq) + CN-(aq) → Ka
HCN(aq)
CN-(aq) + H2O(l)
H2O(l)
HCN(aq) + OH-(aq) → Kb
H+(aq) + OH-(aq)
→ Kw
Where the equilibrium constant for each reaction is shown symbolically at the right. It should be
noted that when two reactions are added, their equilibrium constant are multiplied as;
Ka x Kb = K w
This relationship is general and it shows that the product of acid-base ionization constant in
aqueous solution for conjugate acid-base pairs is equal to the ion-product constant for water kw.
Determination of ion-product kw of water
Self-ionization of water; solutions of a strong acid or base: Although water is often considered
as non-electrolyte (i.e non-conductor of electricity), precise measurements do show a very small
conduction. This conduction results from self-ionization (or auto-ionization), a reaction in
which two like molecules react to give ions. In this case of water, a proton from one H2O
molecule is transferred to another H2O, leaving behind OH- ion and forming hydroxonium ion
H3O+(aq);
H2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
The equilibrium constant kw is given as; kc = [H3O+] [ OH-]
[H2O]2
The value of kc at room temperature is 3.2 x 10-18. This is because the concentration of ion
formed is very small, and the concentration of water remains essentially constant, about 56M at
25oC. Rearranging the above equation gives; [H2O]2 x kc = [H3O+] [ OH-]. The ion-product
constant of water is written as Kw and it has a value of 1.0 x 10-14 where Kw varies with temperature.
At a normal human body temperature (37oC), Kw equals 2.5 x 10-14. At 25oC, the ion-product of
water can be written as; Kw = [H3O+] [ OH-] = 1.0 x 10-14.
Then; using Kw, we can calculate the concentration of H+ and OH- ions in a pure sample of water.
Let x = [ H+] = [ OH-], and
kw = [ H+] [ OH-]. Then 1.0 x 10-14 = x2.
Therefore, x = 1.0 x 10-7.
Thus; the concentration of H+ and OH- are both 1.0 x 10-7 M in a pure sample of water.
Example 1: Calculate the; (a) Kb for CN-(aq) at 25oC whose conjugate acid of CN- of HCN has ka
value of 4.90 x 10-10 is and (b) Ka for NH4+(aq) at 25oC.
Solution
(a)
Kb =
Kw = 1.0 x 10-14
= 2.0 x 10-5.
Ka
4.90 x 10-10
NOTE: It should be noted, the Kb of CN- is approximately equal to the Kb of ammonium ion
(NH4+) which equals to (1.8 x 10-5). The base strength of CN- is comparable with that of ammonia.
(b) The conjugate base of NH4+ is NH3, whose Kb is1.8 x 10-5.
Therefore, Ka = Kw = 1.0 x 10-14
Kb
1.8 x 10-5
= 5. 6 x 10-10.
Example 2: Calculate the Kb of acetate ion C2H3O2-, if the Ka of the acetic acid is 1.8 x 10-5.
Solution
C2H3O2- + H2O(l) = HC2H3O2 + OH-(aq)
Kb = Kw = 1.0 x 10-14
Ka
1.8 x 10-5
= 5. 6 x 10-10.
Exercise
1. Calculate; (a) the Kb for F- if the Ka of its acid is 1.96 x 10-6 and (b) Ka for a conjugate acid of
aniline, C6H5NH+, if its Kb is 1.79 x 10-6.
Calculation of pH of weak acids
Example 1: Calculate the pH of 0.10M solution of NaC2H3O2, if the ka of this e conjugate base of
the acetate ion 5.6 x 10-10.
Solution
NaC2H3O2 → Na+(aq) + C2H3O2-.
C2H3O2- + H2O(l) → HC2H3O2 + OH-(aq).
Ka = [HC2H3O2] [OH-] = 5.6 x 10-10 = x2 → x2 = 0.1 x 5.6 x 10-10
[C2H3O2-]
0.1
2
-11
-6
x = 5.6 x 10 and x = 7.5 x 10
So; [HC2H3O2] = [OH-] = 7.5 x 10-6
pOH = -log (7.5 x 10-6) = 5.12
Then using; pOH + pH = 14; pH = 14 – pOH and pH = 14 – 5.12 = 8.88.
Example 2: Calculate the pH of sodium nicotinate eat 25oC, if the Ka for nicotinate acid was
determined to be 1.4 x10-5 at 25oC.
Solution
The salt ionizes to give Na+ and nicotinate ion in solution but only the nicotinate ion is
hydrolyzed.as; Nic- + H2O(l) → HNic(aq) + OH-(aq).
Since the nicotinate ion acts base, the concentration of species in the solution can be calculated
where kb is for the nicotinate ion. Thus; kakb = kw.
Kb = Kw = 1.0 x 10-14 = 7. 1 x 10-10.
Ka
1.4 x 10-5
The equilibrium concentration is obtained as thus; Let x = [HNic] [OH-]
Kb = [HNic] [OH-] = 7.1 x 10-10 = x2 → x2 = 0.1 x 7.1 x 10-10
[Nic-]
0.1
x2 = 7.1 x 10-11 and x = 8.43 x 10-6
So;
[HNic] = [OH-] = 8.43 x 10-6
pOH = -log (8.43 x 10-6) = 6. 07
Then using; pOH + pH = 14;
pH = 14 – pOH and pH = 14 – 6.07 = 7.93
The solution has a pH greater than seven.
Exercise
1. Calculate the pH of a 0.025M aqueous solution of sodium propionate, NaC3H5O2. What
is the concentration of propionic acid in the solution?
2. Calculate the concentration of hydroxide ion in a 0.080M aqueous solution of
methylamine CH3NH2, hence calculate the pH of the solution.
3. Calculate the pH of a 0.30M solution of ammonium chloride (NH4Cl).
4. Calculate the pH of 0.050M solution of sodium acetate if the Kb value in the solution is 5.
56 x 10-10.
5. The pH of a 0.15M solution of NaX is 9.77. Calculate the ionization constant of the weak
acid HX.
Hydrolysis of strong acids and base
(a) Strong acids: Strong acids such as hydrochloric acid (HCl) can react with water as shown;
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq). Also, the acid can on its own ionizes as thus;
HCl(aq) → H+(aq) + Cl-(aq). For 0.10M HCl, it has a concentration of H+ equal to 0.10M. To calculate
the concentration of OH- in 0.10M HCl, use the relationship Kw = [H+] [OH-] and substitute [H+]
= 0.10M into the equilibrium equation for Kw as;
Kw = [H+][OH-]
1.0 x 10-14 = 0.10 x [OH-]; [OH-] = 1.0 x 10-14. = 1.0 x 10-13.
0.10
Thus; the OH- concentration is 1.0 x 10-13.
For strong bases such as 0.01M NaOH, the concentration of OH- in the solution is 0.01M. The
hydrogen ion [H+], is produced by the self-ionization of water. So, the concentration H+ is
obtained using the equilibrium equation for Kw for 25oC as;
2.0 Kw = [H+][OH-];
1.0 x 10-14 = [H+] x 0.01; and [H+] = 1.0 x 10-14. = 1.0 x 10-12.
0.010
+
-12
Thus; the H concentration is 1.0 x 10 .
Exercise
1. Calculate the concentration of hydrogen ion and hydroxide ion at 25oC in; (a) 0.15MHNO3
and (b) Ca(OH)2.
2. A solution of barium hydroxide at 25oC is 0.125M Ba(OH)2. What is the concentration of H+
and OH- in the solution?
3. A solution has a hydroxide ion concentration of 1.0 x 10-5M at 25oC. Is the solution acidic or
basic?
Strong and weak electrolytes
The conductivity of aqueous solution such as NaCl is explained by the ionic theory of solution
proposed in 1884 by the Swedish chemist Svante Arrhenius (1859 - 1927). According to the
theory, an electrolyte produces ions when it dissolves in water. Most soluble ionic substances
dissolve in water as ions and are therefore electrolytes. Also, some molecular substances dissolve
to produce ions, as hydrogen chloride gas (HCl) react with water to give the ions H3O+ and Clbecause the aqueous solution of the hydrogen chloride called hydrochloric acid contains ions, the
hydrochloric acid is an electrolyte. HCl + H2O(l) → H3O+(aq) + Cl-(aq).
Definition of an electrolyte: An electrolyte is defined as an aqueous solution or molten form of a
compound the dissociates or decomposes into ions and conducts electricity when electric current
is allowed to pass through it.
Types of electrolytes
When electrolytes dissolve in water they produce ions but they do so in varying extents. Based on
this fact, electrolytes of two types namely; (i) Strong electrolytes and (ii) Weak electrolytes.
(a) Strong electrolytes: A strong electrolyte is an electrolyte that exists in solution almost
entirely as ions. Or they are electrolytes that ionize completely in aqueous solution or they
are electrolytes that undergo complete ionization in aqueous solution.
(a) HNO3(aq)
H+(aq) + NO3-(aq). (b) H2SO4(aq)
2H+(aq) + SO42-(aq).
Examples of strong electrolytes are;
(i) Strong acids e.g HCl, HNO3, H2SO4, HI, HBr, HClO4 etc.
(ii) Strong bases: E.g NaOH, LiOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 etc.
(ii) Aqueous solution of salts: E.g KCl, NaCl, MgCl2, CaCl2, AlCl3, etc.
Weak electrolytes
Definition of weak electrolytes: A weak electrolyte is an electrolyte that dissolves in water to
produce a relatively small percentage of ions. Or a weak electrolyte is an electrolyte that
ionizes slightly or that undergoes partial ionization in aqueous solution to produce low
concentration of ions to effect the flow electric current.
(i) CH3COOH(aq)
CH3COO-(aq) + H+(aq). (ii) H2CO3(aq)
H+(aq) + HCO3-(aq).
Examples of weak electrolytes are; (i) Weak acids – e.g HCOOH, H3PO4, H2SO3, HNO2, etc.
(ii) Weak bases –e.g NH4OH, Al(OH)3, Fe(OH)2, Cu(OH)2, Zn(OH)2, Fe(OH)3, etc