TOPIC 5: ATOMIC STRUCTURE
thought that atoms were solid, indivisible particles. But, as a result
of work done mainly by Lord Rutherford, the idea has been greatly
changed in recent years. According to Rutherford, the atom consists of 3
kinds of particles – protons, neutrons and electrons. These are called
centre of the atom is called nucleus. The nucleus contains a cluster of
two sorts of particles, protons and neutrons. Thenucleus is very small,
occupying only about 1% of the volume of an atom. The rest of the atom
is mostly empty space, withelectrons spread out in it.
move around the nucleus in special paths calledelectron shells
(orbits/or orbitals or energy levels). Protons andelectrons have
electric charges. Neutrons have no charges.All the particles in an atom
are very light. Their masses aremeasured in atomic mass units rather
than grams. The proton isa positively charged particle. Its mass is
about equal to that ofhydrogen atom. The neutron is has no charge, it is
neutral. Itsmass is about equal to that of hydrogen atom. The electron
isnegatively charged. Its charge is equal but opposite to the chargeon
the proton. It has a very small mass, about 1⁄1840 of the massof the
single atom is electrically neutral (it has no electrical charge).This
means that in any atom there must be equal numbers ofprotons and
electrons. In this way, the total positive charge onthe protons is
balanced by the total negative charge on theelectrons orbiting the
nucleus. So, the charges must cancel.
arrangement refers to the manner in which electronsare arranged in an
atom. An atom contains a central nucleuscontaining protons and neutrons,
and a cluster of electronsrevolving in orbits around the nucleus. These
electrons aregrouped in shells.
(1913) put forward a theory of electron positioning whichis still
generally accepted and used until now for chemicalpurposes. Bohr’s
Theory on the arrangement of electrons in anatom can be summarized as
- Electrons are in orbit around the nucleus of the atom.
electron orbits are grouped together in shells; a shellis a group of
orbits occupied by electrons withapproximately equal energy.
- The electrons in shells distant from the nucleus havehigher energy than those in shells close to the nucleus.
fill the shells starting with the first shell, whichis closest to the
nucleus. Shells are numbered 1, 2, 3, 4,98etc. outwards from the
nucleus. The shells may berepresented by the letters K, L, M and N
respectivelystarting from the nucleus.
- The maximum possible number of electrons in a shellnumbered n is 2 2n .
first shell can only contain up to 2 electrons. Thesecond shell can
contain a maximum of 8 electrons. Thethird shell can contain up to 18
- In the outermost shell of any atom, the maximumnumber of electrons possible is 8.
- The outer electrons of some atoms can be removed fairlyeasily to form ions.
- Chemical bonding between atoms to form moleculesinvolves the electrons in the outer shell only.
arrangement of electrons around the nucleus is also knownas electronic
configuration. This arrangement depends on themaximum number of
electrons that can occupy a shell. An atomwith 13 electrons will have
the following electronicconfiguration: 2:8:3. This means that there are 2
electrons in thefirst shell, 8 electrons in the second shell and 3
electrons in the third shell.
the first 20 elements, the organization of the electrons becomes
increasingly complicated. The third shell (n = 3) can be occupied by a
maximum number of 18 electrons
this stage, you will not be asked to work out electronarrangements
beyond element 20 (calcium), but you should beable to understand the
electronic structures involving moreelectrons (for example bromine with
the arrangement 2:8:18:7).
atoms of one element have the same number of protons. Thisis called the
atomic number (or proton number) of that element.It is given by the
two elements can have the same atomic number. Sodiumatoms have 11
protons. This is what makes them different fromall other atoms. Only
sodium atoms have 11 protons, and anyatom with 11 protons must be sodium
the same way, an atom with 6 protons must be carbon atom.Also any atom
with 7 protons must be nitrogen atom. So, youidentify an atom by the
number of protons in it. There are 109elements altogether. Of these,
hydrogen has the smallest atoms,with only 1 proton each. Helium atoms
have 2 protons each.Lithium atoms have 3 protons each, and so on up to
meitneriumatoms, which have 109 protons each. Table 5.3 shows the
first20 elements arranged according to the number of protons theyhave.
atom has an equal number of protons and electrons, so theatomic number
also tells us the number of electrons in that atom.In any given atom of
an element, the number of neutrons has no effect on the identity and
properties of that particular element. It is the number of protons and
electrons that determine the identity and properties of any given
element. The number of neutrons only affects the mass, since each one of
them has the same mass as that of a proton.
alone do not make up all the mass of an atom. The neutrons in the
nucleus also contribute to the total mass. The mass of the electrons can
be regarded as so small that it can be ignored. As a proton and a
neutron have the same mass, the mass of a particular atom depends on the
total number of protons and neutrons present. This is called mass
number (or nucleon number). The mass number of an atom is found by
adding together the number of protons and neutrons. It is given by the
symbol A. Table 5.3 shows the mass number of the first 20 elements,
arranged in order of increasing atomic mass (mass number).
mass number = number of protons + neutrons in an atom. Sodium atom has
11 protons and 12 neutrons, so the mass number of sodium is 23. Since
the atomic number is the number of protons only, then:
number – atomic number = number of neutrons. So, for sodium atom, the
number of neutrons = (23-11) =12. You can also take into account the
fact that, because the number of protons is always equal to the number
of electrons, then the number of electrons in sodium atom is simply 11.
The same rule can be applied to work out the sub-atomic composition of
- Number of electrons = number of protons = atomic number
- Number of neutrons = mass number (A) – atomic number (Z).
of the same element may have different numbers of neutrons. In a normal
situation, atoms of the same element will have the same number of
neutrons. However, many cases occur in which two atoms of the same
element contain the same number of protons but different numbers of
neutrons. Having equal number of protons, these atoms must also have
equal numbers of electrons. However, the differing numbers of neutrons
cause the atoms to have different mass numbers. An element showing such
properties is said to show isotopy and the varieties of the atom are
called isotopes of the element.
isotopy can be defined as the tendency of atoms of one element to
posses the same atomic number but different mass numbers (atomic
masses). Isotopes can be defined as atoms of the same element with the
same number of protons but different numbers of neutrons, or as „atoms
of the same element with the same atomic number but different atomic
isotopes of an element have the same chemical properties because they
contain the same number of electrons. It is the number of electrons in
an atom that decides the way in which it forms bonds and reacts with
other atoms. However, some physical properties of the isotopes are
different. The masses of the atoms differ, and therefore other
properties, such as density and rate of diffusion, also vary.
isotopes (like tritium) are unstable. The extra neutrons in their
nuclei cause them to be unstable so that nuclei break spontaneously
(that is, without any extra energy being supplied), emitting certain
types of radiation. They are known as radioisotopes.
order to distinguish between different isotopes of the same element in
writing symbols and formulae, a simple system is adopted. The isotope of
an element, say X will have the symbol X,AZ , where A is the mass
number of the isotope and Z is the atomic number of any atom of X. Thus,
for all isotopes of one element, Z is constant, and A varies because
there are different numbers of neutrons in the different isotopes of the
element. For example, the three isotopes of carbon are expressed as
12C6, 13C6,and 14C6. Chlorine has two isotopes: 35Cl17 and 37Cl17 .
Since A represents the total number of neutrons and protons in the
nucleus of an atom (mass number/atomic mass), and because Z is the
number of protons (atomic number), then the number ofneutrons in the
nucleus of a given isotope is given by:Number of neutrons in the nucleus
= A – Z
we have seen, most elements exist naturally as isotopes.Therefore, the
value we use for the atomic mass of an element isan average mass. This
takes into account the proportions(abundance) of all the naturally
occurring isotopes. If aparticular isotope is present in high
proportion, it will make alarge contribution to the average.
sample of chlorine gas contains 75% of the isotope 35Cl17 and 25% of
the other isotope 37Cl17 . What is the relative atomic mass of chlorine?
average value for the masses of atoms of an element isknown as the
relative atomic mass (Ar).Therefore, the relative atomic mass of
chlorine is 35.5 (i.e., Ar =35.5).
the modern periodic table has been a major scientific achievement. The
first steps towards working out this table were taken long before anyone
had any idea about the structure of atoms. The number of elements
discovered increased steadily during the 19th century. Chemists began to
find out patterns in their properties.
1817, the German scientist Johann Dobereiner noticed that calcium,
strontium and barium had similar properties, and that the atomic weight
of strontium was halfway between the other two. He found the same
pattern with chlorine, bromine and iodine and also with lithium, sodium
and potassium. So, he put forward the law of Triads: “If elements are
arranged in groups of three in order of increasing atomic weights,
having similar properties, then the atomic weight of the middle element
is the arithmetic mean of the atomic weights of the other two elements”,
following are examples of Dobereiner’s triads:(Lithium, Sodium and
Potassium)(Calcium, Strontium and Barium)(Chlorine, Bromine and Iodine)
and(Iron, Cobalt and Nickel)
1863 John Newlands, an English chemist noted that there were many pairs
of similar elements. In each pair, the atomic weights differed by a
multiple of 8. So, he produced a table with the elements in order of
increasing atomic weights, and put forward the Law of Octaves: “If
elements are arranged in order of their increasing atomic weights, the
properties of the 8th element, starting from a given one, are a kind of
repetition of the first element”.
was the first table to show a periodic or repeating pattern of
properties. But it was not widely accepted because there were too many
inconsistencies. For example, he put copper and sodium in the same
group, even though have very different properties. Also iron was placed
in the same group as oxygen and sulphur.
Mendeleev was born in Siberia, Russia, in 1834. By the time he was 32,
he was a professor of Chemistry. In 1869 Mendeleev advanced the work
done by Newlands and contributed very useful new ideas. He began by
listing all the known elements in order of increasing atomic mass. He
spotted that elements with similar properties appear at regular
intervals or periods down the list. His findings were the basis for the
Periodic Law: “The properties of elements are a periodic function of
their atomic masses”.
placed similar elements into groups. He realized that not all elements
had been discovered. So he left gaps for new ones in the correct places
in his table. He also swapped the order of some elements to make them
fit better. He predicted the properties of the missing elements from the
properties of the elements above and below them in the table. He also
listed separately some elements which did not appear to fit into any
group i.e. iron, cobalt, nickel, etc.
table had 9 vertical columns which he called Groups. The groups were
numbered from 0 to 8. The elements in group 0 were not known by then,
but were discovered later on. Groups 1 to 7 were subdivided into A and B
subgroups. Group 0 included the transition elements. Noble gases were
later placed in group 0.
were 7 horizontal rows which he called periods. All vacant positions in
the table stood for new elements yet to be discovered.
- The table summarized a large amount of information about the elements based on their chemical properties.
table was very useful in predicting the existence and properties of
undiscovered elements, for which gaps had been left in the table.
- The table was also used in checking relative atomic masses of elements.
three cases, pairs of elements had to be included in one group based on
inverse order of their atomic weights so as to fit into groups of
elements having similar properties. These pairs were argon (39.9) and
potassium (39.1), cobalt (58.9) and nickel (58.9); plus tellurium
(127.5) and iodine (126.9). This difficulty was resolved when the basis
of classification was based on the atomic number instead of the atomic
- The elements that were placed in group VIII formed an incompatible mixture.
- The placing of two different families in one group e.g. K and Cu; Ca and Zn, etc.
periodic table is the chemists map. It helps you understand the
patterns in chemistry. Today we take it for granted. But it took
hundreds of years, and work of hundreds of chemists, to develop.
Modern Periodic Table is similar to that of Mendeleev, but contains
several improvements. Elements are arranged in order of atomic number
instead of atomic mass. This means that elements no longer have to swap
places to fit correctly. Many new elements have been discovered and
slotted into the spaces left by Mendeleev. Also metals and non-metals
are clearly separated. The Modern Periodic Table is shown in Figure 6.1.
long form of the periodic table is the commonly used form of the
periodic table. The elements in the table are arranged based on their
atomic weights, starting from hydrogen (1), helium (2), lithium (3),
beryllium (4) and so on. The elements appear in vertical columns and
vertical columns in the table are called Groups, numbered I, II, III,
IV, V, VI, VII and 0, which is also known as group VIII. Group I
contains the elements lithium (L), sodium (Na), rubidium (Rb), caesium
(Cs) and francium (Fr). Group II consists of elements starting from
sodium (Na) down to radium (Ra). Some of the groups have special names.
- Group I is often called the alkali metals.
- Group II the alkaline earth metals.
- Group VII the halogens.
- Group 0 the noble gases.
transition metals (or elements) form a separate block in the middle of
the periodic table between group II and III. The atoms of these elements
have more complicated electron arrangements. Note that the group
contains many common metals such as iron (Fe), Nickel (Ni), copper (Cu),
and Zinc (Zn). One of the interesting properties of these elements is
that they form coloured compounds.
- The elements in the table are placed in order of their atomic numbers instead of their atomic masses.
- There are a total of 18 groups and 7 periods.
- There are 5 blocks of similar elements in the periodic table as shown in figure 6.2.
normal (non-transition) elements (groups 1-7) have their outermost
shells incomplete, meaning that they can allow additional electrons to
enter into their outermost orbital (valency shell). But each of their
inner shells is complete.
- The transition metals have their outermost as well as their penultimate (second last) shells incomplete.
of group 0 (noble gases) have their shells complete. These elements
show little reactivity. That is why they wereonce called „inert‟ gases
because they are very unreactive; or „rare gases‟ because they were
- Gaps left by Mendeleev for undiscovered elements
(now occupied by the transition elements and the noble gases) have been
filled by the respective elements following their discovery. Man-made
elements have also found a place in the periodic table.
have been clearly separated from non-metals. Metalloids or semi metals
(poor metals) have also been included. Metalloids are elements whose
properties are intermediate between metals and non-metals. They include
boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb) and
tellurium (Te). In some publications, germanium and antimony are
usually classed as poor metals and the rest as non-metals.
CHEMISTRY FORM TWO ALL TOPICS.
CHEMISTRY FORM TWO TOPIC 1 & 2.
CHEMISTRY FORM TWO TOPIC 3 & 4.
CHEMISTRY FORM TWO TOPIC 5 & 6.
CHEMISTRY FORM TWO TOPIC 7.
CHEMISTRY STUDY NOTES, FORM FOUR.
CHEMISTRY STUDY NOTES, FORM THREE.
CHEMISTRY STUDY NOTES, FORM TWO.
CHEMISTRY STUDY NOTES FORM ONE.