Chemistry Form Two Notes – Periodic Classification

These Chemistry Form Two notes cover Periodic Classification. The material is organised in clear sections for easier reading, understanding and revision.

Periodic Classification

Constructing the modem 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. The Law of Triads

In 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”, E.g.

The following are examples of Dobereiner’s triads:(Lithium, Sodium and Potassium)(Calcium,

Strontium and Barium)(Chlorine, Bromine and Iodine) and(Iron, Cobalt and Nickel) The Law of Octaves

In 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” This finding was comparable to the 8th note of music, hence the use of the word "octave"

This 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. The Periodic Law

Dmitri Mendeleev was bom 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 mass:

Mendeleev 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 6.1: Mendeleev’s short form of the Periodic Table

tr m |v VI AB AB

• Be Cc
• Mg
• Ca
• Zn As
• Sr
• Cd Sb
• Ba Ta
• Hg Bi
• Ro Ra Pa

The 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.

There were 7 horizontal rows which he called periods. All vacant positions in the table stood for

new elements yet to be discovered. Usefulness of Mendeleev's classification

The table summarized a large amount of information about the elements based on their chemical

properties.

The 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.

Limitations of Mendeleev's cl:

In 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 mass. 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, ete

The 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

The 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

Figure 6.1: The Moder Periodic Table

Wa Wa Va _Wa_Yia

We v3 YD YE WE

H- ges liquid

Nor-mtataic BP Treesition nerats Rave Eatn metals Halogens

The 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

horizontal rows.

The vertical columns in the table are called Groups, numbered I, II, III, 1V, 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.

The 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 Main features of the Modern Periodic Table

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

The 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.

Elements 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

Tare gases" because they were rarely found

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

Metals 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.

non-metals

reactive metals

transition metals

poor metals

Figure 6.2: Blocks of the periodic table

Periodicity

General Trends

This refers to change in some properties of elements across the periods and down the groups in the periodic table. These trends become more obvious if we leave aside the noble gases in Group 0. In this case, we shall concentrate our efforts on variations in the most important properties of the elements only. The following is a summary of the change in some properties of elements

down the groups and across the periods.

The Change in Properties of Elements Across the Periods

Atomic and ionic size

The sizes of atoms and ions may be given in terms of atomic radius and ionic radius units respectively. The number of shells an atom or ion posses and the nuclear charge determines the

ze of an atom or ion. This is how the two properties vary along the period and down the group:

Along the period: Considering the normal elements only, the size of the atoms decrease from left to right across the period. This is because as atomic number increases across the period, the

nuclear charge (due to increasing protons) increases and electrons in shells are pulled closer to

the nucleus

Tonic size

a.

Positive ions (cations):Across the period; The ionic size does not change, i.e. remains the same,

as you move across the period from either direction.

Negative ions (anions):A negative ion is larger compared to the corresponding neutral atom because on forming an ion, one or more electrons are added to the atom. The added electron(s) is/are repelled by the electron(s) already present in the outermost shell, hence leading to an increase in the size of an atom, even though no new shell is formed.Down the group and along the period: \onic size increases down the group, and along the period, i.e. from left to right Atomic radii (singular: radius)

Along the period: In the period, atomic radii decrease from left to right with increase in the

atomic number.

Electronegativity Electronegativity is the tendency of an atom to attract the shared pair of electrons towards itself

in a molecule. The electronegativity values of elements in group 0 (inert gases) is zero.

Along the period: Electronegativity increases while moving across the period from left to right in

the periodic table.

Metallic character (or electropositivity)

Electropositivity is the tendency of an element to lose the valency electron(s) and donate the same to other elements (usually non-metallic elements). This process occurs during the formation of new substances e.g. molecules and compounds. Literally, such reactions occur between metals and non-metals whereby metals donate electrons and non-metals receive these

electrons. So, metals are electropositive elements while non- metals are electronegative elements.

Along the period: Generally, metallic character decreases along the period from left to right.The gradation in metallic properties across the period is as follows: Metals —> poor metals —

metalloids + non-metals — noble ga

Chemical reactivity

Reactivity is the tendency of an element to lose or gain electrons in a chemical reaction.

Along the period: For metals, the reactivity decreases from left to right in a period while it

increases for non-metals.

Ionization Energy or Ionization Potential (I.E or I.P) This refers to the minimum amount of energy required to remove the most loosely bound electron from an isolated atom or ion in its gaseous state. The smaller the value of ionization

energy, the easier it is to remove the electron from the atom.M(g) +M+(g) + e-

Along the period: \t increases along the period from left to right with the increase in atomic

number.

Electron affinity (Ea): This is just opposite to LE. It is defined as the amount of energy released when an extra electron

is added to an isolated neutral atom in its gaseous state. Along the period: The value increases along the period from left to right

Density and melting point The density of a substance is the ratio of its mass to its volume, while the melting point is the

temperature at which a solid substance turns into liquid at standard atmospheric pressure. Density-Across the period: Densities decrease across the period from left to right

Meting point-Across the period: Melting points of elements decrease across the period from

left to right

The Change in Properties of Elements Down the Groups

Atomic and ionic size

Atomic size-Down the group: Atomic size increases as you move down the group.

Ionic size- Positive ions (cations)-Down the group: On descending the group, the nuclear charge increases and the number of shells increase by one at each step so, the ionic size also increases. A positive ion is smaller than the corresponding neutral atom because on forming the ion, the metal atom loses both the valency electron(s) and the outermost shell. Valency electron(s) refer(s) to the electron(s) in the outer-most shell of an atom. Any further removal of

electron(s) from the ion will decrease the ionic size further.Negative ions (anions)-Down the

group and along the period: Ionic size increases down the group, and along the period, ie from left to right.

Atomic radii (singular: radius)

Atomic radius is the distance from the centre of the nucleus to the outermost shell (valen shell). Down the group: Atomic radii of elements increase down the group with increase in

atomic size.

Electronegativity

Down the group: Electronegativity decreases while moving downwards in a group.

Metallic character (or electropositivity)

Down the group: Metallic character (electropositivity) increases down the group

Ionization Energy or Ionization Potential (I.E or IP)

Down the group: It decreases gradually down the group.

Why is there a decrease in LE as you go down the group? This is because electrons are held in their shells by their attraction to the positive nucleus, and as you go down the group, the size of the atom increases (increasing atomic radius). So, the outermost electron(s) of an atom gets further and further away from the attraction or pull of the positive nucleus, hence requiring little

energy to remove from the atom.

Electron affinity (Ea)

Down the group: The value of electron affinity decreases down the group. Density and melting point Density-Down the group: Densities of elements increase down the group

Meting point-Down the group: Melting points of elements decrease down the group as the elements become less metallic in nature. Electronic Configuration to Locate the Positions of Elements in Periodic Table

ration to locate the positions of elements in periodic table

The modem periodic table is based on electronic configurations of the elements. Look at table

6.3 and study the electronic configurations of the first twenty elements and where they are placed

in the periodic table.

Beryllium, magnesium and calcium have two electrons in the outer shell. These elements are in

Group 2

This pattern continues to Group 3, Group 4 and so on. The group number in the periodic table is the same as the number of electrons in the outermost shell. The halogens are the elements in Group 7. Bromine is one of the halogens. How many electrons does each bromine atom have in

its outer shell?

‘As we move down each group, the number of shells increases by one at each step. Each atom of

an element has one complete shell than the one above it.

As we move across each period, the outer shell is being filled by one electron at each step. ctronic configurations are found to be more stable than othe:

the end of each period have full outer shells. They have stable duplet (2 electrons) or octet (8

electrons) in their outermost shells. This makes them more difficult to break up, and this fits well

with the fact that they are so unreactive.

The outer electrons of an atom are mainly responsible for the chemical properties of an element.

Therefore, elements in the same group will have similar chemical properties