Chap 6: The Periodic Table of Elements
The Periodic Table is an arrangement of elements in order of their increasing proton (atomic) number.
Groups:
The vertical columns of elements in the Periodic Table are called groups. There are eight groups in the periodic table (there are more than eight vertical columns, but the block of elements between group II and group III are not classified according to groups, so that we do not count the columns in that bock when counting groups). The group number of an element is determined by its number of valence electrons. For example, Sodium (Na) has 1 valence electron, so it falls in group I. The group number is generally written in roman numbers. The elements in the same group have similar chemical properties, as the chemical properties majorly depend upon the number of valence electrons. Their physical properties, however, are very different. We will learn about all this as the chapter proceeds.
Periods:
The horizontal rows of elements in the Periodic Table are called periods. There are seven periods in the Periodic Table. The period number indicates the number of shells each atom of an element contains. All the elements in one period have the same number of shells. For example, Potassium (K) and Calcium (Ca) both have 4 shells, so they both are placed in period 4.
From left to right across a period, there is a decrease in metallic properties and an increase in non-metallic properties.
Fig 6.1. The Periodic Table.
As you can see in the periodic table, different parts of the table are coloured differently. Let’s learn these classifications in detail.
Transition Metals:
The block of metals between group II and group III is known as transition metals. These are typical metals. They are strong and hard, good conductors of heat and electricity, and have high melting points. The section of the periodic table in which the transition metals are placed does not have any group numbers. There are some properties typical of all transition metals, and the other metals (the ones placed in groups) do not generally have these. E.g., the transition metals form coloured compounds. They are used as catalysts to speed up chemical reactions. They have variable oxidation states. You will learn more about oxidation states in the chapters of bonding, but I will give you a general idea.
You see, atoms of most elements can either loose or gain electrons, giving them a positive or negative charge. This charged atom is known as an ion. The oxidation state is the charge an atom of an element would have if it existed as an ion in a compound. For example, oxygen in its ion form, i.e. O2-, has the oxidation state -2. This state is fixed for many elements, but not for transition metals.
An example of transition metals is Iron (Fe). It forms coloured compounds. It can exist in form of two ions, Iron (II) and Iron (III). The number in the brackets refers to its oxidation state. Iron (II) compounds are green, and Iron (III) compounds are reddish-brown. Iron is used as a catalyst in the Haber Process (you will learn more about this process later). Other typical properties of metals are also true for this metal.
Groups I and II:
These two groups consist of metals. These are reactive metals, and their reactivity increases as you go down the group. Group I metals are also known as Alkali Metals, while Group II metals are called Alkaline Earth Metals. They are called this because they react with water to form an alkali and hydrogen gas. Alkali Metals are very reactive, even reacting with air and cold water. They are soft, and can be cut easily. When freshly cut, each of these elements shows a shiny and silvery surface that rapidly tarnishes in air. They also have low melting and boiling points, and low densities. Why then are they called metals, you will learn so in a later chapter. Examples of Group I metals are Sodium (Na) and Potassium (K). Examples of Group II metals are Magnesium (Mg) and Calcium (Ca).
Note: Hydrogen does NOT belong to Group I. It has 1 valence electron so it is generally placed in group I, but it does not have the properties typical of Group I elements.
Group VII:
These elements are called Halogens. The Halogens have low melting points (M.P) and boiling points (B.P). They are also coloured. The M.Ps and B.Ps of the Halogens increase as we go down the group, and their colour gets darker. Halogens react with most metals to form halides. Their ions are called halide ions. In a solution, a more reactive halogen replaces a less reactive halogen. You will learn more about this later, and can relate all later learnings back to this.
Examples of Halogens are Fluorine (F) and Chlorine (Cl). They, however, always exist as F2 and Cl2. This is a property typical of all halogens, i.e. they are diatomic (refer to the chapter of bonding).
Group VIII (also called Group 0):
These elements are called Noble Gases. They are also referred to as inert gases because they are unreactive. The noble gases:
- Are monatomic elements
- Are all colourless gases at room temperature
- Have low M.Ps and B.Ps that increase on going down the group
- Are insoluble in water.
Their atoms have full outer shells of electrons, therefore they do not need to lose, gain or share electrons, hence they are unreactive. They do not form ions or molecules.
Namely, the Noble Gases are Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn).