Chem

Introduction Simplified version of the modern periodic table is shown below:

Period 1 2 3 4

1 H* LI Na K

2

3

4

5

6

7

Be Mg Ca

B Al

C Si

N P

O S

F Cl Br

8 He Ne Ar Kr

*- Hydrogen displays properties of metals and non-metals. Trends in the Periodic Table As stated before the periodic table is arranged in order of increasing atomic number. The atomic number increases from left to right and from top to bottom. Elements in the same row (horizontal) are said to be in the same period and elements in the same column (vertical) are said to be in the same group. The group number corresponds to the number of electrons in the outer shell of the atom, so for example, a group 5 element possesses 5 electrons in its outer shell. Since we know that elements are arranged in groups and group number corresponds to the number of electrons in the outer shell, we can safely say all group 5 elements have 5 electrons in their outer shell. Elements found in the first 3 groups of the periodic table are metals, those found in groups 4-7 are non-metals. Those found in group 8 are the Nobel gases. The Nobel gases are stable and have a complete outer shell of electrons, as a result they are found in their elementary form. In other words, these elements do not react naturally to form compounds. Other elements try to achieve this stability by gaining, losing or sharing electrons from other elements via a process called chemical bonding bonding for short . For example Sodium loses 1 electron and Chlorine gains 1 electron to form sodium chloride. E.g. Sodium + Chlorine → Sodium Chloride Bonding is explained in some detail in the next section. It should be noted however that metals, like sodium above, bond by losing electrons to non-metals, like Chlorine above. Non-metals gain electrons when bonding with metals; however, they may also bond by sharing electrons with other non-metals to complete their outer shell. We already know from earlier, that elements in the same group possess the same amount of outer electrons and as such their reactivity is closely related. We will show examples of this by examining groups 1, 2 and 7.

Group 1 Metals This is a group of highly reactive metals also called the alkali metal series, and contains the metals Lithium, Sodium, Potassium, Rubidium, Cesium and Francium. Hydrogen is also a member of this group however, it shares both the properties of a group 1 metal and a group 7 non-metals. These metals possess 1 electron in their out electron shell as shown below:

The electronic configurations for the elements in group 1 are shown below:

Elements Li – Lithium Na – Sodium K – Potassium

Electronic configuration 2,1 2,8,1 2,8,8,1

As stated before, all elements try to achieve the stable electronic configuration of the closest Nobel gases (filled outer shells) . Since the elements in this group posses 1 electron in their outer shell, they will achieve stability by losing its single electron to a non-metal during reactions. For example, Sodium has an electronic configuration of 2,8,1. The closest noble gas to sodium is neon which has a electronic configuration of 2,8. Sodium will thus lose its single electron to a non-metal to achieve this stability. Please refer to the section on Atomic Structure if you are having difficulty understanding electronic configuration. Reactivity increases with increased atomic size as it is easier to remove the outer electron with increased atomic size. This is so because the outer electrons are further away from the nucleus and there is a resulting loss of electrostatic forces of attraction which exists between the negatively charged electrons and the positively charged nucleus. With this in mind the order of reactivity is as a follows:

Li