periodic table

What is Periodic Table and Periodicity?

A very important aspect of chemistry is the concept of the Periodic table and Periodicity.

But what exactly do we mean when we talk about the periodic table? The periodic table is the placement of elements in groups and periods according to their atomic numbers.

Elements are arranged in the periodic table by groups and periods. Groups of elements are determined by the valence electrons in their outermost shell while periods are determined by the number of shells the atoms possess.

There is something interesting about the periodic table, it can be used to predict the property of an element.

We have eight (8) groups and seven (7) periods in the periodic table. There are currently 118 elements in the periodic table.

The different groups in the periodic table have different properties and we will be analyzing these properties of these groups shortly.

There is no stress in calculating the oxidation numbers of the first twenty elements or getting their valencies due to the arrangements of elements in groups.

How to predict group and Period of Elements

The Group of elements is determined by the number of electrons in the outermost shell, this is called the valence electrons.

Example 1

In sodium atom, 11Na

The configuration is 2,8,1. One electron in the outermost shell

Therefore Sodium is in group 1

Example 2

In calcium atom, 20Ca

The configuration is 2,8,8,2, two electrons in the outermost shell

Therefore Calcium is in group 2

Example 3

In Chlorine atom, 17Cl

The configuration is 2,8,7

Therefore chlorine belongs to group 7

How about Periods?

How do you determine the period of elements?

Recall that period is the number of shells

So using the above examples

Like in example 1

In sodium atom 11Na

The shell configuration is 2, 8, 1 so the period is 3

Why? Because it contains three shells

In example 2

In Calcium atom 20Ca, so the period is 4

Finally in the same vein, in example 3

In Chlorine atom 17Cl

2,8,7, so the period is 3

Analysis of the Periodic Table

Actually there are 7 periods and 8 groupsthe period table (excluding the transition metals) while 7 periods and 18 groups if transition metals are grouped .

Remember that I mentioned that elements are placed according to groups in the periodic table.

So we are going to be looking at the various groups in the periodic

We mentioned before that the periodic table is arranged in groups and periods with the columns the group and the rows the period.

Also, recall that we have 7 periods and 8 groups in the periodic table.

Group 1 elements (Alkali metals)

They are called group 1 elements because they all have an electron in the outermost Shell.

Hydrogen is sometimes placed in group 1 simply because it has one valence electron nut that’s not a permanent location since it can also behave like group 7 elements that require one electron to attain stable configuration.

Why are group I elements called alkali metals?

Group 1 elements are called alkali metals on the periodic table because they dissolve or reacts violently with cold water to form alkali and liberate.

2Na + 2H2O==2NaOH + H2

2K+ 2H2O==2KOH+ H2

An alkali is a base that is soluble in water and only group 1 hydroxides and Ammonium hydroxide

The group 1 elements are lithium, Sodium, Potassium, Rubidium, Cesium, and Francium.

The Group 1 reactions are very explosive .and is usually very fast.The three observations usually associated with itare: effervescence of a hydrogen gas, darting or fast movement of the solid metal and dissolution of the metal.

7 Basic characteristics of Group 1 elements

what are the characteristics of Group Iof I metals?

The alkali metals also have some properties that are peculiar to them alone.

1.They combine with non metals to form salts.

  2K+Cl2 –> 2KCl

  2Na  – Cl2 –> 2NaCl

2.The carbonates of group I do not decompose

The Group I carbonates do not decompose becuase they are very stable to heat.

Examples of group I carbonates are potassium carbonate, sodium carbonate, lithium carbonate etc

3.The nitrates of group I elements decompose with difficulty to give their dioxonitrate III and oxygen.

The nitrates of group I elements decompose stingingly to release oxygen.

2KNO3 ==2KNO2 + O2

 2NaNO3 ==2NaNO2 + O2

4.Group I elements form alkalis when they react with cold water

The Group I elements form alkali when they dissolve in water . Example potassium can dissolve in water to form potassium hydroxide and liberate hydrogen

2K+ 2H2O===2KOH + H2

5. Group I elements are very soft

Group I elements are very soft; this is because of the weak metallic bonding caused by the presence of onlynone valence electron in the structure of metals.

Metallic bonding is caused by the attraction between the lattice positive ions and sea of electrons

6. Group I metals conduct electricity

Group I metals conduct electricity due to the presence of mobile electrons in the lattice.

7. The density of group I metals increase down the group

As you move down the group of alkali metals, the number of shells increases thus increasing the mass of the atom.

Density and mass is directly proportional to each other, so density increases down the group.

8. Reactivity of group I metals increase down the group.

In the same trend as density, reactivity of group I metals increase as you move down the group.

For example, Potassium can displace sodium from its salt

K + NaCl ====KCl + Na

K + NaBr ====KBr + Na

Sodium can displace lithium from its salt .

Na +LiiCl====NaCl + Li

Na + LiBr====NaBr + Li

Summary of the periodicity of Alkali metals

Periodic property Periodicity of alkali metals down the group
Density increases down the group
Soft nature Alkali metals become softer down the group
Melting point Melting point of alkali metals decreases
Reactivity or metallic character
Reactivity increases down the group
Conductivity Conductivity increases down the group
Atomic size Atomic size increases down the group

Group 2 elements (Alkaline-earth metals)

Group 2 elements show similar chemical properties because they have two electrons in their outermost Shell.

The group 2 metals are called alkaline earth metals because they dissolve in water to form alkaline solutions and because they are gotten from the earth surface.

Beryllium is not a typical alkaline earth metal because its oxides and hydroxides are amphoteric.

The group 2 elements are magnesium, Calcium, strontium, Barium, etc 

The group elements behave just like the group 1 elements but just that they are not as electropositive as the group I elements.

For example sodium Potassium and Lithium are more reactive than calcium,

The group two elements like Calcium dissolve in cold water but not as vigorously as group one metals.

What are the properties of group 2 alkaline earth metals?

There are basic characteristics of Group II alkaline earth metals

These are the facts about group 2 elements

1.They have low ionisation energy that decreases down the group

Generally, metals have low ionisation energy when compared to non metals.

ionisation energy is the energy required to remove the most loosely bound electrons from the outermost shell of a gaseous atom.

The reason why group II elements or metals in general have low ionisation energy is due to the small size though this ionisation energy increases as you move from metals to non metals.

2. They have low melting and boiling points

Generally metals have low melting and boiling points basically due to the weak metallic force of attraction.

The metallic bonding is of the four types of bonding caused by attraction between lattice positive ions and sea of electrons.

3. Reactivity of alkaline earth metals decrease down the group

Yes, reactivity of metals decrease as you move down the group. Magnesium is less reactive than calcium;and so is Calcium less reactive than Strontium.

For example Calcium can displace magnesium from its salt.

Ca + MgCl2 ====CaCl2 + Mg

Sr + CaBr =====SrBr2 + Ca

Also, calcium can react with cold water though slowly and very fast with steam but magnesium can only react with steam.

Ca + 2H2O –> Ca (OH)2 + H2

        Mg + 2H2O –> Mg (OH)2 + H2

4. They are good conductors of electricity

Due to the presence of mobile electrons , metals are good conductors of electricity .

Graphite and metals conduct electricity due to the presence of mobile electrons while molten or aqueous ionic compound conduct electricity by tge presence of mobile ions.

5. Alkaline earth metals are soft metals

Alkaline earth metals are soft metals compared to group 3 metals but group I metals maybe a little softer than group 2 metals because of a weaker netallic bond.

6. Density of alkaline earth metals decreases down the group

The densitof group 2 metals increase as you move down the group snd this is typically caused by the increase in mass number of elements as you go down .

7. Group II Oxides do not dissolve in water

Group II Oxides do not dissolve in water to form alkali and their Oxides are resistant to heat.

Overview of the Group II alkaline earth metals

Periodic propertiesTrend down the group
Reactivity Reactivity increases down the group
Density Density increases down the group
soft nature Alkaline earth metals become softer down the group
Atomic size Atomic size of group 2 elements increase down the group
Ionisation energy ionisation energy of group 2 elements decreases down the group
Melting and boiling point Since the strength of metallic bond decreases down the group( due to increase in atomic suze), there idecrease in melting and boiling points
Electronegativity Electronegativity decreases down the group

Group 13 elements (Boron Family)

group 3 elements are sometimes called the boron group or family while group 13 called the Scandium group.

while these groups are classified in terms of number , they really have some differences.

For example Scandium is actually a transition metals while Aluminium is a typical normal metal.

The group 3 elements include Boron, Aluminium, Gallium, Indium, and Titanium.

I I will like to say that there is another type of group 3 elements; group 13 as seen in the complete periodic table.

Group 13 elements are Scandium, Yttrium and Lutetium commonly exhibit similar chemical properties.

They are trivalent and have the potential to donate three electrons during bond formation.

Properties of group 13 elements

group 3 elements is a combination of both metalloids and metals.

Boron is an example of a non metal or it has basic properties if a non metal.

1. Group 3 elements are silvery white soft solid

The group 3 elements are silvery white soft solids but the hardness increases down the group.

2. Group 3 elements liberate hydrogen when they react with acids

2Al +6HCl ==2AlCl3 + H2

3.All the elements burn in oxygen to form ocude of the formula X2O3

Facts about Group 3 elements

  1. Boron do not usually react with water
  2. Aluminium forms a protective oxide layer
  3. Group 3 compounds are among the most abundant in nature

In group three, it is important to note that boron is a metalloid but aluminum us a metal.

However the most common element in group 3 is Aluminium,

Aluminum and its compounds (Al2O3  and Al(OH)3) are amphoteric

Amphoteric means it has a dual nature; it can behave like a base and an acid.

Al2O3  + 3H2SO4  –>     Al2(SO4)3 + 3H2O

Al2O3  + 2NaOH + 3H2O –>   2NaAl(OH)4

Group 14 elements

The group 4 elements include Carbon, Silicon, Germanium, Tin and Lead.

There is a special change in the trend of the character of elements in this group. As you move down the group, there is a progressive change from non-metallic character to metallic character.

Carbon and silicon form +4 valiancy and equally form giant structures.

The carbon family, Group 14 in the p-block, contains carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl).

Each of these elements has only two electrons in its outermost p orbital: each has the electron configuration ns2np2.

Group 5 elments

The group 5 elements are Nitrogen, Phosphorus, Arsenic, Antimony, and Bismuth.

The common ones are Nitrogen and Phosphorus with two stable valencies of 3 and 5.

Nitrogen and phosphorus form oxides and hydrides when they combine with oxygen and hydrogen respectively.

Group 6 elements

The group 6 elements are Oxygen, Sulphur, Selenium, and Polonium. The most common ones are Oxygen and Sulphur. They ionize by gaining one electron and are very reactive when compared to other groups

Group 7 elements (Halogens)

Group 7 Elements are called Halogens because they readily combine with metals to form salts.

Group 7 Elements are very electronegative and can attain stability easily by gaining one electron in order to complete their shells.

There is a trend in the reactivity of elements

For example, Flourine is the brightest among them.

Fluorine –yellow, Chlorine –greenish-yellow, Bromine-Iodine Iodine-black

Their states also differ, for example, Fluorine and Chlorine are gases while Bromine is liquid and Iodine is solid.

Group 8/group 0 elements –Noble gases

Group 8 elements (Noble gases)

The group 8 elements are Neon, Argon, Krypton, etc

Group 8 Elements are called noble gases because they have zero valencies and this is because they have stable configurations.

They hardy combine because they have full outer filled configurations and as such cannot lose or gain electrons.

I think I have done a little about the groups of elements but I will talk about the periodicity of elements.

The Transition Metals (Between group 2 and 3)

What are transition metals? Transition metals are a special group of metals that exist between group two and group three. This is because they have multiple valencies.

Transition metals have special properties like variable oxidation states, formation of colored compounds, catalytic abilities, and ability to exhibit paramagnetism.

Property of Transition Elements. Reason behind the Properties
Transition metals are paramagnetic
This is due to the unpaired elections in the 3d orbital
Variable oxidation states
They have multiple valencies
This is due to the transition in the d orbital
Catalytic ability
similar atomic size that easily mix with other metals
Formation of coloured compounds or ions Due to the presence of vacant d-orbitals from the d-d transition of electrons which causes the colo
This happens by the absorption of visible light radiation, which promotes an electron from one d-orbital to another.

Periodicity of Elements

What is Periodicity? Periodicity is defined as the trend of the properties of elements across the period and down the group.

When I talk about the trend I mean the way the properties vary across the period and down the group.

What are the properties we talk about?

By periodic properties we mean

  1. Ionization energy
  2. Atomic size
  3. Electron affinity
  4. Metallic character /electropositivity
  5. Non-metallic character/electronegativity
  6. Melting and boiling points
  7. Thermal and Electrical conductivity

A quick Overview of the Periodic Properties

These periodic properties are governed by parameters like screening effect, nuclear charge and distance of electrons from the nucleus.

I’ll do a quick summary of all the periodic properties in a jiffy.

Please take note that the periodic properties vary from left to right across the period and from top to bottom down the group.

Ionization energy

Ionization energy is the energy required to remove the most loosely bound electron from the outermost shell of an atom.

Ionization decreases from top to bottom down the group and increases from left to right across the period.

Atomic size /Atomic radius

Atomic size is one-half the distance of the closest approach between the two atoms in an elemental substance.

Atomic size generally decreases across the period from left to right and increases down the group. Attractions. his is due to the increase in nuclear charge

Electron affinity

Electron affinity is the energy released when an atom gains an electron.

Electron affinity increases across the period and decreases down the group.

Metallic character/electropositivity

Electropositivity is the power of an atom to give out electrons.

Electropositivity generally decreases across the period and increases down the group.

Non-metallic character / Electronegativity

Electronegativity is the power of an atom to gain electrons or attract electrons to itself.

Electronegativity increases across the period and decreases down the group.

Fluorine is the most electronegative element.

Melting and boiling point

The trend or gradation of melting and boiling point is not uniform like the other properties.

Melting and boiling point increases from left to right for metals, actually, it is highest in group 4 and decreases henceforth for non-metals.

The reverse is the case down the group; it decreases down the group for metals and increases for non-metals.

Summary of the periodic properties

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