Periodic Classification of Elements

Science Part 1 • Class 10th

In this chapter

• Classification of elements
• Dobereiner’s Triads
• Newlands’ Law of Octaves
• Mendeleev’s Periodic Table
• Modern Periodic Law and Table
• Groups, periods and blocks
• Electronic configuration
• Periodic trends
• Metallic and nonmetallic character

Classification of Elements

Today, 118 elements are known to the scientific world.

Around the year 1800, only about 30 elements were known.

As more elements were discovered, information about their properties also increased.

Scientists started searching for a pattern in the properties of elements.

Early classification

• Initially, elements were classified as metals and nonmetals.
• Later, another class called metalloids was identified.
• Classification helps us study a large number of elements systematically and compare their properties.

Dobereiner’s Triads

• In 1817, German scientist Dobereiner proposed the concept of triads.
• He arranged three elements having similar chemical properties.
• The elements were arranged in increasing order of atomic mass.
• The atomic mass of the middle element was approximately equal to the mean of the atomic masses of the first and third elements.

Triads – Calculation

Example: Li – Na – K

Example: Ca – Sr – Ba

Example: Cl – Br – I

Limitation of Dobereiner’s Triads

• Dobereiner’s method was useful for some groups of elements.
• However, all known elements could not be classified into triads.
• Therefore, a more general classification was needed.

Newlands’ Law of Octaves

• English scientist John Newlands proposed a new method of classification.
• In 1866, he arranged elements in increasing order of their atomic masses.
• He started with hydrogen and ended with thorium.
• He observed that every eighth element had properties similar to the first element.
• He compared this repetition with the octaves in music.

Newlands – Examples and Limitations

Examples

These pairs showed similar properties.

Limitations

1. The law was applicable only up to calcium.
2. Newlands arranged the elements in a table of 7 × 8 = 56 boxes.
3. Some boxes had two elements.
4. Elements having different properties were sometimes placed in the same box.
5. There was no provision for newly discovered elements.
6. Newly discovered elements did not fit properly into the law.

Mendeleev’s Periodic Table

• Russian scientist Dmitri Mendeleev developed the periodic table during 1869–1872 A.D.
• He considered atomic mass as the fundamental property.
• He arranged known elements in increasing order of atomic masses.
• At that time, 63 elements were known.
• He also considered the physical and chemical properties of elements.

Groups and Periods

• Elements in the same group have similar properties.
• The arrangement helped scientists compare properties systematically.

Merits of Mendeleev’s Periodic Table

1. Correction of atomic masses

• Some atomic masses were revised to give elements their proper positions.
• Example: Beryllium was originally given atomic mass 14.09.
• Its correct value 9.4 was used and Beryllium was placed before boron.

2. Prediction of undiscovered elements

• Mendeleev left vacant places for elements not yet discovered.
• He predicted their properties.

Predicted Elements

The actual properties of these elements matched well with Mendeleev’s predictions.

3. Position of noble gases

• Noble gases were discovered later.
• Mendeleev created a separate zero group for them.
• The original periodic table was not disturbed.

Demerits of Mendeleev’s Table

1. Position of Co and Ni

Cobalt and nickel have the same whole-number atomic mass, causing ambiguity regarding their sequence.

2. Position of isotopes

Isotopes have the same chemical properties but different atomic masses, so their placement was a problem.

3. Non-uniform increase in atomic mass

Atomic masses do not increase uniformly, making it difficult to predict how many elements could be discovered between two heavy elements.

Position of Hydrogen

Hydrogen shows similarities with two different groups.

Modern Periodic Law

• After the discovery of the electron, scientists studied the relationship between atomic number and electronic configuration.
• In 1913, English scientist Henry Moseley performed experiments using an X-ray tube.
• He showed that atomic number corresponds to the positive charge on the nucleus.
• Atomic number is therefore a more fundamental property than atomic mass.

Modern Periodic Table

• It is also called the long form of the periodic table.
• Elements are arranged according to atomic number.
• The modern table explains periodic properties more accurately.
• There are 118 elements.

Periods and Groups

Lanthanide and actinide series are shown separately at the bottom.

Blocks of the Modern Periodic Table

• s-block: Groups 1 and 2
• p-block: Groups 13 to 18
• d-block: Groups 3 to 12; d-block elements are called transition elements.
• f-block: Lanthanide series and actinide series

Groups and Electronic Configuration

Electronic configuration determines the position of an element.

• Elements in the same group have similar outermost electronic configurations.
• They generally have the same number of valence electrons.
• Be, Mg and Ca belong to Group 2 and have two electrons in the outermost shell.
• F and Cl belong to Group 17 and have seven electrons in the outermost shell.

Periods and Electron Shells

Second period

Li, Be, B, C, N, O, F, Ne → electrons in K and L shells.

Third period

Na, Mg, Al, Si, P, S, Cl, Ar → electrons in K, L and M shells.

Electron Capacity of Shells

First period has 2 elements. Second period has 8 elements. Third period has 8 elements because of the law of electron octet.

Valency

• Chemical reactivity depends on the number of valence electrons.
• Across a period, valency generally increases and then decreases.
• Down a group, valency remains the same.

Periodic Trends

The chapter discusses:

1. Valency
2. Atomic size
3. Metallic–nonmetallic character

Atomic Size

• Atomic radius is used to represent atomic size.
• It is expressed in picometer (pm).
• 1 pm = 10⁻¹² m.

Across a period

Atomic radius decreases from left to right because nuclear charge increases while electrons enter the same outermost shell.

Atomic Size Down a Group

Down a group

• A new electron shell is added.
• The distance between the nucleus and outermost electron increases.
• Therefore, atomic size increases from top to bottom.

Metallic Character

• Metals generally have 1 to 3 valence electrons.
• They tend to lose electrons.
• They form positive ions called cations.
• Their electronic configuration becomes similar to a noble gas.

Nonmetallic Character

• Nonmetals tend to accept electrons.
• Their tendency to attract electrons is related to electronegativity.

Electropositivity and Electronegativity

Gradation in Halogen Family

• Group 17 contains the halogen family.
• Halogens have the general molecular formula X₂.
• Fluorine → gas
• Chlorine → gas
• Bromine → liquid
• Iodine → solid

Gradation in Alkaline Earth Metals

Group 2 contains the alkaline earth metals:

• Be does not react with water.
• Mg reacts with steam.
• Ca, Sr and Ba react with water at room temperature.
• The rate of reaction increases down the group.

Periodic Trends – One Look