Classification of Elements – Early Attempts

Introduction

In Class IX, we have learned that matter is composed of elements, compounds, and mixtures. Elements consist of only one type of atom. As of today, 118 elements are known, out of which 94 occur naturally, while the rest are synthetic.
With the discovery of more elements, scientists needed a way to systematically study and understand their properties. This led to various classification attempts, aiming to find a pattern in the properties of elements.

5.1 Making Order Out of Chaos – Early Classification Attempts

Need for Classification

As more elements were discovered, organizing them systematically became a challenge.
Scientists looked for patterns in properties to categorize elements efficiently.
The first classification divided elements into metals and non-metals based on physical and chemical properties.
Further refinements led to better classification methods.

Early Classification Attempts

1. Lavoisier’s Classification (1789)

Antoine Lavoisier, the "Father of Modern Chemistry," classified elements into metals and non-metals based on their properties.
Metals: Lustrous, good conductors of heat and electricity, malleable, and ductile.
Non-metals: Dull, poor conductors, brittle.
Limitations:

Some elements like metalloids showed properties of both metals and non-metals, making classification incomplete.
No classification for compounds.

5.1.1 Döbereiner’s Triads

Introduction

In 1817, Johann Wolfgang Döbereiner, a German chemist, attempted to classify elements based on similar properties.
He grouped elements into sets of three, which he called "Triads".
When arranged in increasing atomic mass, the middle element’s atomic mass was approximately the average of the other two elements.

Example of Döbereiner’s Triad

Consider the triad of Lithium (Li), Sodium (Na), and Potassium (K):

Element	Atomic Mass
Lithium (Li)	6.9
Sodium (Na)	23.0
Potassium (K)	39.0

The average atomic mass of Li and K: $\frac{6.9 + 39.0}{2} = 22.95 \approx 23.0$
This is approximately equal to the atomic mass of Sodium (Na).

Triads in Other Groups

Döbereiner arranged elements into different groups (A, B, and C) based on their atomic masses (Table 5.1).

Table 5.1: Groups of Three Elements

Group A	Atomic Mass	Group B	Atomic Mass	Group C	Atomic Mass
Nitrogen (N)	14.0	Calcium (Ca)	40.1	Chlorine (Cl)	35.5
Phosphorus (P)	31.0	Strontium (Sr)	87.6	Bromine (Br)	79.9
Arsenic (As)	74.9	Barium (Ba)	137.3	Iodine (I)	126.9

Groups B and C form Döbereiner Triads because the atomic mass of the middle element is roughly the average of the first and third elements.

Döbereiner’s Identified Triads (Table 5.2)

Triad	Elements	Atomic Masses
Alkali Metals	Lithium (Li), Sodium (Na), Potassium (K)	6.9, 23.0, 39.0
Alkaline Earth Metals	Calcium (Ca), Strontium (Sr), Barium (Ba)	40.1, 87.6, 137.3
Halogens	Chlorine (Cl), Bromine (Br), Iodine (I)	35.5, 79.9, 126.9

Limitations of Döbereiner’s Triads

Only a few elements could be arranged into triads.
Many known elements did not fit into this pattern.
As more elements were discovered, the classification became inadequate.

Johann Wolfgang Döbereiner – His Contribution

Born: 1780, Germany
Studied pharmacy at Münchberg, later pursued chemistry at Strasbourg.
Became a professor of chemistry and pharmacy at the University of Jena.
Major contributions:
- Discovered platinum as a catalyst.
- Proposed triads, which later contributed to the development of the Periodic Table.
Died: 1849

Conclusion

Döbereiner’s Triads were one of the first systematic attempts to classify elements.
Though limited, this concept influenced later classification models, leading to the development of the Modern Periodic Table.

5.1.2 Newlands’ Law of Octaves

Introduction

Inspired by Döbereiner’s Triads, other scientists attempted to find patterns in element properties.
In 1866, John Newlands, an English scientist, arranged elements in order of increasing atomic masses.
He started with Hydrogen (lowest atomic mass) and ended at Thorium (56th element).
He observed that every eighth element had properties similar to the first element, just like musical notes repeat every eighth note.
Based on this, he proposed the "Law of Octaves", which is known as Newlands’ Law of Octaves.

Newlands’ Octaves – Pattern in Elements

Newlands compared element periodicity to musical octaves.
In Indian music, the seven notes are: Sa, Re, Ga, Ma, Pa, Dha, Ni.
In Western music, the seven notes are: Do, Re, Mi, Fa, So, La, Ti.
Similarly, Newlands observed that every eighth element had similar properties to the first.

Table 5.3: Newlands’ Octaves

Musical Notes	Elements
Sa (Do)	H
Re (Re)	Li
Ga (Mi)	Be
Ma (Fa)	B
Pa (So)	C
Dha (La)	N
Ni (Ti)	O
Sa (Do, repeated)	F
Re (Re, repeated)	Na
Ga (Mi, repeated)	Mg
Ma (Fa, repeated)	Al
Pa (So, repeated)	Si
Dha (La, repeated)	P
Ni (Ti, repeated)	S
Sa (Do, repeated)	Cl
Re (Re, repeated)	K

The properties of Lithium (Li) and Sodium (Na) were found to be similar.
The properties of Beryllium (Be) and Magnesium (Mg) were also similar.
This pattern was found in lighter elements, but not in heavier elements.

Limitations of Newlands’ Law of Octaves

Applicable only up to Calcium (Ca):
- The pattern worked only for lighter elements (H to Ca).
- After Calcium (Ca), elements did not follow the octaves pattern.
Assumed only 56 elements existed:
- Newlands believed that no new elements would be discovered.
- However, many new elements were discovered later that did not fit his table.
Forced placement of elements:
- Newlands placed two elements in the same slot to fit his pattern.
- Example: Cobalt (Co) and Nickel (Ni) were placed in the same slot.
Incorrect grouping of dissimilar elements:
- He placed fluorine (F), chlorine (Cl), and bromine (Br) in the same column as cobalt (Co) and nickel (Ni), despite their very different properties.
- Iron (Fe), which is similar to Co and Ni, was placed far away.
Noble gases were undiscovered:
- The discovery of noble gases (He, Ne, Ar, etc.) made the Law of Octaves irrelevant, as they did not fit into the pattern.

Conclusion

Newlands’ Law of Octaves was one of the first systematic attempts at element classification.
While it worked for lighter elements, it failed for heavier elements.
The discovery of noble gases and additional elements proved the law incorrect.
However, this work influenced the development of the Modern Periodic Table.

Q U E S T I O N S

1. Did Döbereiner’s triads also exist in the columns of Newlands’ Octaves? Compare and find out.

Yes, some of Döbereiner’s triads were also present in Newlands’ Octaves.
Example: The alkali metal triad (Li, Na, K) appears in Newlands’ table.
However, not all triads fit into the Octaves pattern as Newlands’ arrangement forced some dissimilar elements together.

2. What were the limitations of Döbereiner’s classification?

Döbereiner’s classification could identify only a few triads from known elements.
Many elements did not fit into the triad pattern.
It did not work for heavier elements or newly discovered elements.
It did not explain gradual changes in properties between elements.

3. What were the limitations of Newlands’ Law of Octaves?

Applicable only up to Calcium (Ca):
- After Calcium (Ca), the periodicity of properties did not follow the Octaves pattern.
Assumed only 56 elements existed:
- Newlands believed no new elements would be discovered.
- Later-discovered elements did not fit into his classification.
Forced placement of elements:
- Some elements were forced into the same slot (e.g., Cobalt (Co) and Nickel (Ni)).
Incorrect grouping of dissimilar elements:
- He placed Fluorine (F), Chlorine (Cl), and Bromine (Br) in the same column as Cobalt (Co) and Nickel (Ni), despite their very different properties.
- Iron (Fe), which is similar to Co and Ni, was placed far away.
Noble gases were undiscovered:
- The discovery of noble gases (He, Ne, Ar, etc.) disrupted the Octaves pattern, making it irrelevant.

5.2 Making Order Out of Chaos – Mendeleev’s Periodic Table

Introduction

After the rejection of Newlands’ Law of Octaves, scientists continued searching for a classification pattern.
The most significant contribution came from Dmitri Ivanovich Mendeleev, a Russian chemist.
Mendeleev classified elements based on atomic mass and chemical properties.
His work led to the development of the Periodic Table, which became a cornerstone of modern chemistry.

Dmitri Ivanovich Mendeleev (1834-1907)

Born on 8 February 1834 in Tobolsk, Western Siberia, Russia.
Pursued higher education due to his mother’s efforts and sacrifices.
He believed science could remove superstitions, errors, and untruths.
His Periodic Table unified chemistry and motivated the discovery of new elements.

Mendeleev’s Classification of Elements

When Mendeleev started his work, 63 elements were known.
He examined the relationship between atomic mass and element properties.
Focused on compounds formed by elements with hydrogen and oxygen because:
- Hydrogen (H) and Oxygen (O) are highly reactive.
- Most elements form compounds with them (hydrides and oxides).

Mendeleev’s Method

He wrote each element’s properties on a card.
Sorted elements with similar properties together.
Arranged elements increasing atomic mass order.
Found a periodic recurrence of similar elements.
Based on this, he proposed:

Mendeleev’s Periodic Law

"The properties of elements are the periodic function of their atomic masses."

Structure of Mendeleev’s Periodic Table

Vertical columns → "Groups"
Horizontal rows → "Periods"
Formulae for Oxides & Hydrides:
- Mendeleev used ‘R’ to represent any element in a group.
- Example:
  - Carbon Hydride (CH₄) → RH₄
  - Carbon Oxide (CO₂) → RO₂
Published in 1872 in a German journal.

Significance of Mendeleev’s Work

His classification laid the foundation for the modern Periodic Table.
The table predicted the existence of undiscovered elements.
Provided a systematic way to study elements and their properties.

Mendeleev's Periodic Table

Corrections in Atomic Mass Order

While arranging elements, Mendeleev placed some elements with slightly higher atomic mass before elements with lower atomic mass to maintain similar properties within groups.
Example:
- Cobalt (Co) – atomic mass 58.9 was placed before Nickel (Ni) – atomic mass 58.7 to group it with similar elements.
- Another example of such an anomaly can be found in Table 5.4 of his periodic table.

Prediction of Undiscovered Elements

Mendeleev left gaps in his Periodic Table and boldly predicted the existence of unknown elements.
Instead of treating these gaps as errors, he considered them as elements yet to be discovered.
He named these elements by adding the Sanskrit prefix "Eka-" (meaning "one") to the name of the element just above it in the same group.
Later, these predicted elements were discovered and matched Mendeleev’s predictions closely.

Examples of Predicted Elements

Predicted Element (Eka-element)	Discovered Element
Eka-Boron	Scandium (Sc)
Eka-Aluminium	Gallium (Ga)
Eka-Silicon	Germanium (Ge)

Comparison of Eka-Aluminium and Gallium (Table 5.5)

Property	Eka-Aluminium (Predicted)	Gallium (Discovered)
Atomic Mass	68	69.7
Formula of Oxide	E₂O₃	Ga₂O₃
Formula of Chloride	ECl₃	GaCl₃

The close resemblance between Mendeleev’s predicted properties and the actual properties of gallium convinced scientists of the accuracy and usefulness of his Periodic Table.

Recognition of Mendeleev’s Work

The success of Mendeleev’s predictions proved the correctness and importance of his Periodic Table.
Chemists widely accepted his classification, making him the founder of the concept of periodicity.
The discovery of noble gases (He, Ne, Ar, etc.) did not disturb the table.
- Instead, they were placed in a new group without affecting the existing arrangement.

5.2.2 Limitations of Mendeleev’s Classification

1. Position of Hydrogen

Hydrogen resembles alkali metals (Group I) in its electronic configuration and compound formation:
- HCl (hydrogen chloride) vs. NaCl (sodium chloride)
- H₂O (water) vs. Na₂O (sodium oxide)
- H₂S (hydrogen sulfide) vs. Na₂S (sodium sulfide)
However, hydrogen also resembles halogens (Group VII) as:
- It exists as diatomic molecules (H₂ like Cl₂, Br₂).
- Forms covalent bonds with metals and non-metals.
Mendeleev could not assign hydrogen a fixed position in the table.

2. Challenge of Isotopes

Isotopes are atoms of the same element with different atomic masses but identical chemical properties.
Example: Chlorine isotopes
- Cl-35 and Cl-37 have different atomic masses but similar chemical properties.
- If Mendeleev’s classification was strictly based on atomic mass, they should be placed in different positions.
- But since they had identical properties, they needed to be placed in the same slot.
This contradicted Mendeleev’s Periodic Law, which was based solely on atomic mass.

3. Irregularity in Atomic Masses

Atomic masses do not increase uniformly between elements.
This made it difficult to predict how many elements existed between two known elements.
The issue was more significant for heavier elements, where the pattern of increasing atomic masses became unpredictable.

Conclusion

Despite its limitations, Mendeleev’s Periodic Table was a revolutionary achievement in chemistry.
His work paved the way for the modern Periodic Table, which is now based on atomic number instead of atomic mass (as proposed by Henry Moseley).

Q U E S T I O N S

1. Use Mendeleev’s Periodic Table to predict the formulae for the oxides of the following elements: K, C, Al, Si, Ba.

Potassium (K) – Group I → Oxide formula: K₂O
Carbon (C) – Group IV → Oxide formula: CO₂
Aluminium (Al) – Group III → Oxide formula: Al₂O₃
Silicon (Si) – Group IV → Oxide formula: SiO₂
Barium (Ba) – Group II → Oxide formula: BaO

2. Besides gallium, which other elements have since been discovered that were left by Mendeleev in his Periodic Table? (Any two)

Scandium (Sc) – Eka-Boron
Germanium (Ge) – Eka-Silicon

These elements were predicted by Mendeleev before their discovery, and their properties closely matched his predictions.

3. What were the criteria used by Mendeleev in creating his Periodic Table?

Mendeleev arranged elements based on:

Increasing atomic mass – Elements were placed in order of their atomic masses.
Grouping based on chemical properties – Elements in the same column (group) had similar properties.
Oxide and hydride formation – Elements were classified based on their ability to form compounds like oxides (RO, RO₂, etc.) and hydrides (RH, RH₂, etc.).
Prediction of undiscovered elements – He left gaps for elements yet to be discovered and accurately predicted their properties.

4. Why do you think the noble gases are placed in a separate group?

Noble gases (He, Ne, Ar, Kr, Xe, Rn) were discovered after Mendeleev’s Table was created.
They are chemically inert and do not readily form compounds.
Their electron configurations are stable, making them different from all other elements.
Since they did not fit into any existing group, they were placed in a new group (Group 0 or Group VIII) without disturbing the periodic table’s arrangement.

5.3 Making Order Out of Chaos – The Modern Periodic Table

Introduction

In 1913, Henry Moseley showed that the atomic number (Z) is a more fundamental property than atomic mass.
This led to the modification of Mendeleev’s Periodic Law, and the Modern Periodic Law was formulated as:

“Properties of elements are a periodic function of their atomic number.”
The atomic number represents the number of protons in an atom.
When elements are arranged in increasing atomic number, a more accurate classification is obtained.
This arrangement is called the Modern Periodic Table.

Significance of the Modern Periodic Table

The position of elements is now based on atomic number rather than atomic mass.
Prediction of element properties became more precise.
Isotopes (same atomic number but different atomic masses) were placed in the same position instead of separate slots.
The anomalous position of cobalt (Co) and nickel (Ni) was resolved.
The number of elements between two elements could now be accurately predicted.

5.3.1 Position of Elements in the Modern Periodic Table

Structure of the Modern Periodic Table

The Modern Periodic Table consists of:
- 18 vertical columns → called Groups.
- 7 horizontal rows → called Periods.
What decides the position of an element?
- Group: Determined by the number of valence electrons in the outermost shell.
- Period: Determined by the number of electron shells.

Understanding Groups and Periods

Elements in a group have the same number of valence electrons.
- Example:
  - Fluorine (F) and Chlorine (Cl) belong to Group 17.
  - Both have 7 valence electrons.
Elements in a period have the same number of shells.
- Example:
  - Li, Be, B, C, N, O, F, Ne belong to the second period.
  - All have two shells.
As atomic number increases across a period:
- The number of valence electrons increases by one.
- The number of shells remains the same.
As we go down a group:
- The number of shells increases.
- The number of valence electrons remains the same.

How Many Elements Are There in Each Period?

The number of elements in each period follows the rule 2n², where n is the shell number.

Shell (n)	Formula (2n²)	Number of Elements
K (n=1)	2 × (1)² = 2	2 (First Period)
L (n=2)	2 × (2)² = 8	8 (Second Period)
M (n=3)	2 × (3)² = 18	8 (Third Period)
N (n=4)	2 × (4)² = 32	18 (Fourth Period)

The third, fourth, fifth, sixth, and seventh periods contain 8, 18, 18, 32, and 32 elements, respectively.

5.3.2 Trends in the Modern Periodic Table

1. Valency

Valency is determined by the number of valence electrons in the outermost shell.
In a period:
- Increases from 1 to 4 and then decreases from 4 to 0.
In a group:
- Remains constant.

Examples of Valency Calculation

Magnesium (Mg, Z = 12) → Electronic configuration (2, 8, 2) → Valency 2.
Sulfur (S, Z = 16) → Electronic configuration (2, 8, 6) → Valency 2.

2. Atomic Size (Atomic Radius)

Atomic size is the distance between the nucleus and the outermost shell.

Trends in Atomic Size

Across a period (left to right): Decreases
- Nuclear charge increases, pulling electrons closer to the nucleus.
Down a group (top to bottom): Increases
- Additional electron shells increase the atomic size.

Example of Atomic Radius in Period 2

Element	B	Be	O	N	Li	C
Atomic Radius (pm)	88	111	66	74	152	77

Largest atom: Lithium (Li)
Smallest atom: Oxygen (O)

3. Metallic and Non-Metallic Properties

Metals → Tend to lose electrons → Found on the left side of the Periodic Table.
Non-metals → Tend to gain electrons → Found on the right side of the Periodic Table.
Metalloids (semi-metals) → Exhibit properties of both metals and non-metals.

Position of Metals and Non-Metals

Metals → Left side (e.g., Na, Mg).
Non-metals → Right side (e.g., Cl, S).
Metalloids → In between, along the zig-zag line.
- Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po).

4. Electropositivity and Electronegativitiy

Electropositivity (Tendency to Lose Electrons)

Across a period (left to right): Decreases
- Higher nuclear charge makes it harder to lose electrons.
Down a group (top to bottom): Increases
- Outer electrons are farther from the nucleus and easily lost.

Electronegativity (Tendency to Gain Electrons)

Across a period (left to right): Increases
- Nuclear charge increases, attracting electrons more strongly.
Down a group (top to bottom): Decreases
- Larger atomic size reduces the attraction for additional electrons.

5. Nature of Oxides

Metals form basic oxides (e.g., Na₂O, CaO).
Non-metals form acidic oxides (e.g., CO₂, SO₂).
Metalloids form amphoteric oxides (e.g., Al₂O₃, ZnO).

Conclusion

The Modern Periodic Table corrected many limitations of Mendeleev’s Periodic Table.
Elements are arranged based on atomic number, solving issues like the placement of cobalt & nickel and isotopes.
The periodic table helps predict element properties based on trends in valency, atomic size, metallic character, and electronegativity.
Metals, non-metals, and metalloids are well classified, providing a clearer understanding of chemical behavior.

Q U E S T I O N S

1. How could the Modern Periodic Table remove various anomalies of Mendeleev’s Periodic Table?

The Modern Periodic Table corrected Mendeleev’s anomalies in the following ways:

Arranged elements by atomic number (Z) instead of atomic mass → Resolved the incorrect placement of cobalt (Co) and nickel (Ni).
Isotopes were placed in the same position → Since isotopes have the same atomic number, they were not given different slots.
New group for noble gases → When noble gases were discovered, they were placed in Group 18 without disturbing the existing table.
Hydrogen’s position remained flexible → It can be placed in Group 1 or 17, recognizing its dual properties.

2. Name two elements you would expect to show chemical reactions similar to magnesium. What is the basis for your choice?

Calcium (Ca) and Barium (Ba) would show similar chemical reactions to magnesium (Mg).
Basis:
- All belong to Group 2 (Alkaline Earth Metals).
- Have 2 valence electrons, leading to similar reactivity and bonding behavior.

3. Name

(a) Three elements that have a single electron in their outermost shells:

Hydrogen (H), Lithium (Li), Sodium (Na) → All belong to Group 1 and have one valence electron.

(b) Two elements that have two electrons in their outermost shells:

Beryllium (Be), Magnesium (Mg) → Both belong to Group 2 and have two valence electrons.

(c) Three elements with filled outermost shells:

Helium (He), Neon (Ne), Argon (Ar) → All belong to Group 18 (Noble Gases) and have full valence shells, making them unreactive.

4. (a) Lithium, sodium, potassium are all metals that react with water to liberate hydrogen gas. Is there any similarity in the atoms of these elements?

Yes, all three elements (Li, Na, K) belong to Group 1 (Alkali Metals).
They each have one valence electron, which makes them highly reactive with water to produce hydrogen gas (H₂).

(b) Helium is an unreactive gas, and neon is a gas of extremely low reactivity. What, if anything, do their atoms have in common?

Both belong to Group 18 (Noble Gases) and have completely filled outermost electron shells.
Helium (He) has 2 electrons (1s²), and Neon (Ne) has 8 electrons (2s² 2p⁶).
Full valence shells make them chemically inert (unreactive).

5. In the Modern Periodic Table, which are the metals among the first ten elements?

The first ten elements are: H, He, Li, Be, B, C, N, O, F, Ne.

Metals among them: Lithium (Li) and Beryllium (Be).
Li (Group 1, Alkali Metal) and Be (Group 2, Alkaline Earth Metal) are the only metals.

6. By considering their position in the Periodic Table, which one of the following elements would you expect to have maximum metallic characteristic?

Elements given: Ga (Gallium), Ge (Germanium), As (Arsenic), Se (Selenium), Be (Beryllium).

Metallic character increases down a group and decreases across a period.
Comparing positions in the periodic table:
- Ga (Group 13) is a metal.
- Ge (Group 14) is a metalloid.
- As (Group 15) is a metalloid.
- Se (Group 16) is a non-metal.
- Be (Group 2) is a metal but lighter than Ga.

✅ Answer: Gallium (Ga) has the highest metallic character.

E X E R C I S E S

1. Which of the following statements is not a correct statement about the trends when going from left to right across the periods of the Periodic Table?

✅ Correct answer: (c) The atoms lose their electrons more easily.

Explanation:
- As we move left to right across a period:
  - Metallic character decreases (option a is correct).
  - Number of valence electrons increases (option b is correct).
  - Oxides become more acidic (option d is correct).
- Atoms lose electrons less easily due to increasing nuclear charge, making option (c) incorrect.

2. Element X forms a chloride with the formula XCl₂, which is a solid with a high melting point. X would most likely be in the same group of the Periodic Table as:

✅ Correct answer: (b) Mg (Magnesium)

Explanation:
- The formula XCl₂ suggests that X has a valency of 2.
- Elements in Group 2 (Alkaline Earth Metals) form similar compounds (e.g., MgCl₂).
- Magnesium (Mg) belongs to Group 2, making it the most likely answer.

3. Which element has:

(a) Two shells, both of which are completely filled with electrons?

✅ Answer: Neon (Ne) (Electronic configuration: 2, 8)

(b) The electronic configuration 2, 8, 2?

✅ Answer: Magnesium (Mg, Atomic number 12)

(c) A total of three shells, with four electrons in its valence shell?

✅ Answer: Silicon (Si, Atomic number 14) (Electronic configuration: 2, 8, 4)

(d) A total of two shells, with three electrons in its valence shell?

✅ Answer: Boron (B, Atomic number 5) (Electronic configuration: 2, 3)

(e) Twice as many electrons in its second shell as in its first shell?

✅ Answer: Carbon (C, Atomic number 6) (Electronic configuration: 2, 4)

4. (a) What property do all elements in the same column of the Periodic Table as boron have in common?

✅ Answer: They all have 3 valence electrons and show similar chemical reactivity.

Elements in Group 13 (Boron family) → B, Al, Ga, In, Tl.

(b) What property do all elements in the same column of the Periodic Table as fluorine have in common?

✅ Answer: They all have 7 valence electrons and are highly reactive non-metals (halogens).

Elements in Group 17 (Halogens) → F, Cl, Br, I, At.

5. An atom has electronic configuration 2, 8, 7.

(a) What is the atomic number of this element?

✅ Answer: Atomic number = 17 (Chlorine, Cl)

(b) To which of the following elements would it be chemically similar? (Atomic numbers given: N(7), F(9), P(15), Ar(18))

✅ Answer: Fluorine (F, Atomic number 9)

Explanation:
- Both Chlorine (Cl) and Fluorine (F) belong to Group 17 (Halogens).
- They have 7 valence electrons and show similar chemical properties.

6. The position of three elements A, B, and C in the Periodic Table are shown below –

Group 16	Group 17
-	-
-	A
-	-
B	C

(a) State whether A is a metal or non-metal.

✅ Answer: A is a non-metal (Group 17, Halogen).

(b) State whether C is more reactive or less reactive than A.

✅ Answer: C is less reactive than A because reactivity in Group 17 decreases down the group.

(c) Will C be larger or smaller in size than B?

✅ Answer: C will be smaller than B, as atomic size decreases across a period.

(d) Which type of ion, cation or anion, will be formed by element A?

✅ Answer: A will form an anion (A⁻) because halogens gain electrons to form negative ions.

7. Nitrogen (atomic number 7) and phosphorus (atomic number 15) belong to Group 15.

(a) Write the electronic configuration of these two elements.

✅ Nitrogen (N, Z = 7): 2, 5
✅ Phosphorus (P, Z = 15): 2, 8, 5

(b) Which of these will be more electronegative? Why?

✅ Answer: Nitrogen (N) is more electronegative than Phosphorus (P).

Electronegativity decreases down a group because atomic size increases, reducing nuclear attraction.

8. How does the electronic configuration of an atom relate to its position in the Modern Periodic Table?

✅ Answer:

Group number → Determined by valence electrons (e.g., Group 1 → 1 valence electron).
Period number → Determined by number of electron shells (e.g., 3 shells → Period 3).

9. In the Modern Periodic Table, calcium (atomic number 20) is surrounded by elements with atomic numbers 12, 19, 21, and 38.

✅ Answer: Elements with atomic numbers 12 (Magnesium, Mg) and 38 (Strontium, Sr) will have similar properties to calcium (Ca).

All belong to Group 2 (Alkaline Earth Metals) and have 2 valence electrons.

10. Compare and contrast the arrangement of elements in Mendeleev’s Periodic Table and the Modern Periodic Table.

Aspect	Mendeleev’s Periodic Table	Modern Periodic Table
Basis of Classification	Atomic mass	Atomic number
Groups & Periods	8 groups, variable periods	18 groups, 7 periods
Isotopes	Could not be placed correctly	Placed in the same position
Noble Gases	Not included (discovered later)	Included in Group 18
Position of Hydrogen	Uncertain placement	Still uncertain (can be in Group 1 or 17)
Metals & Non-Metals	No clear distinction	Clear separation with a zig-zag line

✅ Conclusion:

The Modern Periodic Table is more accurate and resolves Mendeleev’s limitations.
It correctly arranges elements by atomic number, making periodic trends more predictable.

chapter 5: Periodic Classification of Elements

Classification of Elements – Early Attempts

Introduction

5.1 Making Order Out of Chaos – Early Classification Attempts

Need for Classification

Early Classification Attempts

1. Lavoisier’s Classification (1789)

5.1.1 Döbereiner’s Triads

Introduction

Example of Döbereiner’s Triad

Triads in Other Groups

Table 5.1: Groups of Three Elements

Döbereiner’s Identified Triads (Table 5.2)

Limitations of Döbereiner’s Triads

Johann Wolfgang Döbereiner – His Contribution

Conclusion

5.1.2 Newlands’ Law of Octaves

Introduction

Newlands’ Octaves – Pattern in Elements

Table 5.3: Newlands’ Octaves

Limitations of Newlands’ Law of Octaves

Conclusion

Q U E S T I O N S

1. Did Döbereiner’s triads also exist in the columns of Newlands’ Octaves? Compare and find out.

2. What were the limitations of Döbereiner’s classification?

3. What were the limitations of Newlands’ Law of Octaves?

5.2 Making Order Out of Chaos – Mendeleev’s Periodic Table

Introduction

Dmitri Ivanovich Mendeleev (1834-1907)

Mendeleev’s Classification of Elements

Mendeleev’s Method

Mendeleev’s Periodic Law

Structure of Mendeleev’s Periodic Table

Significance of Mendeleev’s Work

5.2.1 Achievements of Mendeleev’s Periodic Table

Corrections in Atomic Mass Order

Prediction of Undiscovered Elements

Examples of Predicted Elements

Comparison of Eka-Aluminium and Gallium (Table 5.5)

Recognition of Mendeleev’s Work

5.2.2 Limitations of Mendeleev’s Classification

1. Position of Hydrogen

2. Challenge of Isotopes

3. Irregularity in Atomic Masses

Conclusion

Q U E S T I O N S

1. Use Mendeleev’s Periodic Table to predict the formulae for the oxides of the following elements: K, C, Al, Si, Ba.

2. Besides gallium, which other elements have since been discovered that were left by Mendeleev in his Periodic Table? (Any two)

3. What were the criteria used by Mendeleev in creating his Periodic Table?

4. Why do you think the noble gases are placed in a separate group?

5.3 Making Order Out of Chaos – The Modern Periodic Table

Introduction

Significance of the Modern Periodic Table

5.3.1 Position of Elements in the Modern Periodic Table

Structure of the Modern Periodic Table

Understanding Groups and Periods

How Many Elements Are There in Each Period?

5.3.2 Trends in the Modern Periodic Table

1. Valency

Examples of Valency Calculation

2. Atomic Size (Atomic Radius)

Trends in Atomic Size

Example of Atomic Radius in Period 2

3. Metallic and Non-Metallic Properties

Position of Metals and Non-Metals

4. Electropositivity and Electronegativitiy

Electropositivity (Tendency to Lose Electrons)

Electronegativity (Tendency to Gain Electrons)

5. Nature of Oxides

Conclusion

Q U E S T I O N S

1. How could the Modern Periodic Table remove various anomalies of Mendeleev’s Periodic Table?

2. Name two elements you would expect to show chemical reactions similar to magnesium. What is the basis for your choice?

3. Name

(a) Three elements that have a single electron in their outermost shells:

(b) Two elements that have two electrons in their outermost shells:

(c) Three elements with filled outermost shells:

4. (a) Lithium, sodium, potassium are all metals that react with water to liberate hydrogen gas. Is there any similarity in the atoms of these elements?

(b) Helium is an unreactive gas, and neon is a gas of extremely low reactivity. What, if anything, do their atoms have in common?

5. In the Modern Periodic Table, which are the metals among the first ten elements?