The element with Z = 20 is

1. Atomic Number and Atomic Structure (basic)

At the heart of every element lies its Atomic Number (Z), which serves as its unique chemical fingerprint. This number represents the total number of protons found in the nucleus of an atom. In a neutral atom, the number of protons is exactly equal to the number of electrons orbiting the nucleus. This balance of positive and negative charges is what makes the atom stable. For instance, if we look at Carbon, its atomic number is 6, meaning it has 6 protons and 6 electrons Science, Carbon and its Compounds, p.59.

Electrons are not scattered randomly; they are arranged in specific energy levels called shells, labeled K, L, M, and N. Each shell has a maximum capacity: the K-shell can hold up to 2 electrons, while the L-shell can hold up to 8. This arrangement is known as the electronic configuration. Let's look at how this works for two common elements:

• Carbon (Z=6): The first 2 electrons fill the K-shell, leaving 4 for the L-shell. Configuration: 2, 4 Science, Carbon and its Compounds, p.59.
• Nitrogen (Z=7): The first 2 fill the K-shell, leaving 5 for the L-shell. Configuration: 2, 5 Science, Carbon and its Compounds, p.60.

The electrons in the outermost shell are called valence electrons. These are the most important part of the atom's structure because they determine the combining capacity (valency) of the element—how it reacts and bonds with others to reach a stable state, often aiming for an "octet" of 8 electrons Science, Carbon and its Compounds, p.59-60. For example, a larger atom like Calcium (Z=20) distributes its 20 electrons as 2, 8, 8, 2, leaving it with 2 valence electrons in its fourth (N) shell.

Element	Atomic Number (Z)	K-shell	L-shell	M-shell	N-shell	Valence Electrons
Carbon	6	2	4	-	-	4
Nitrogen	7	2	5	-	-	5
Calcium	20	2	8	8	2	2

Sources: Science, Carbon and its Compounds, p.59; Science, Carbon and its Compounds, p.60

2. Electronic Configuration and Shells (basic)

To understand how the periodic table works, we must first look inside the atom. Imagine an atom as a tiny solar system where the nucleus (containing protons and neutrons) is the sun, and electrons are the planets. However, these electrons don't just orbit anywhere; they reside in specific energy levels called shells. These shells are designated by letters: K, L, M, and N, starting from the one closest to the nucleus. Each shell has a maximum capacity for electrons, calculated by the formula 2n², where 'n' is the shell number. Thus, the K-shell (n=1) can hold 2, the L-shell (n=2) can hold 8, and the M-shell (n=3) can hold up to 18 electrons.

The electronic configuration of an element describes exactly how its electrons are distributed among these shells. Electrons follow a specific filling order, occupying the innermost shells first because they are at a lower energy state. For instance, consider Calcium (Atomic Number Z = 20). Its 20 electrons are arranged as 2, 8, 8, 2 in the K, L, M, and N shells, respectively. You might wonder why the M-shell only takes 8 electrons before the N-shell starts filling—this is due to the Octet Rule, which states that atoms are most stable when they have 8 electrons in their outermost shell.

The outermost shell is known as the valence shell, and the electrons inhabiting it are called valence electrons. These valence electrons are the "negotiators" of the chemical world; they determine how an atom reacts, bonds, or shares with others. For example, elements like Carbon share their valence electrons to attain a stable "noble gas configuration," which is the gold standard of atomic stability Science, Carbon and its Compounds, p.59. Understanding this configuration is the key to predicting an element's position in the periodic table and its chemical personality.

Sources: Science (NCERT 2025 ed.), Carbon and its Compounds, p.59

3. The Modern Periodic Law (basic)

In our previous discussions, we looked at how early scientists tried to find a pattern in the chaos of elements. The real breakthrough came when we moved away from 'Atomic Mass' and embraced the Atomic Number (Z). Just as nature uses periodic cycles like the phases of the moon to mark time Science Class VIII, Keeping Time with the Skies, p.178, the Modern Periodic Law reveals a repeating rhythm in the chemical world.

The law states: "The physical and chemical properties of the elements are a periodic function of their atomic numbers." This means that if you arrange elements in increasing order of their atomic number, elements with similar properties will reappear at regular, predictable intervals. This shift was revolutionary because the atomic number represents the number of protons in the nucleus (and electrons in a neutral atom), which directly dictates how an element reacts. For example, atoms of elements like Gold or Magnesium are classified as metals because of how their electrons are structured Science Class VIII, Nature of Matter, p.123.

To understand this periodicity, let's look at Calcium (Z=20). Its electrons are arranged in shells as 2, 8, 8, 2 Science Class X, Metals and Non-metals, p.47. Because it has two electrons in its outermost shell, it behaves very similarly to Magnesium (Z=12), which also has two valence electrons. The Modern Periodic Law ensures these 'chemical cousins' are placed in the same vertical column, known as a Group. Specifically, Calcium belongs to Group 2, the alkaline earth metals.

Feature	Mendeleev’s Periodic Law	Modern Periodic Law
Basis of Classification	Atomic Mass	Atomic Number (Z)
Reason for Properties	Weight of the atom	Electronic configuration
Scientific Accuracy	Had anomalies (e.g., isotopes)	Explains isotopes and positions perfectly

Sources: Science Class VIII, Keeping Time with the Skies, p.178; Science Class VIII, Nature of Matter: Elements, Compounds, and Mixtures, p.123; Science Class X, Metals and Non-metals, p.47

4. Periodic Trends in Properties (intermediate)

To understand why elements behave the way they do, we must look at Periodic Trends—the predictable patterns in chemical and physical properties as we move across the table. These trends are primarily driven by two factors: the number of shells (which increases as we go down a group) and the effective nuclear charge (the 'pull' the nucleus exerts on electrons, which increases as we move left to right). For instance, as we move across a period, the atomic radius tends to decrease because the increasing nuclear charge pulls the electron cloud closer to the center.

One of the most fundamental trends is Valency. Valency is determined by the number of electrons in the outermost shell, known as valence electrons. Elements in the same group share the same number of valence electrons, which is why they exhibit similar chemical properties. For example, Calcium (Z=20) has a configuration of 2, 8, 8, 2 Science, Class X (NCERT 2025 ed.), Chapter 3, p. 47. Because it has two valence electrons, it belongs to Group 2, known as the Alkaline Earth Metals. These metals are characterized by their tendency to lose those two electrons to form +2 ions, making them highly reactive and leading to the formation of basic (alkaline) oxides.

We also distinguish elements by their Metallic and Non-metallic character. Metals are 'electropositive,' meaning they lose electrons easily, while non-metals are 'electronegative' and tend to gain electrons. This transition is clearly visible as you move from the left side of the periodic table (metals like Calcium) to the right side (non-metals like Oxygen or Chlorine). Metals are generally lustrous, malleable, and good conductors of electricity, whereas non-metals often lack these properties and form acidic oxides Science, Class VII (NCERT 2025 ed.), The World of Metals and Non-metals, p. 54.

Property	Across a Period (Left to Right)	Down a Group (Top to Bottom)
Atomic Radius	Decreases	Increases
Valency	Increases then decreases	Remains the same
Metallic Character	Decreases	Increases
Nature of Oxides	Basic → Amphoteric → Acidic	Generally remains same

Sources: Science, Class X (NCERT 2025 ed.), Metals and Non-metals, p.47; Science, Class VII (NCERT 2025 ed.), The World of Metals and Non-metals, p.54

5. Chemical Bonding and Reactivity (intermediate)

At the heart of chemistry lies a single quest: the pursuit of stability. Most elements are naturally unstable because their outermost electron shells are incomplete. To find peace, they strive to achieve a Noble Gas Configuration—a state where the outermost shell is completely full (usually 8 electrons, known as an octet). This drive is what triggers chemical reactions and the formation of bonds Science, Metals and Non-metals, p.47.

There are two primary ways atoms achieve this stability: Ionic Bonding and Covalent Bonding. In Ionic Bonding, atoms completely transfer electrons from one to another. For instance, a Sodium (Na) atom, which has one lonely electron in its outer shell, will gladly give it away to a Chlorine (Cl) atom, which is just one electron short of a full set Science, Metals and Non-metals, p.46. This transfer creates charged particles called ions (Na⁺ and Cl⁻). Because opposite charges attract, they stick together in a powerful electrostatic grip. This explains why ionic compounds have very high melting points and can conduct electricity when dissolved in water—the ions are free to move and carry a charge Science, Carbon and its Compounds, p.58.

On the other hand, Covalent Bonding involves the sharing of electron pairs. Carbon is the master of this technique. Instead of giving electrons away, it shares them with other atoms to fill its shell Science, Carbon and its Compounds, p.60. While the bonds inside these molecules are very strong, the forces between the molecules are quite weak. Consequently, covalent compounds usually have lower melting points and do not conduct electricity because they don't form free-floating ions Science, Carbon and its Compounds, p.59.

Feature	Ionic Compounds	Covalent Compounds
Formation	Transfer of electrons (Metal to Non-metal)	Sharing of electrons (usually Non-metals)
Melting/Boiling Point	High (strong electrostatic forces)	Low (weak intermolecular forces)
Electrical Conductivity	High (in molten/solution state)	Poor (no ions formed)

An element's position in the Periodic Table tells us exactly how it will react. For example, Calcium (Z=20) has an electronic configuration of 2, 8, 8, 2. With two electrons in its outermost shell, it belongs to Group 2 (Alkaline Earth Metals). It is highly reactive because it "wants" to lose those two extra electrons to reveal a stable octet underneath, forming a Ca²⁺ ion Science, Metals and Non-metals, p.47.

Sources: Science, Metals and Non-metals, p.46; Science, Metals and Non-metals, p.47; Science, Carbon and its Compounds, p.58; Science, Carbon and its Compounds, p.59; Science, Carbon and its Compounds, p.60

6. Families of the Periodic Table (exam-level)

In the study of chemistry, we organize elements into Families (or Groups) based on their chemical behavior. This behavior is primarily dictated by their electronic configuration—specifically the number of electrons in their outermost shell, known as valence electrons. Elements in the same family share similar properties because they seek to achieve stability in the same way, often by gaining or losing the same number of electrons to reach a stable 'noble gas' configuration Science, Chapter 3, p.47.

Two of the most prominent metallic families are the Alkali Metals (Group 1) and the Alkaline Earth Metals (Group 2). While both are reactive, they have distinct identities. Alkali metals like Sodium (Na) and Potassium (K) have only one valence electron, making them so reactive and soft that they can be cut with a knife Science, Chapter 3, p.40. In contrast, Alkaline Earth Metals like Magnesium (Mg) and Calcium (Ca) have two valence electrons. For example, Calcium (Atomic Number 20) has a configuration of 2, 8, 8, 2. Because it has two electrons in its outermost 'N' shell, it belongs to Group 2 Science, Chapter 3, p.47.

Understanding these families helps us predict how an element will react. For instance, Group 18 elements are known as Noble Gases (like Helium, Neon, and Argon). They have completely filled outer shells, which makes them chemically unreactive or 'inert' Science, Chapter 3, p.47. On the other hand, metals are grouped by their reactivity levels—from highly reactive metals that must be stored carefully to low-reactivity metals like Gold or Platinum Science, Chapter 3, p.50.

Family Name	Group No.	Valence Electrons	Key Characteristics
Alkali Metals	1	1	Very reactive, soft, low melting points.
Alkaline Earth Metals	2	2	Reactive, form +2 ions, basic oxides.
Noble Gases	18	8 (except He: 2)	Chemically inert, stable octet configuration.

Sources: Science (NCERT 2025 ed.), Metals and Non-metals, p.39-40; Science (NCERT 2025 ed.), Metals and Non-metals, p.47; Science (NCERT 2025 ed.), Metals and Non-metals, p.50

7. Solving the Original PYQ (exam-level)

Now that you have mastered the fundamentals of atomic structure and periodic classification, this question serves as a perfect application of those building blocks. To solve this, you must first recall that the atomic number (Z = 20) identifies the element as Calcium (Ca). By applying the Bohr-Bury scheme of electronic configuration, you can distribute these 20 electrons across shells as 2, 8, 8, 2. The crucial observation here is the presence of two electrons in the outermost shell, which signifies its position in the periodic table. According to Science, Class X (NCERT), the number of valence electrons determines the group; hence, Calcium belongs to Group 2, which is exclusively comprised of alkaline earth metals.

UPSC often tests your ability to distinguish between adjacent groups, so elimination is your best tool. Option (A), alkali metals, refers to Group 1 elements which possess only one valence electron (like Potassium, Z=19). Options (C) and (D) represent the non-metallic side of the table; halogens occupy Group 17 with seven valence electrons, while inert gases reside in Group 18 with a stable, complete octet. By systematically matching the electron distribution to the specific chemical family, you can confidently identify that an element with two valence electrons must be an alkaline earth metal, making (B) the correct answer.