Q1. What is the atomic number (Z) and why is it important in identifying an element? Explain how it relates to protons, electrons (in a neutral atom), and the position in the periodic table.
Answer:
The atomic number (Z) is the number of protons in the nucleus of an atom. It uniquely identifies an element — no two different elements have the same Z.
In a neutral atom, the number of electrons equals the number of protons (so electrons = Z). This balance of charges makes the atom electrically neutral.
The atomic number determines an element’s chemical identity and largely its chemical behaviour because electrons (especially valence electrons) control bonding and reactions.
In the periodic table, elements are arranged in order of increasing atomic number, so Z decides the element’s position (period and group trends follow from electron configuration).
Example: For carbon, Z = 6, so it has 6 protons and, in a neutral atom, 6 electrons. Knowing Z helps predict how the atom will bond and where it appears in the periodic table.
Q2. How do you calculate the number of neutrons in an atom if you know its mass number (A) and atomic number (Z)? Give two worked examples including isotopic notation using nuclear symbols.
Answer:
The mass number (A) equals protons + neutrons. To find neutrons: neutrons = A − Z. This is because A counts nucleons (protons and neutrons).
Example 1: For carbon-14 written as 146C, A = 14 and Z = 6. Neutrons = 14 − 6 = 8 neutrons. This isotope is called carbon-14.
Example 2: For aluminium with symbol 2713Al, A = 27, Z = 13. Neutrons = 27 − 13 = 14 neutrons.
Steps to solve: identify A from the superscript, identify Z from the subscript, subtract to get neutrons.
This calculation works for any isotope and shows why isotopes of the same element have identical Z but different A.
Q3. Distinguish between mass number (A) and atomic mass (relative atomic mass). Why do textbooks sometimes use “u” (atomic mass unit) and when is A an integer?
Answer:
Mass number (A) is a whole number equal to the sum of protons and neutrons in a particular nucleus (e.g., 126C has A = 12). It is always an integer because nucleons are discrete particles.
Atomic mass (or relative atomic mass) is the average mass of atoms of an element as they occur naturally, taking into account the abundances of isotopes. It is measured in atomic mass units (u) and is often not an integer (for carbon it’s ≈ 12.011 u).
The unit 1 u is defined as 1/12 the mass of a 126C atom, so masses of protons and neutrons are close to 1 u each. Textbooks use u to show masses on an atomic scale.
In summary, A is an integer for one isotope; atomic mass in u is an average and may be decimal.
Q4. Explain how to read and interpret the nuclear symbol 147N. What information does it give about protons, neutrons, and electrons in different cases (neutral atom vs ion)?
Answer:
The nuclear symbol 147N gives two key numbers: the superscript 14 is the mass number (A) and the subscript 7 is the atomic number (Z).
From Z = 7 we know there are 7 protons in the nucleus. From A = 14, neutrons = A − Z = 14 − 7 = 7 neutrons. So this is the isotope nitrogen-14.
For a neutral atom, electrons = protons = 7, so it has 7 electrons. If the atom is an ion, the number of electrons changes: e.g., N+ (a positive ion) would have 6 electrons, while N− (rare) would have 8 electrons.
Thus the nuclear symbol compactly tells you identity (element), nucleon count (A), and proton count (Z); electron count depends on the atom’s charge.
Q5. Explain the role of neutrons in the nucleus. Why are they important even though they have no electric charge? Include a brief note on how neutron number affects stability and isotopes.
Answer:
Neutrons are neutral particles in the nucleus that, together with protons, are called nucleons. Although neutrons carry no charge, they play a crucial role in nuclear stability.
Neutrons add strong nuclear force attraction that helps bind the nucleus together and reduce the electrostatic repulsion between positively charged protons. Without enough neutrons, repulsion between protons can make the nucleus unstable.
Changing the number of neutrons produces isotopes — atoms of the same element (same Z) but different mass numbers (A). Some isotopes are stable, while others are radioactive and decay because the neutron–proton ratio is outside the stable range.
Hence, neutrons are essential for providing mass and stability to nuclei; they affect nuclear properties but not chemical behaviour (which depends on electrons).
High Complexity (Analytical & Scenario-Based)
Q6. A sample of element X contains two stable isotopes: 3517X (75% abundance) and 3717X (25% abundance). Calculate the average atomic mass of X in u. Explain each step and describe how this average relates to the mass number A.
Answer:
To find the average atomic mass, take a weighted mean of isotope masses using their natural percent abundances. Convert percentages to fractions: 75% = 0.75 and 25% = 0.25.
Multiply each isotope’s mass number (which approximates its mass in u) by its fraction and add:
Contribution of 3517X = 35 × 0.75 = 26.25 u
Contribution of 3717X = 37 × 0.25 = 9.25 u
Sum = 26.25 + 9.25 = 35.50 u. So the average atomic mass ≈ 35.50 u.
This average is not an integer because it reflects a mixture of isotopes with different A values. While each isotope has an integer mass number (A), the observed atomic mass of a natural sample is a weighted average of those integer masses.
Q7. Consider an atom with atomic number 17 and mass number 35. State the number of protons, neutrons, and electrons in: (a) a neutral atom, (b) the anion formed when it gains 1 electron, and (c) the cation formed when it loses 1 electron. Also explain how gaining or losing electrons affects chemical behaviour.
Answer:
For the atom 3517X (atomic number Z = 17, mass number A = 35): protons = 17 always, neutrons = A − Z = 35 − 17 = 18.
(b) Anion (gains 1 electron): electrons = 17 + 1 = 18 electrons. Composition: 17 protons, 18 neutrons, 18 electrons. This ion has a −1 charge.
(c) Cation (loses 1 electron): electrons = 17 − 1 = 16 electrons. Composition: 17 protons, 18 neutrons, 16 electrons. This ion has a +1 charge.
Gaining or losing electrons changes the atom’s charge and chemical reactivity: ions interact strongly via electrostatic forces, form ionic bonds, and have different properties from the neutral atom. The nucleus (protons + neutrons) remains unchanged during ion formation.
Q8. Define isotopes, isobars, and isotones and give one clear example of each. Explain why isotopes have similar chemical behaviour but different physical properties.
Answer:
Isotopes are atoms with the same atomic number (Z) but different mass numbers (A) — same element, different neutron count. Example: 126C and 146C. They have identical chemical behaviour because they have the same number of electrons and electronic structure, but different mass and some different nuclear properties (e.g., radioactivity).
Isobars are atoms with the same mass number (A) but different atomic numbers (Z) — different elements with the same total nucleons. Example: 4018Ar and 4020Ca. Chemically they differ because Z (and electrons) differ, while nuclear mass is same.
Isotones are atoms with the same number of neutrons but different Z. Example: 146C (8 neutrons) and 157N (8 neutrons).
Isotopes behave chemically similar because chemical reactions depend on electrons, not neutrons; physical properties like density, mass, and nuclear stability change with neutron number.
Q9. A student finds a radioactive sample and measures that an atom in it has 92 protons and 146 neutrons. (a) Identify the element and write its nuclear symbol. (b) Calculate its mass number. (c) Discuss briefly why heavy nuclei with high neutron and proton numbers can be unstable.
Answer:
(a) An atom with 92 protons is the element uranium (Z = 92). Its nuclear symbol is 23892U if we use the numbers below.
(b) Mass number A = protons + neutrons = 92 + 146 = 238. So the full notation is 23892U (uranium-238).
(c) Heavy nuclei like uranium have many protons, which create large electrostatic repulsion between positive charges. To bind this many protons, a higher number of neutrons is needed to provide extra strong nuclear force attraction. However, when the neutron–proton balance becomes unfavourable, internal forces compete and the nucleus may become unstable and undergo radioactive decay (alpha, beta, etc.). Large nuclei have many possible decay paths, making them more commonly radioactive than light nuclei.
Q10. Two atoms have the same mass number but different atomic numbers (they are isobars). If you were given 4018Ar and 4020Ca, compare their nuclear composition, **electron co...
