SSS 1: ATOMIC STRUCTURE

Atom

An atom is the smallest particle of an element which can take part in a chemical reaction.

Components of an atom

An atom has two main parts:

I. The nucleus (made of proton and neutron)

II. The shells (electrons) which surrounds the nucleus at definite distance.

Structure of an atom

Sub-particles of an atom

An atom is made up of three sub-particles namely:

I. Proton

II. Electron

III. Neutron

Proton and neutron are the nuclear particles found in the nucleus, while the electrons are in constant and continuous motion in the shells at various energy levels around the nucleus.

Characteristics of the sub-particles

NOTE: The nucleus is also positively charged due to the protons.

Atomic number and Mass number

Atomic number: Atomic number is the number of protons in the nucleus of an atom. It has the symbol Z

Mass number (Atomic mass or Nucleon number): Mass number is the total number of protons and neutrons in the nucleus of an atom. It has the symbol A

Examples:

1. Given:

I. Proton number = 29

II. Electron number = 29

III. Neutron number = 63 - 29 = 34

IV. Mass number = 63

V. Atomic number = 29

2. The mass number of an element is 31. If its atomic number is 15, what are the compositions of:

(a) the nucleus of the atom?

(b) its atom?

Solution:

(a) Nucleus = proton + neutron

Proton number = atomic number = 15

Neutron number = mass - atomic

number number

= 31 - 15

= 16

The nucleus is made up of 15 protons and 16 neutrons

(b) An atom is made up of proton, neutron and electron.

In a neutral atom, the number of electrons equals the number of protons

Number of electrons = 15

The compositions of the atom are: 15 protons, 15 electrons and 16 neutrons

(3) The total number of protons and electrons in a neutral atom is 26. If its mass number is 27, find the number of neutrons.

Solution:

Protons + electrons = 26

Protons (x) = electrons (x) in a neutral atom

x + x = 26

2x = 26

x = 26/2

x = 13, number of protons= 13

Mass number = proton + neutron

27 = 13 + neutron

neutron = 27 - 13

neutron = 14

Isotope

Isotopes are atoms of the same element with the same atomic number but different mass number.

Isotope could also be defined as atoms of the same element with the same number of protons but different number of electrons.

Isotopy

Isotopy is a phenomenon whereby atoms of the same element have the same atomic number but different mass number.

Examples of elements exhibiting isotopy:

Atomic structure and symbols of the three isotopes of hydrogen: Protium(p), Deuterium(D), Tritium(T)

Relative Atomic Mass of Isotopic

Elements

Atoms of the same isotope have the same mass number. However, the relative atomic mass of an element that has two or more isotopes depend on the percentage or relative abundance of each isotope in the naturally occuring element.

Mass spectrometer is used to determine the number of isotopes of an element and their relative abundance.

Simple Calculations:

1. Naturally occuring chlorine atoms contains 75% of ³⁵₁₇Cl and 25% of ³⁷₁₇Cl by mass respectively. Calculate the mean relative atomic mass of chlorine.

Solution:

³⁵₁₇Cl ³⁷₁₇Cl

(75% x 35) + (25% x 37)

26.25 + 9.25

35.5

Mean relative atomic mass of Cl = 35.5

2. An element X has two isotopes ²⁰X and ²²X, which occur in the ratio 1:3, calculate the mean relative atomic mass of X.

Solution:

Given ratio: 1 : 3

Total ratio = 1 + 3 = 4

²⁰X ²²X

(1/4 × 20) + (3/4 × 22)

5 + 16.5

21.5

3. An element X with relative atomic mass of 16.2 contains isotopes ¹⁶₈X with relative abundance of 90% and ᵐ₈X with relative abundance of 10%. What is the value of m?

Solution:

Relative atomic mass of X = 16.2

¹⁶₈X ᵐ₈X

(90% x 16) + (10% x m) = 16.2

(14.40) + (m/10) = 16.2

m/10 = 16.2 - 14.40

m/10 = 1.8

m = 18

4. An element B exists in two isotopic forms in ratio 2:3. If the isotopes have these relative atomic masses as 15 and 17 respectively. Calculate the relative atomic mass of the element B.

Solution:

Sum of ratio of isotopes = 2 + 3 = 5

¹⁵B ¹⁷B

(2/5 x 15) + (3/5 x 17)

6 + 10.2

16.2

Relative atomic mass of B = 16.2

5. If a sample of an element Q exists in a ratio 1:2:3. Calculate the relative atomic masses of each isotope of Q given that the relative atomic mass of Q is 1.1

Solution:

Sum of ratio of isotopes = 1+2+3 = 6

First isotope of Q = 1/6 x 1.1 = 0.18

Second isotope of Q = 2/6 x 1.1 = 0.39

Third isotope of Q = 3/6 x 1.1 = 0.55

Quantum numbers

Quantum numbers are numbers assigned to all the electrons in an atom. They describe certain characteristics of the electrons.

The arrangement of the electrons around the nucleus of an atom is determined by the four quantum numbers of each electron in the atom.

The four quantum numbers are:

1. Principal quantum numbers

2. Subsidiary (Azimuthal) quantum

numbers

3. Magnetic quantum numbers

4. Spin quantum numbers

Principal quantum numbers(n)

1. The principal quantum number (n) describes the energy level of an electron or shells.

2. It has integer values n = 1, 2, 3, 4, 5, and so on, corresponding to the K, L, M, N shells, respectively.

3. The maximum number of electrons a shell can hold is determined by the formula 2n², where n is the principal quantum number.

For example, the third shell (n = 3) can hold up to 2 × 3² = 18 electrons.

4. The principal quantum number gives an idea of the electron density around the nucleus

5. It shows the number of subshells

NOTE:

1. As n increases, the energy and distance of electrons from the nucleus also increase.

2. The total numbers of orbitals is given by n². where 'n' is the principal quantum number of the shell.

3. The orbit closest to the nucleus is also the orbit of minimum energy (or the ground state of the atom)

Subsidiary (Azimuthal) Quantum

Number (l)

1. The subsidiary quantum number (l) describes the shape of the orbital in which an electron is found.

2. It also determines the subshell (s, p, d, or f) to which the electron belongs.

3. It has integral values ranging from 0 to (n – 1). i.e. l = 0, 1, 2, 3, ... (n - 1). The electrons with subsidiary quantum numbers; 0,1,2, and 3 are usually refer to as the s,p,d, and f electrons respectively.

Examples:

1. If n = 1: l = 0 → 1s (only s-subshell)

2. If n = 2: l = 0, 1 → 2s and 2p subshells

3. If n = 3: l = 0, 1, 2 → 3s, 3p, and 3d subshells

4. If n = 4: l = 0, 1, 2, 3 → 4s, 4p, 4d, and 4f subshells

Magnetic quantum number(mₗ)

1. The magnetic quantum number, represented by mₗ, describes the orientation of an orbital in space relative to the three axes (x, y, and z).

2. It shows the number of orbitals that are present within a subshell.

3. It has whole number value from –1 through 0 to +1, including zero.

The number of possible mₗ values for a given subshell = the number of orbitals in that subshell. i.e.

s subshell → 1 orbital → 2 electrons

p subshell → 3 orbitals → 6 electrons

d subshell → 5 orbitals → 10 electrons

f subshell → 7 orbitals → 14 electrons

Spin quantum number

1. The spin quantum number, represented by mₛ, describes the direction of spin (rotation) of an electron about its own axis within an orbital.

2. Electrons behave as if they spin on their axes, and this spinning motion creates a tiny magnetic field.

3. Each electron in an orbital can spin in one of two opposite directions:

mₛ = +½ or –½

(I) +½ (spin-up) → represents an electron spinning in a clockwise direction (↑)

(II) –½ (spin-down) → represents an electron spinning in the anticlockwise or opposite direction (↓)

4. Each orbital can hold a maximum of two electrons. These two electrons must spin in opposite directions (one +½ and the other –½).

This opposite spin is in accordance with Pauli’s Exclusion Principle, which states that no two electrons in an atom can have the same set of four quantum numbers.

(↑↓)→ Two electrons in an orbital with opposite spins.

5. Opposite spins reduce electron repulsion within the same orbital.

Orbitals

An orbital is a region or space around the nucleus of an atom where electrons can be found with a certain amount of energy.

Each orbital can hold a maximum of two electrons. These two electrons must spin in opposite directions (↑↓).

Types of orbitals

The four types of orbitals are:

I. s-orbital

II. p-orbital

III. d-orbital

IV. f-orbital

s-orbital

Shape of s-orbital

Properties of s-orbital:

1. The s-orbital is spherical in shape.

2. It is spherically symmetrical about the nucleus.

3. Each energy level has only one s-orbital (e.g., 1s, 2s, 3s...)

4. Size increases with energy level:

The size of the s-orbital increases as the principal quantum number (n) increases.

(1s < 2s < 3s < 4s)

5. Each s-orbital can hold a maximum of 2 electrons with opposite spins (↑↓).

p-orbital

shapes of p-orbitals

Properties of p-orbital:

1. The p-orbital has a dumbbell shape (two lobes on opposite sides of the nucleus).

2. There are three p-orbitals in each energy level (except the first): Px, Py, and Pz.

3. All three p-orbitals in the same shell have equal energy

4. p-orbitals first appear in the second energy level (n = 2).

5. p-orbitals contain a maximum of six (6) electrons with each holding 2 electrons

6. The electron density is concentrated on both sides of the nucleus along a specific axis.

d-orbitals

Shapes of d-orbitals

Properties of d–Orbitals

1. Most d–orbitals have a cloverleaf (four-lobed) shape.

2. There are five d–orbitals

3. The d–orbitals contain a maximum of 10 electrons in total with each holding 2 electrons

4. d–orbitals first appear from the third energy level (n = 3).

5. d–orbitals play a key role in the formation of transition elements, showing variable oxidation states, colored ions, and magnetic properties.

f-orbitals

Shapes of f-orbitals

Properties of f–Orbitals

1. The shapes of f–orbitals are complex and multi-lobed (some have 6 or 8 lobes).

2. There are seven f–orbitals

3. The f-orbitals can hold a maximum of 14 electrons,

4. f–orbitals first appear in the fourth energy level (n = 4) i.e. in the 4f–subshell.

5. They contribute to the formation of inner transition elements: the lanthanides and actinides.

6. f–orbitals are responsible for properties such as magnetism, color, and variable oxidation states in these elements.

7. Their energy levels are very close to those of d–orbitals, leading to complex electron configurations.

Rules governing electron filling in

orbitals

The three principles or rules that govern electron filling into orbitals are:

I. Pauli Exclusion principle

II. Hund's rule

III. Aufbau principle

Pauli Exclusion principle

The Pauli Exclusion Principle was proposed by Wolfgang Pauli in 1925.

Pauli Exclusion principle states that two electrons in the same orbital of an atom cannot have the same value for all set of four quantum numbers.

Explanation: According to Pauli’s principle, two electrons can share the same orbital only if they have opposite spins (↑↓)

Hund's rule

Hund’s Rule of Maximum Multiplicity states that electrons occupy orbitals singly before they pair up in degenerate orbitals.

NOTE:

1. All the unpaired electrons in singly occupied orbitals have parallel spins.

2. Orbitals of the same energy level (like the three p-orbitals, five d-orbitals, or seven f-orbitals) are called degenerate orbitals.

Aufbau principle

Aufbau means building up.

Aufbau principle states that electrons always occupy the lower energy level before filling higher energy orbitals

Aufbau principle states that electrons occupy orbitals in order of increasing energy.

Aufbau principle can easily be derived using this:

1s 2s 3s 4s 5s...

2p 3p 4p 5p...

3d 4d 5d...

4f 5f ...

1s 2s 2p 3s 3p 4s 3d 4p 5s...

This quantum numbers can be used to determine the electronic configurations of the atoms of all the known elements.

1s< 2s< 2p< 3s< 3p< 4s< 3d...

<————————————————

Order of increasing energy of orbitals

1s has the lowest energy.

3d has higher energy than 4s.

NOTE:

1s² : means in the first energy level, the s-orbital has 2 electrons

2p³ : means in the second energy level, the p-orbital has 3 electrons

3p⁶ : means in the third energy level, the p-orbital has 6 electrons

4d² : means in the fourth energy level, the d-orbital has 2 electrons

Electronic configuration

Electronic configuration is the arrangement of electrons in the orbitals of an atom.

Example;

Write the electronic configurations of the following elements:

1. Hydrogen (₁H)

2. Sodium (₁₁Na)

3. Calcium (₂₀Ca)

Solution:

1. Hydrogen: 1s¹

2. Sodium: 1s² 2s² 2p⁶ 3s¹

3. Calcium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Orbital diagrams

Electron Shell

An electron shell (also called an energy level or orbit) is the region around the nucleus where electrons are likely to be found.

Each shell is denoted by the principal quantum number (n), such as n = 1, 2, 3, 4... or by letters K, L, M, N... respectively.

The closest shell to the nucleus is the K shell.

Duplet rules

The duplet rule states that an atom is stable when it has two electrons in its outermost shell.

This type of stability is observed in light elements such as hydrogen (H), helium (He), and lithium (Li).

Examples:

Hydrogen (H): Has 1 electron. It needs 1 more electron to complete its outermost shell (2 electrons). It forms H₂ molecule by sharing electrons (H–H).

Helium (He): Already has 2 electrons in its outermost shell. The most stable (noble gas).

Lithium (Li): Has 2,1 electron configuration. It loses 1 electron to form Li⁺, achieving a duplet.

Octet Rule

The octet rule states that an atom becomes stable when it has eight electrons in its outermost shell, similar to the noble gases (except helium).

Most main-group elements follow this rule when forming compounds.

Examples:

1. Sodium (Na): 1s² 2s² 2p⁶ 3s¹. It loses 1 electron to form Na⁺ (with 8 electrons in the next inner shell).

2. Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵. It gains 1 electron to form Cl⁻ (3s² 3p⁶ = 8 electrons).

In NaCl, both Na⁺ and Cl⁻ achieve stable octet configurations.

Other examples:

Oxygen (O) gains 2 electrons → O²⁻

Nitrogen (N) shares 3 electrons → NH₃ (ammonia)

Iso-electronic species

Iso-electronic species are atoms, ions or molecules that contain the same number of electrons but may have different nuclear charges (different number of protons) and different chemical formulas. For example, O²⁻, F⁻, Ne, Na⁺, Mg²⁺ have 10 electrons and the same electronic configuration: 1s² 2s² 2p⁶.

PART 1: Objective Questions on Atomic Structure

1. Which statement best explains why isotopes of the same element exhibit identical chemical properties?

(a) They possess the same number of neutrons.

(b) They possess identical electronic configurations.

(d) They possess equal nuclear binding energies.

---

2. The experimentally determined relative atomic mass of an element is best described as:

(a) The mass number of its most abundant isotope.

(b) The weighted mean of the masses of its naturally occurring isotopes.

(d) The mass of one mole of the element.

---

3. Which quantum number primarily determines the size and energy of an atomic orbital?

(a) Spin quantum number

(b) Magnetic quantum number

(d) Azimuthal quantum number

---

4. Which of the following statements correctly describes the Pauli Exclusion Principle?

(a) Electrons occupy orbitals of equal energy singly before pairing.

(b) No two electrons in the same atom can have the same set of four quantum numbers.

(d) Electrons move around the nucleus in fixed circular paths.

---

5. Which set of quantum numbers is NOT permissible?

(a) n = 3, l = 2, mₗ = +2

(b) n = 4, l = 3, mₗ = –1

(d) n = 5, l = 1, mₗ = 0

---

6. The maximum number of orbitals in the fourth principal energy level is:

(a) 8

(b) 12

(d) 32

---

7. Which orbital possesses the highest angular momentum?

(a) s

(b) p

(d) f

---

8. The electronic configuration of chromium is considered exceptional because of:

(a) Nuclear instability

(b) Half-filled d-orbital stability

(d) Isotopic abundance

---

9. The magnetic quantum number describes the:

(a) Spin of electrons

(b) Shape of orbitals

(d) Energy of electrons

---

10. Which species is NOT isoelectronic with Ne?

(a) Na⁺

(b) Mg²⁺

(d) Al³⁺

---

11. The maximum number of electrons that can occupy all orbitals when n = 5 is:

(a) 25

(b) 32

(d) 72

---

12. Which instrument is primarily employed for determining isotopic abundance?

(a) X-ray diffractometer

(b) Spectrophotometer

(d) Electron microscope

---

13. The Aufbau Principle is based upon:

(a) Increasing orbital size

(b) Increasing orbital energy

(d) Increasing spin

---

14. Which of the following possesses the greatest nuclear charge?

(a) O²⁻

(b) F⁻

(d) Na⁺

---

15. The total number of electrons that can occupy all p-orbitals is:

(a) 2

(b) 4

(d) 10

---

16. Which orbital first appears in the third principal energy level?

(a) s

(b) p

(d) f

---

17. Hund's Rule predicts that electrons will occupy degenerate orbitals:

(a) In pairs first

(b) Randomly

(d) According to nuclear charge

---

18. Which isotope contributes most to the relative atomic mass of chlorine?

(a) ³⁴Cl

(b) ³⁵Cl

(d) ³⁷Cl

---

19. Which property is primarily affected by isotopic variation?

(a) Chemical reactivity

(b) Atomic number

(d) Electron configuration

---

20. The total number of possible orbitals in the fifth shell is:

(a) 15

(b) 20

(d) 50

---

Answers

1. (b)

2. (b)

3. (c)

4. (b)

5. (c)

6. (c)

7. (d)

8. (b)

9. (c)

10. (d)

11. (c)

12. (c)

13. (b)

14. (d)

15. (c)

16. (c)

17. (c)

18. (b)

19. (c)

20. (c)

WAEC, NECO & UTME Standard Theory Questions (1–10)

Topic: Atomic Structure

Question 1

(a) Define an atom.

(b) State the three subatomic particles found in an atom and indicate the location of each.

---

Question 2

(a) Define the following terms: (i) Atomic number (ii) Mass number

(b) Calculate the number of neutrons in an atom with mass number 56 and atomic number 26.

---

Question 3

(a) Define isotope.

(b) Explain the term isotopy.

---

Question 4

(a) Explain how a mass spectrometer is used in determining the relative atomic mass of an element.

(b) Calculate the relative atomic mass of chlorine containing 75% ³⁵Cl and 25% ³⁷Cl.

---

Question 5

(a) Define electronic configuration.

(b) Write the electronic configurations of: (i) Hydrogen (ii) Sodium (iii) Calcium

---

Question 6

(a) State Pauli's Exclusion Principle.

(b) State Hund's Rule of Maximum Multiplicity.

---

Question 7

(a) What is an orbital?

(b) List the four types of orbitals.

---

Question 8

(a) State the four quantum numbers.

(b) State the function of each quantum number.

---

Question 9

(a) State the octet rule.

(b) Explain how sodium and chlorine obey the octet rule during the formation of sodium chloride.

---

Question 10

(a) Define isoelectronic species.

(b) State five species that are isoelectronic with neon.

---

Answer Guide

Question 1

(a) An atom is the smallest particle of an element that can take part in a chemical reaction.

(b)

Proton — nucleus

Neutron — nucleus

Electron — shells around the nucleus

(c)

Proton is positively charged while neutron has no charge.

Proton determines atomic number while neutron contributes to mass number.

---

Question 2

(a) (i) Atomic number: Number of protons in an atom. (ii) Mass number: Sum of protons and neutrons.

(b) Neutrons = 56 − 26 = 30

---

Question 3

(a) Isotopes are atoms of the same element having the same atomic number but different mass numbers.

(b) Isotopy is the existence of isotopes of the same element.

Medical diagnosis and treatment

Carbon dating

Industrial thickness measurement

Tracers in research

Food irradiation

---

Question 4

(a) A mass spectrometer separates isotopes according to their mass-to-charge ratio and measures their relative abundance to determine the relative atomic mass.

(b) (75 × 35 + 25 × 37)/100 = 35.5

(c)

Determines isotopic abundance.

Determines molecular masses.

---

Question 5

(a) Electronic configuration is the arrangement of electrons in the orbitals of an atom.

(b) (i) H = 1s¹

(ii) Na = 1s² 2s² 2p⁶ 3s¹

(iii) Ca = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

---

Question 6

(a) No two electrons in an atom can have the same set of four quantum numbers.

(b) Electrons occupy degenerate orbitals singly before pairing.

---

Question 7

(a) An orbital is a region around the nucleus where there is a high probability of finding an electron.

(b)

s-orbital

p-orbital

d-orbital

f-orbital

Dumbbell-shaped

Three p-orbitals (px, py, pz)

Holds a maximum of six electrons

Begins from n = 2

---

Question 8

(a)

Principal quantum number (n)

Azimuthal quantum number (l)

Magnetic quantum number (mₗ)

Spin quantum number (mₛ)

(b)

n: Energy level

l: Shape of orbital

mₗ: Orientation of orbital

mₛ: Direction of electron spin

---

Question 9

(a) An atom becomes stable by having eight electrons in its outermost shell.

(b) Sodium loses one electron while chlorine gains one electron, producing Na⁺ and Cl⁻ with stable octet configurations.

---

Question 10

(a) Isoelectronic species are atoms, ions or molecules with the same number of electrons.

(b)

O²⁻

F⁻

Na⁺

Mg²⁺

(c) They have the same number of electrons but different numbers of protons (nuclear charges), leading to different chemical properties.

Certainly. Since 20 full WAEC/NECO/UTME theory questions with three sub-questions each and answers are too long for one response, here's how I'll present them:

Part 1: Questions 1–10 (already provided)

Part 2: Questions 11–20 with three sub-questions each and an answer guide.

---

WAEC, NECO & UTME THEORY QUESTIONS (11–20)

Question 11

(a) Define the nucleus of an atom.

(b) State the particles found in the nucleus.

---

Question 12

(a) State the five properties of the s-orbital.

(b) Draw the shape of an s-orbital.

---

Question 13

(a) State five properties of the p-orbital.

(b) Name the three p-orbitals.

---

Question 14

(a) State five properties of the d-orbitals.

(b) How many d-orbitals are present in a d-subshell?

---

Question 15

(a) State five properties of the f-orbitals.

(b) How many f-orbitals are present in an f-subshell?

---

Question 16

(a) State the maximum number of electrons that can be accommodated in the following shells: (i) K (ii) L (iii) M (iv) N

(b) State the formula used in calculating the maximum number of electrons in a shell.

---

Question 17

(a) Write the electronic configuration of the following elements: (i) Oxygen (ii) Aluminium (iii) Chlorine

(b) State the valence electrons of each element.

---

Question 18

(a) Differentiate between orbit and orbital.

(b) State three differences between them.

---

Question 19

(a) Calculate the relative atomic mass of an element having isotopes ²⁴X and ²⁶X in the ratio 3 : 2.

(b) Explain why the relative atomic mass obtained is not a whole number.

---

Question 20

(a) State the differences between: (i) Atom and ion (ii) Isotope and isobar

(b) Explain why isotopes have similar chemical properties but different physical properties.

---

ANSWER GUIDE

Question 11

(a) The nucleus is the dense central part of an atom.

(b) Protons and neutrons.

Positively charged.

Contains almost all the mass of the atom.

Very small compared with the whole atom.

Contains protons and neutrons.

Extremely dense.

---

Question 12

(a)

Spherical shape.

Symmetrical.

One orbital per energy level.

Holds two electrons.

Size increases with principal quantum number.

(b) Spherical shape.

---

Question 13

(a)

Dumbbell shape.

Three orbitals.

Equal energy.

Begins from n = 2.

Holds six electrons.

(b)

---

Question 14

(a)

Cloverleaf shape.

Five orbitals.

Holds ten electrons.

Begins from n = 3.

Important in transition elements.

(b) 5

---

Question 15

(a)

Complex shape.

Seven orbitals.

Holds fourteen electrons.

Begins from n = 4.

Found in lanthanides and actinides.

(b) 7

---

Question 16

(a)

K = 2

L = 8

M = 18

N = 32

(b) 2n²

(c)

2 × 5² = 50 electrons

---

Question 17

(a)

O = 1s² 2s² 2p⁴

Al = 1s² 2s² 2p⁶ 3s² 3p¹

Cl = 1s² 2s² 2p⁶ 3s² 3p⁵

(b)

O = 6

Al = 3

Cl = 7

(c)

O → Period 2, Group 16

Al → Period 3, Group 13

Cl → Period 3, Group 17

---

Question 18

(a)

Orbit: Fixed energy level around the nucleus.

Orbital: Region of high probability where an electron is found.

(b)

Orbit is two-dimensional; orbital is three-dimensional.

Orbit contains one or more orbitals.

Orbit is represented by shells; orbital is represented by s, p, d and f.

(c)

Both describe regions where electrons are found.

---

Question 19

(a)

(b)

Because it is a weighted average of isotopic masses.

(c)

Relative abundance of isotopes.

Mass of each isotope.

---

Question 20

(a)

(i) Atom: Neutral particle.

Ion: Charged particle formed by gain or loss of electrons.

(ii) Isotope: Same atomic number but different mass numbers.

Isobar: Same mass number but different atomic numbers.

(b)

They have the same number of electrons (same electronic configuration) but different numbers of neutrons.

(c)

Any four:

Cancer treatment (radiotherapy)

Medical diagnosis

Carbon dating

Industrial thickness measurement

Leak detection in pipelines

Food preservation and irradiation

Search This Blog

CHEMISTRY

SSS 1: ATOMIC STRUCTURE

Comments

Post a Comment

Popular posts from this blog

SSS 1: CURRICULUM

SSS 1: PARTICULATE NATURE OF MATTER (II)