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
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:
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
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
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
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
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
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
OR
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
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
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