Atomic Structure

Chemistry • Secondary School

1. Introduction

All substances are made from atoms. An atom is extremely small, but its internal structure explains many chemical properties of elements. Understanding atomic structure is therefore one of the foundations of Chemistry.

Atom: The smallest electrically neutral particle of an element that retains the chemical properties of that element.

Atoms contain three main subatomic particles: protons, neutrons and electrons.

Learning objectives

Identify the three main subatomic particles.
State their relative masses and charges.
Describe the structure of an atom.
Calculate proton, neutron and electron numbers.
Explain isotopes.
Write electronic configurations.
Relate electronic structure to chemical behaviour.

2. Subatomic Particles

Particle	Symbol	Relative charge	Relative mass	Location
Proton	p⁺	+1	1	Nucleus
Neutron	n⁰	0	1	Nucleus
Electron	e⁻	−1	1/1840 approximately	Electron shells

Remember: The nucleus contains protons and neutrons. Electrons occupy regions around the nucleus called shells or energy levels.

3. Atomic Number

Atomic number (Z): The number of protons in the nucleus of an atom.

Every element has its own atomic number. Therefore, the number of protons determines the identity of an element.

\(Z=\text{number of protons}\)

For a neutral atom, the number of electrons is equal to the number of protons.

Example: An atom has atomic number 11. How many protons and electrons does it contain?

Solution:

Protons = 11.

Because the atom is neutral:

Electrons = 11

The element is sodium, Na.

4. Mass Number

Mass number (A): The total number of protons and neutrons in the nucleus.

\(A=Z+N\)

Therefore:

\(N=A-Z\)

where \(N\) is the number of neutrons.

Example: An atom of chlorine has mass number 35 and atomic number 17. Calculate the number of neutrons.

Solution:

\(N=A-Z\)

\(N=35-17=18\)

Therefore, the atom contains 18 neutrons.

5. Nuclide Notation

An atom can be represented using the notation:

\[ {}^{A}_{Z}X \]

where:

X = chemical symbol of the element
A = mass number
Z = atomic number

Example:

\[ {}^{23}_{11}\mathrm{Na} \]

This represents sodium with:

11 protons
11 electrons if neutral
12 neutrons

6. Isotopes

Isotopes: Atoms of the same element that have the same number of protons but different numbers of neutrons.

Because isotopes have the same number of protons, they have the same atomic number. They differ in mass number because their numbers of neutrons are different.

Example: Chlorine isotopes

\[ {}^{35}_{17}\mathrm{Cl} \qquad {}^{37}_{17}\mathrm{Cl} \]

Both have 17 protons, but the first has 18 neutrons while the second has 20 neutrons.

Feature	Same for isotopes	Different for isotopes
Protons	✓
Atomic number	✓
Electrons in neutral atoms	✓
Neutrons		✓
Mass number		✓

7. Relative Atomic Mass

Many elements exist naturally as mixtures of isotopes. The relative atomic mass takes account of the masses and abundances of the different isotopes.

\[ A_r= \frac{\sum(\text{isotopic mass}\times\text{percentage abundance})} {100} \]

Example: Chlorine contains 75% chlorine-35 and 25% chlorine-37. Calculate its relative atomic mass.

\[ A_r= \frac{(35\times75)+(37\times25)}{100} \]

\[ A_r=\frac{2625+925}{100}=35.5 \]

Therefore, the relative atomic mass is approximately 35.5.

Examination focus: Always convert percentage abundance into the weighted calculation. Do not simply find the ordinary average unless the isotopes have equal abundance.

8. Electron Arrangement

Electrons occupy energy levels around the nucleus. For the early elements commonly studied at secondary level, the shell arrangement can be represented using numbers such as:

Shell	Maximum electrons commonly used in basic shell notation
1st shell	2
2nd shell	8
3rd shell	8 for the first 20 elements in simple school-level notation
4th shell	Begins to fill for elements beyond the first 20

Example: Sodium

Sodium has atomic number 11, so a neutral sodium atom has 11 electrons.

\[ 2,\;8,\;1 \]

The outer shell contains one electron.

Example: Chlorine

Chlorine has atomic number 17.

\[ 2,\;8,\;7 \]

Therefore chlorine has seven electrons in its outer shell.

9. Valence Electrons

Valence electrons: Electrons in the outermost occupied electron shell of an atom.

Valence electrons are particularly important because they are involved in chemical bonding and strongly influence chemical reactivity.

Element	Atomic number	Electron arrangement	Valence electrons
Li	3	2,1	1
C	6	2,4	4
O	8	2,6	6
Ne	10	2,8	8
Na	11	2,8,1	1
Cl	17	2,8,7	7

10. Formation of Ions

Ion: An electrically charged particle formed when an atom or group of atoms gains or loses electrons.

Cations

A cation is a positively charged ion formed when an atom loses one or more electrons.

\[ \mathrm{Na}\rightarrow\mathrm{Na}^{+}+e^{-} \]

Anions

An anion is a negatively charged ion formed when an atom gains one or more electrons.

\[ \mathrm{Cl}+e^{-}\rightarrow\mathrm{Cl}^{-} \]

Important: Atoms become ions by changing their number of electrons. The number of protons does not change during ordinary ion formation.

11. Electronic Configuration and Chemical Behaviour

Atoms tend to form stable electron arrangements. Main-group atoms may lose, gain or share electrons to obtain a more stable outer shell.

Example:

Sodium has the arrangement \(2,8,1\). It can lose one electron to form \(Na^+\), giving the stable arrangement \(2,8\).

Chlorine has the arrangement \(2,8,7\). It can gain one electron to form \(Cl^-\), giving the stable arrangement \(2,8,8\).

This helps explain why sodium and chlorine react to form sodium chloride.

\[ \mathrm{Na}^{+}+\mathrm{Cl}^{-}\rightarrow\mathrm{NaCl} \]

12. Historical Development of the Atomic Model

Scientist/model	Main idea
Dalton	Matter was described as being made of atoms; atoms of each element were considered distinct.
J. J. Thomson	Discovered the electron and proposed a model containing electrons in positive matter.
Rutherford	Evidence from scattering experiments supported a small, dense, positively charged nucleus.
Bohr	Electrons were proposed to occupy specific energy levels.
Modern model	Electrons occupy regions of probability called orbitals rather than simple fixed circular paths.

Exam tip: When comparing atomic models, focus on what new evidence caused the model to change.

13. Worked Examination Questions

Question 1: An atom has 19 protons and 20 neutrons. Determine its atomic number and mass number.

Solution:

Atomic number = number of protons = 19.

\[ A=19+20=39 \]

Mass number = 39.

Question 2: An ion contains 17 protons and 18 electrons. Determine its charge.

Solution:

There is one more electron than proton.

\[ 17-18=-1 \]

Therefore the ion has charge −1, written as \(Cl^-\).

Question 3: Write the electron arrangement of magnesium, atomic number 12.

Solution:

\[ 2,\;8,\;2 \]

14. MANEB Examination Focus

Define atomic number and mass number.
Calculate numbers of protons, neutrons and electrons.
Interpret nuclide notation.
Explain isotopes.
Calculate relative atomic mass from isotope abundance.
Write electron arrangements.
Identify valence electrons.
Explain formation of cations and anions.
Relate electronic arrangement to chemical behaviour.
Compare historical atomic models.

Common mistakes: Do not confuse atomic number with mass number. Atomic number counts protons only, while mass number counts protons plus neutrons.

15. Practice Exercises

1. Define an atom.

2. State the charge and relative mass of a proton.

3. State the charge of an electron.

4. An atom has atomic number 13 and mass number 27. Calculate its numbers of protons, neutrons and electrons.

5. Explain why isotopes have the same chemical properties but different masses.

6. Write the electron arrangement of oxygen, atomic number 8.

7. Write the electron arrangement of calcium, atomic number 20.

8. Explain how a sodium atom forms \(Na^+\).

9. Explain how chlorine forms \(Cl^-\).

10. An element has two isotopes of masses 10 and 11 with abundances of 20% and 80%, respectively. Calculate its relative atomic mass.

Answers

An atom is the smallest electrically neutral particle of an element that retains its chemical properties.
Charge = +1; relative mass = 1.
−1.
Protons = 13, electrons = 13, neutrons = 14.
They have the same number of protons and therefore the same atomic identity and similar chemical behaviour, but different numbers of neutrons give different masses.
\(2,6\).
\(2,8,8,2\).
Sodium loses one outer electron to form \(Na^+\).
Chlorine gains one electron to form \(Cl^-\).
\[ A_r=\frac{(10\times20)+(11\times80)}{100}=10.8 \]

16. Examination Tips

Memorise the definitions of atomic number, mass number and isotopes.
Remember: protons identify the element.
Use \(N=A-Z\) to calculate neutrons.
For a neutral atom, protons = electrons.
For ions, compare proton and electron numbers to determine charge.
Check electron arrangements carefully before predicting ion formation.
Show all working in numerical questions.

17. Topic Summary

Atoms consist of a tiny nucleus containing protons and neutrons, surrounded by electrons in energy levels. The number of protons is the atomic number, while the total number of protons and neutrons is the mass number.

Isotopes have the same atomic number but different mass numbers. Electron arrangement, especially the number of valence electrons, helps explain how elements form ions and chemical bonds.