Learning goals. After studying these notes you should
be able to:
- Explain
what an atom is and why the idea of the atom changed over time.
- Describe
the experiments of Thomson, Rutherford and Bohr and what each proved.
- Name
the three subatomic particles and state their charge, mass and location.
- Calculate
protons, neutrons and electrons from atomic number and mass number.
- Write
electron configurations using shells and subshells.
- Distinguish
isotopes, isobars and isotones, and ions from neutral atoms.
Describe periodic trends and radioactive decay at a basic level.
What is atom?
An atom is the smallest unit of ordinary matter that has the properties of a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms are extremely small, typically around 100 picometers across. A human hair is about a million carbon atoms wide.There are about 118 known elements, and each one is made of its own kind of atom.
The word "atom" comes from the Greek word atomos, meaning "uncuttable". The ancient Greeks believed matter could not be divided forever. We now know atoms can be divided into smaller particles, but if you split an atom of gold into its particles, it is no longer gold.Atoms are composed of a central nucleus that is surrounded by a cloud of electrons. The nucleus is made up of protons and neutrons, which are particles that are much heavier than electrons. The number of protons in the nucleus determines the atomic number of the atom, which is what identifies the element. The number of neutrons in the nucleus determines the isotope of the atom.
Electrons are held in place around the nucleus by the electrostatic attraction between the negatively charged electrons and the positively charged protons. Electrons can only exist in certain energy levels or orbitals. When an electron absorbs energy, it can jump to a higher energy level. When an electron loses energy, it can jump to a lower energy level. This movement of electrons is responsible for the emission of light.
Atoms can combine with other atoms to form molecules. Molecules are held together by chemical bonds, which are the forces that attract atoms to each other. There are three main types of chemical bonds: ionic bonds, covalent bonds, and metallic bonds.
Ionic bonds are formed between atoms that have different electronegativities. Electronegativity is a measure of an atom's ability to attract electrons. When an atom with a high electronegativity attracts electrons from an atom with a low electronegativity, the two atoms become oppositely charged and are attracted to each other.
Covalent bonds are formed between atoms that share electrons. When two atoms share electrons, they create a strong attraction between each other.
Metallic bonds are formed between atoms that are arranged in a regular lattice structure. The electrons in a metallic bond are free to move throughout the lattice, which allows the metal to conduct electricity and heat.
Atoms are the building blocks of all matter. They are incredibly small, but they are the basis for all of the chemistry that we see around us
How small is an atom?
- A
typical atom has a diameter of about 0.1 to 0.5 nanometres (1 nm = 10⁻⁹
m).
- About
5 million hydrogen atoms placed side by side would stretch across the
thickness of one full stop (period) on this page.
- A
single drop of water contains more than a billion billion (10²¹) atoms.
- Atoms
are mostly empty space. If the nucleus were the size of a marble placed at
the centre of a football stadium, the electrons would be like tiny specks
near the outer seats.
Because atoms are so small, scientists measure them using
special units. The atomic mass unit (u) is used for mass, and the picometre
(pm) or angstrom (Å) is used for length (1 Å = 10⁻¹⁰ m).
History of the Atom: How Ideas Changed
Science progresses when new evidence disproves older models.
The story of the atom is one of the best examples.
1 Democritus and Leucippus (about 400 BCE)
These Greek philosophers proposed that matter is made of
tiny, indivisible particles moving in empty space. They had no experiments to
prove it. It was a thoughtful idea, not yet a scientific theory.
2 John Dalton's Atomic Theory (1803)
Dalton revived the atomic idea using experimental evidence
from chemical reactions. His main postulates were:
- All
matter is made of extremely small indivisible particles called atoms.
- Atoms
of the same element are identical in mass and properties.
- Atoms
of different elements differ in mass and properties.
- Atoms
combine in simple whole-number ratios to form compounds.
- Atoms
are neither created nor destroyed in a chemical reaction (law of
conservation of mass).
What Dalton got right: the law of conservation of
mass, the law of definite proportions, and the idea of atoms combining in fixed
ratios.
What was later corrected: atoms are divisible into
subatomic particles, and atoms of the same element can differ in mass
(isotopes).
3 J.J. Thomson and the Electron (1897)
Thomson experimented with cathode ray tubes, glass tubes
with almost no air and two metal plates connected to high voltage. He observed
a beam moving from the negative plate (cathode) to the positive plate (anode).
The beam bent toward a positively charged plate, proving that it was made of negatively
charged particles. These were later named electrons. The beam was
the same no matter which gas or metal he used, so electrons are present in all
atoms.
Thomson's plum pudding model (1904): Since atoms are
neutral overall, there must be positive charge to balance the electrons.
Thomson proposed that an atom is a sphere of positive charge with electrons
embedded inside it, like plums in a pudding, or seeds in a watermelon.
PLUM PUDDING MODEL (Thomson)
.-"""""""""""""""""-.
.' + + - + + - + '.
/ + - + + + - + + - \
| + + + - + + + + + |
| - + + + + - + + - |
\ + + - + + + - + + /
'. + - + + + + - .'
'-.....................-'
(+) = positive charge spread everywhere
(-) = electrons scattered inside
4 Goldstein and the Proton (1886)
Eugen Goldstein used a modified cathode ray tube with a perforated cathode and
found rays moving in the opposite direction. These canal rays were positively
charged. The lightest positive particle, from hydrogen, was later named the proton
(identified by Rutherford in 1917–1919).
5 Rutherford's Gold Foil Experiment (1909–1911)
Ernest Rutherford, with his students Hans Geiger and Ernest Marsden, fired
fast-moving, positively charged alpha (α) particles at a very thin sheet of gold foil,
about 100 nm thick. A circular screen coated with zinc sulphide around the foil
glowed with a flash wherever an alpha particle hit it.
RUTHERFORD'S GOLD FOIL EXPERIMENT
Alpha source Gold foil Detector screen
+-------+ | (ZnS coated)
| a --->| - ------------------ - - - -> | - - - - ------------------ - - - -> (O) most pass straight
| | |
| a --->| - - - ------------------ - -> |\
| | | \ . . ........................... . > (O) a few deflect (small angle)
| a --->| - - ------------------ - - -> (N) <- - - ------------------- - (O) very few bounce back
+-------+ |
(nucleus)
Observations and conclusions
Observation
Conclusion
Most alpha particles passed straight through
Most of the atom is empty space
Some particles were deflected by small angles
Positive charge exists in a small region that repels alpha
particles
Very few (about 1 in 20,000) bounced almost straight back
The positive charge and most of the mass is concentrated
in a tiny, dense centre called the nucleus
Rutherford said it was "as if you had fired a 15-inch shell at a piece of tissue paper
and it came back and hit you."
Rutherford's nuclear model:
- An atom has a tiny, dense, positively charged nucleus at the centre containing nearly all
the mass.
- Electrons revolve around the nucleus in circular paths, like planets around the Sun.
- The size of the nucleus is very small compared to the atom (about 1/100,000 of the
atom's diameter).
Drawback: According to classical physics, an electron moving in a circle is
accelerating and should continuously lose energy, spiralling into the nucleus.
That would make atoms unstable, which is not what we observe.
6 Niels Bohr's Model (1913)
Bohr fixed Rutherford's instability problem with these postulates:
- Electrons revolve in certain fixed circular paths called orbits or energy levels
(shells) without radiating energy. These are stationary orbits. - Each orbit has a definite energy. They are labelled n = 1, 2, 3... or K, L, M, N...
starting nearest the nucleus.
- Energy of an orbit increases with distance from the nucleus.
- An electron absorbs energy to jump to a higher level, and releases energy
(as light) when it falls to a lower level. The energy difference is ΔE = hν, where
h is Planck's constant and ν is the frequency of light.
- Only orbits where angular momentum equals nh/2π are allowed.
BOHR MODEL (concentric shells)
n=3 (M)
.---------------.
/ n=2 (L) \
/ .-----------. \
| / n=1 (K) \ \
| | .-----. \ |
| | | (N) | \ |
| | '-----' / |
| \ / /
\ '-----------' / /
\ /
'---------------'
N = nucleus (protons + neutrons)
Electrons sit on the circular shells
Success of Bohr's model: It explained the line spectrum of hydrogen and why atoms are
stable.
Limitations: It worked well only for hydrogen-like species (one electron). It could not
explain spectra of multi-electron atoms, the splitting of spectral lines in magnetic fields
(Zeeman effect) or electric fields (Stark effect), or chemical bonding properly. It also
violated the Heisenberg uncertainty principle by assigning exact paths to electrons.
2.7 The Quantum Mechanical Model (1920s to present)
Scientists including de Broglie, Heisenberg, Schrödinger and Born developed the
modern model.
- de Broglie (1924): Electrons show wave-particle duality. They behave as both particles
and waves. - Heisenberg uncertainty principle (1927): It is impossible to know the exact
position and exact momentum of an electron at the same time.
- Schrödinger equation (1926): Describes electron behaviour as a wave function (ψ).
The square of the wave function, ψ², gives the probability of finding an electron at a point.
So instead of fixed orbits, we have orbitals: three-dimensional regions around the nucleus
where there is a high probability (usually 90 to 95 percent) of finding an electron.
ORBIT (Bohr) ORBITAL (Quantum)
.-"""-. . : : : : .
/ e \ : : : * : : :
| * | : : : : : : :
\ / ' : : : : '
'--------'
Fixed circular path Cloud of probability
Timeline Summary
Year
Scientist
Contribution
~400 BCE
Democritus
Idea of indivisible atoms
1803
Dalton
Atomic theory based on experiments
1897
Thomson
Discovered electron; plum pudding model
1886 / 1917
Goldstein / Rutherford
Canal rays; discovery of proton
1911
Rutherford
Nucleus discovered; nuclear model
1913
Bohr
Quantised orbits
1932
Chadwick
Discovered neutron
1920s
Schrödinger, Heisenberg
Quantum mechanical model
3. Subatomic Particles
An atom consists of a nucleus (made of protons and neutrons, together called nucleons)
surrounded by electrons.
Property
Electron (e⁻)
Proton (p⁺)
Neutron (n⁰)
Discovered by
J.J. Thomson (1897)
Rutherford (1917–19)
James Chadwick (1932)
Charge
−1 (−1.6 × 10⁻¹⁹ C)
+1 (+1.6 × 10⁻¹⁹ C)
0
Relative mass
1/1836
1
1
Mass (kg)
9.11 × 10⁻³¹
1.673 × 10⁻²⁷
1.675 × 10⁻²⁷
Mass (u)
0.00055
1.0073
1.0087
Location
Outside nucleus, in shells/orbitals
Inside nucleus
Inside nucleus
ATOM DIAGRAM (Helium-4)
e-
. ' .
.' '.
/ +---------------+ \
| | p n | |
e- -- | | n p | | -- (shell path)
| +-----------------+ |
\ nucleus /
'. .'
' . .....................'
Helium: 2 protons, 2 neutrons, 2 electrons
Key points to remember
- Atoms are electrically neutral because the number of protons equals the
number of electrons.
- The mass of the atom is concentrated in the nucleus because electrons are
almost massless in comparison.
- Protons and neutrons are themselves made of smaller particles called quarks
(a proton has two up quarks and one down quark; a neutron has one up and
two down quarks). At school level you rarely need this, but it is good
general knowledge.
- Neutrons act like "glue". They add strong nuclear attraction without
adding electrical repulsion, which helps hold the positively charged
protons together.
- The only atom with no neutron in its most common form is hydrogen-1 (protium).
4. Atomic Number, Mass Number and Notation
4.1 Atomic Number (Z)
The atomic number is the number of protons in the nucleus. It defines the element.
Every atom with 6 protons is carbon, whatever its mass. In a neutral atom, Z also
equals the number of electrons.
4.2 Mass Number (A)
The mass number is the total number of protons plus neutrons: A = Z + N, so N = A − Z
4.3 Standard Notation
A
X A = mass number (protons + neutrons)
Z Z = atomic number (protons)
X = chemical symbol
Example: 23 Na -> A = 23, Z = 11
11
Protons = 11, Electrons = 11, Neutrons = 23 - 11 = 12
4.4 Quick Reference Examples
Atom
Z
A
Protons
Neutrons
Electrons
Hydrogen (¹H)
1
1
1
0
1
Carbon (¹²C)
6
12
6
6
6
Oxygen (¹⁶O)
8
16
8
8
8
Sodium (²³Na)
11
23
11
12
11
Chlorine (³⁵Cl)
17
35
17
18
17
Calcium (⁴⁰Ca)
20
40
20
20
20
Uranium (²³⁸U)
92
238
92
146
92
4.5 Atomic Mass vs Mass Number
Mass number is always a whole number. Atomic mass (relative atomic mass) is the
weighted average mass of all naturally occurring isotopes of an element compared
with 1/12 the mass of a carbon-12 atom, so it is usually a decimal. For example,
chlorine has an atomic mass of about 35.5 because it exists as roughly 75 percent ³⁵Cl
and 25 percent ³⁷Cl.
Worked example: average atomic mass Boron has two isotopes: ¹⁰B (19.9 percent)
and ¹¹B (80.1 percent).
Average = (10 × 0.199) + (11 × 0.801) = 1.99 + 8.811 = 10.80u
5. Isotopes, Isobars and Isotones
5.1 Isotopes
Isotopes are atoms of the same element with the same atomic number but different mass
numbers, because they have different numbers of neutrons. They have nearly identical
chemical properties (same electrons) but slightly different physical properties (different mass).
HYDROGEN ISOTOPES
Protium 1H Deuterium 2H Tritium 3H
e- e- e-
* * *
(p) (p n) (p n n)
1 p, 0 n 1 p, 1 n 1 p, 2 n
99.98% natural 0.015% natural trace, radioactive
Other examples: carbon-12, carbon-13 and carbon-14; chlorine-35 and chlorine-37;
uranium-235 and uranium-238.
Uses of isotopes
- Carbon-14: radiocarbon dating of fossils and ancient objects (half-life about 5,730
years).
- Cobalt-60: cancer radiotherapy and sterilising medical equipment.
- Iodine-131: treatment of thyroid disorders.
- Uranium-235: fuel in nuclear reactors.
- Deuterium: heavy water used as a moderator in some nuclear reactors.
5.2 Isobars
Isobars are atoms of different elements that have the same mass number but different atomic
numbers. Example: ⁴⁰Ar (Z = 18) and ⁴⁰Ca (Z = 20) both have A = 40.
5.3 Isotones
Isotones are atoms of different elements with the same number of neutrons. Example:
³⁹K (19 protons, 20 neutrons) and ⁴⁰Ca (20 protons, 20 neutrons) both have 20 neutrons.
5.4 Memory Trick
- Isotopes: same total protons (same element).
- Isobars: same baryon (nucleon) count, meaning the same mass number.
- Isotones:same neutron count ("n" sounds like "tone").
Type
Same
Different
Isotopes
Atomic number (Z)
Mass number (A), neutrons
Isobars
Mass number (A)
Atomic number (Z)
Isotones
Number of neutrons
Z and A
6. Electron Arrangement in Atoms
6.1 Shells and the 2n² Rule
Electrons are placed in shells (energy levels) around the nucleus. The maximum number of
electrons in a shell is given by 2n²,where n is the shell number.
Shell
n
Maximum electrons (2n²)
K
1
2
L
2
8
M
3
18
N
4
32
6.2 Bohr–Bury Rules
- The maximum number of electrons in a shell is 2n².
- The outermost shell (valence shell) can hold at most 8 electrons (2 for the first shell).
- The second-last shell can hold at most 18 electrons.
- Electrons fill lower energy shells first, then higher ones.
6.3 Shell Arrangement of the First 20 Elements
Element Z K L M N Short form
Hydrogen 1 1 1
Helium 2 2 2
Lithium 3 2 1 2,1
Beryllium 4 2 2 2,2
Boron 5 2 3 2,3
Carbon 6 2 4 2,4
Nitrogen 7 2 5 2,5
Oxygen 8 2 6 2,6
Fluorine 9 2 7 2,7
Neon 10 2 8 2,8
Sodium 11 2 8 1 2,8,1
Magnesium 12 2 8 2 2,8,2
Aluminium 13 2 8 3 2,8,3
Silicon 14 2 8 4 2,8,4
Phosphorus 15 2 8 5 2,8,5
Sulphur 16 2 8 6 2,8,6
Chlorine 17 2 8 7 2,8,7
Argon 18 2 8 8 2,8,8
Potassium 19 2 8 8 1 2,8,8,1
Calcium 20 2 8 8 2 2,8,8,2
Drawing a shell diagram: carbon (2,4)
e- e-
.--------------.
/ .-------------. \ |
| / e- \ |
e- | | (6p | | e-
| | 6n) | |
| \ e- / |
\ \'--------------' / |
'-----------------'
e- e-
Inner shell (K): 2 electrons
Outer shell (L): 4 electrons (valence electrons)
6.4 Valence Electrons and Valency
Electrons in the outermost shell are valence electrons.They take part in chemical bonding.
Valency is the combining capacity of an atom.
- If valence electrons are 1 to 4, valency equals the number of valence electrons
(Na = 1, Mg = 2, Al = 3, Si = 4). - If valence electrons are 5 to 7, valency = 8 − valence electrons (N = 3, O = 2, Cl = 1).
- If the outer shell is full (2 or 8), valency = 0 (noble gases such as He, Ne
and Ar).
Octet rule: Atoms tend to gain, lose or share electrons to achieve eight electrons in
their outer shell (a duplet for hydrogen and helium), the stable configuration of noble
gases.
7. Modern View: Subshells, Orbitals and Quantum Numbers
The quantum mechanical model describes each electron by four quantum numbers, which act
like the "address" of the electron.
Quantum number
Symbol
Describes
Allowed values
Principal
n
Shell (size and energy)
1, 2, 3...
Azimuthal
l
Subshell (shape)
0 to n − 1
Magnetic
mₗ
Orbital orientation
−l to +l
Spin
mₛ
Spin of electron
+½ or −½
7.1 Subshells and Orbital Shapes
Subshell
l
Number of orbitals
Max electrons
Shape
s
0
1
2
Spherical
p
1
3
6
Dumbbell
d
2
5
10
Cloverleaf (mostly)
f
3
7
14
Complex
ORBITAL SHAPES
s orbital p orbitals (px, py, pz)
.--. |
/ \ O--+--O
| * | | (px along x-axis,
\ / py along y-axis,
'--' pz along z-axis)
Sphere around the nucleus Each p orbital is a dumbbell on one axis
7.2 Rules for Filling Electrons
- Aufbau principle: Electrons occupy the lowest-energy orbitals first.
- Pauli exclusion principle: No two electrons in an atom can have the same set
of four quantum numbers. Each orbital holds a maximum of two electrons,
with opposite spins.- Hund's rule: Within a subshell, electrons occupy empty orbitals singly first,with
parallel spins, before pairing.
7.3 Order of Filling
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
Read diagonally (arrows from top right to bottom left):
1s > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f > 5d > 6p > 7s...
7.4 Electron Configuration Examples
Element
Z
Configuration
Hydrogen
1
1s¹
Carbon
6
1s² 2s² 2p²
Nitrogen
7
1s² 2s² 2p³
Oxygen
8
1s² 2s² 2p⁴
Neon
10
1s² 2s² 2p⁶
Sodium
11
1s² 2s² 2p⁶ 3s¹
Chlorine
17
1s² 2s² 2p⁶ 3s² 3p⁵
Calcium
20
1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Iron
26
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
Orbital box diagram for nitrogen (Hund's rule)
1s 2s 2p
[^v] [^v] [^][^][^]
Each 2p orbital gets one electron before any pairing.
Orbital box diagram for oxygen
1s 2s 2p
[^v] [^v] [^v][^][^]
Exceptions: Chromium (3d⁵ 4s¹) and copper (3d¹⁰ 4s¹) are exceptions because half-filled
and completely filled d subshells have extra stability.
8. Energy Levels and Atomic Spectra
When an electron absorbs energy (heat, electricity or light), it jumps to a higher energy level.
This is the excited state. It quickly returns to a lower level and releases the extra energy as a
photon of light. The colour of the light depends on the energy difference.
ENERGY LEVEL DIAGRAM (Hydrogen)
Energy
^ n=4 ------------------
| n=3 ------------------
| ^ absorb v emit photon
| n=2 ------------------
|
|
| n=1 ------------------ (ground state)
Energy gap between levels gets smaller as n increases.
Hydrogen spectral series
Series
Falls to
Region
Lyman
n = 1
Ultraviolet
Balmer
n = 2
Visible
Paschen
n = 3
Infrared
Brackett
n = 4
Infrared
Pfund
n = 5
Infrared
Each element has a unique pattern of spectral lines, like a fingerprint. Astronomers use these
patterns to find out which elements are present in distant stars. The bright colours of fireworks
and flame tests come from electron transitions: sodium glows yellow, copper blue-green,
lithium crimson-red and potassium lilac.
9. Ions: Atoms That Gained or Lost Electrons
An ion is a charged atom formed when an atom gains or loses electrons. The number
of protons never changes.
- Cation (positive ion): formed by losing electrons. Metals usually form
cations. Na → Na⁺ + e⁻.
- Anion (negative ion): formed by gaining electrons. Non-metals usually form
anions. Cl + e⁻ → Cl⁻.
SODIUM ATOM SODIUM ION (Na+)
Protons: 11 Protons: 11
Electrons: 11 Electrons: 10
Config: 2,8,1 Config: 2,8 (like neon)
Charge: 0 Charge: +1
CHLORINE ATOM CHLORIDE ION (Cl-)
Protons: 17 Protons: 17
Electrons: 17 Electrons: 18
Config: 2,8,7 Config: 2,8,8 (like argon)
Charge: 0 Charge: -1
Species
Protons
Neutrons
Electrons
Charge
²³Na
11
12
11
0
²³Na⁺
11
12
10
+1
³⁵Cl
17
18
17
0
³⁵Cl⁻
17
18
18
−1
²⁷Al³⁺
13
14
10
+3
¹⁶O²⁻
8
8
10
−2
Isoelectronic species have the same number of electrons. Na⁺, Mg²⁺, F⁻, O²⁻ and Ne all
have 10 electrons.
Ionic bonding preview: When sodium (which wants to lose one electron) meets chlorine
(which wants to gain one), an electron transfers and the oppositely charged ions attract to
form sodium chloride (table salt).
10. The Atom and the Periodic Table
The modern periodic table arranges elements by increasing atomic number. Elements in
the same group (vertical column) have the same number of valence electrons, so similar
chemical behaviour. Elements in the same period (horizontal row) have the same number
of occupied shells.
Group 1 Group 17 Group 18
Li (2,1) F (2,7) He (2)
Na (2,8,1) Cl (2,8,7) Ne (2,8)
K (2,8,8,1) Br (...,7) Ar (2,8,8)
Same outer electrons = same chemical family
Trends Across a Period and Down a Group
Property
Across a period (left to right)
Down a
Reason
Atomic radius
Decreases
Increases
Across: more nuclear charge pulls electrons in. Down:
extra shells are added.
Ionisation energy
Increases
Decreases
Harder to remove electrons when held tightly or when
closer to the nucleus.
Electronegativity
Increases
Decreases
Same reasoning as above.
Metallic character
Decreases
Increases
Easier loss of electrons down a group.
Ionisation energy is the energy needed to remove the outermost electron from a gaseous
atom. Electron affinity is the energy change when an atom gains an electron.
Electronegativity is the tendency of an atom to attract bonding electrons,
11. Forces Inside the Atom
Four fundamental forces operate in nature. Two are important in the atom:
- Electromagnetic force: attracts negative electrons to the positive nucleus, and makes
protons repel each other.
- Strong nuclear force: a very powerful but very short-range attraction between
nucleons that overcomes proton–proton repulsion and holds the nucleus
together.
Gravity is negligible at the atomic scale, and the weak force is involved in certain types of
radioactive decay.
Nuclear stability. Small nuclei are stable when neutrons and protons are about equal. Larger
nuclei need more neutrons than protons. Nuclei beyond bismuth-209 (Z = 83) are all
radioactive because the strong force cannot hold so many protons together permanently.
12. Radioactivity and Nuclear Changes
Unstable nuclei emit radiation to become more stable. This spontaneous process is
radioactive decay.
Type
Symbol
What is emitted
Charge
Penetrating power
Stopped by
Alpha
α
Helium nucleus (2p + 2n)
+2
Low
Paper, skin
Beta
β⁻
Fast electron
−1
Medium
Aluminium sheet
Gamma
γ
High-energy electromagnetic wave
0
Very high
Thick lead, concrete
Penetrating Power
alpha ---------------->| stopped by paper
beta ---------------->| stopped by aluminium
gamma ------------------------>| stopped by thick lead
Alpha decay example: ²³⁸U → ²³⁴Th + ⁴He (mass number falls by 4, atomic number by 2).
Beta decay example: ¹⁴C → ¹⁴N + e⁻ + antineutrino (a neutron changes into a proton; atomic
number rises by 1, mass number is unchanged).
Half-life is the time taken for half of the radioactive atoms in a sample to decay. If the half-life
of a substance is 10 days and you start with 80 g, after 10 days you have 40 g, after 20 days
20 g, after 30 days 10 g.
Amount remaining
100% | ########
50% | ####
25% | ##
12.5% | #
+---+---+---+---> number of half-lives
1 2 3
Nuclear fission and fusion
- Fission: a heavy nucleus (such as uranium-235) splits into smaller nuclei, releasing
large energy. Used in nuclear power stations.
- Fusion: light nuclei (such as hydrogen isotopes) join to form a heavier one, releasing
even more energy. It powers the Sun and other stars.Einstein's equation E = mc² explains the energy: a tiny amount of mass converts to a huge
amount of energy.
13. Atoms, Molecules and the Mole
- A molecule is a group of two or more atoms chemically bonded together (O₂, H₂O, CO₂).
- A compound contains atoms of different elements combined in a fixed ratio.
- The mole is the amount of substance that contains 6.022 × 10²³ particles (Avogadro's
number).
Key relations
- Number of moles = mass in grams ÷ molar mass.
- Number of particles = moles × 6.022 × 10²³.
- Molar mass in g/mol is numerically equal to atomic or molecular mass in u.
Worked example.
How many atoms are in 12 g of carbon?
Moles = 12 ÷ 12 = 1 mol. Atoms = 1 × 6.022 × 10²³ = 6.022 × 10²³ atoms.
Worked example.
How many molecules are in 18 g of water? Molar mass of H₂O = (2 × 1) + 16 = 18 g/mol.
Moles = 18 ÷ 18 = 1. Molecules = 6.022 × 10²³. Since each molecule has 3 atoms,
total atoms = 1.8066× 10²⁴.
14. Video Learning Guide:-
Watching a good animation after reading these notes helps your brain connect words to moving
pictures. Since video links change over time, use the search phrases below on YouTube or on
the channel's own website. Trusted educational channels include Khan Academy, Crash
Course Chemistry, TED-Ed, Kurzgesagt, Veritasium, Professor Dave Explains, The
Organic Chemistry Tutor, Physics Wallah and Vedantu (useful for exam-style explanations).
Topic in these notes
Suggested search phrase
What to look for
Approx. length
Sections 1 and 2: History of the atom
"TED-Ed history of the atom"
Timeline animation from Democritus to the quantum model
5 to 6 min
Section 2.5: Rutherford
"Rutherford gold foil experiment animation"
Alpha particles passing through and bouncing back
3 to 5 min
Sections 3 and 4: Subatomic particles
"Khan Academy atomic number mass number"
How to read an element box and compute neutrons
8 to 10 min
Section 5: Isotopes
"Isotopes explained Crash Course Chemistry"
Hydrogen isotopes and average atomic mass
5 to 10 min
Sections 6 and 7: Electron configuration
"Electron configuration Organic Chemistry Tutor"
Aufbau, Pauli and Hund's rule with worked examples
15 to 20 min
Section 7: Orbitals
"Atomic orbitals s p d f 3D animation"
Rotating orbital shapes
4 to 8 min
Section 8: Spectra
"Hydrogen emission spectrum Bohr model
animation"
Electron jumps and coloured lines
5 to 8 min
Section 12: Radioactivity
"Alpha beta gamma decay animation" and
"half-life explained"
Decay modes and the half-life curve
5 to 10 min
Whole-topic overview
"Structure of an atom class 9"
Full chapter revision
20 to 30 min
Interactive simulation (free): The PhET Interactive Simulations project from the
University of Colorado Boulder offers "Build an Atom", "Rutherford Scattering" and "Models
of the Hydrogen Atom". You add protons, neutrons and electrons yourself and watch the
element, charge and mass change in real time. It is one of the best ways to understand the
atom by experimenting.
How to watch actively
- Before the video, write two questions you hope it will answer.
- Pause at every diagram and sketch it yourself.
- After the video, write a three-sentence summary without looking at the notes.
- Compare your summary with the key points at the end of these notes
15. Worked Examples
Example 1: Find particles in an atom. An atom has atomic number 15 and mass number 31.
Find protons, neutrons and electrons, and the electron arrangement.
- Protons = Z = 15. Electrons = 15 (neutral atom).
- Neutrons = A − Z = 31 − 15 = 16.
- Electron arrangement: K = 2, L = 8, M = 5, so 2, 8, 5. Valency = 8 − 5 = 3.
The element is phosphorus.
Example 2: Ions. An ion has 12 protons and 10 electrons. What is its charge and element?
Charge = 12 − 10 = +2. The element with Z = 12 is magnesium, so the ion is Mg²⁺.
Example 3: Identify isotopes. Species A has 17 protons and 18 neutrons. Species B has
17 protons and 20 neutrons. Species C has 18 protons and 22 neutrons.
- A and B have the same Z (17) but different A (35 and 37), so they are isotopes
(chlorine-35 and chlorine-37).
- C has Z = 18 and A = 40, so it is a different element (argon-40). It is an isobar
of calcium-40.
- C has 22 neutrons, so it is an isotone of any atom with 22 neutrons,
such as scandium-43 (21 protons, 22 neutrons).
Example 4: Average atomic mass. Copper has ⁶³Cu (69.2 percent) and ⁶⁵Cu (30.8 percent). Average = (63 × 0.692) + (65 × 0.308) =
43.596 + 20.02 = 63.62 u.
Example 5: Electron configuration. Write the configuration of sulphur (Z = 16). It is 1s² 2s² 2p⁶ 3s² 3p⁴. That is 2 + 2 + 6
+ 2 + 4 = 16 electrons. Valence electrons = 6 (3s² 3p⁴).
Example 6: Half-life. A radioactive sample of 160 g has a half-life of 5 years. How much
remains after 20 years? Number of half-lives = 20 ÷ 5 = 4. Remaining = 160 ÷ 2⁴ = 160 ÷ 16 = 10 g.
16. Common Misconceptions
Misconception
Correct idea
Electrons orbit the nucleus like planets.
This is a useful simple picture (Bohr model), but
electrons exist as probability clouds called orbitals.
Atoms are solid little balls.
Atoms are mostly empty space.
Isotopes are different elements.
Isotopes are the same element with different neutron
numbers.
Mass number and atomic mass are the same.
Mass number is a whole-number count of nucleons. Atomic
mass is a weighted average over isotopes.
Atoms can never be split.
They can be split in nuclear reactions, but the result is
no longer the same element.
An ion has a different number of protons.
Ions differ only in electrons. Changing protons changes
the element.
Neutrons carry a small charge.
Neutrons are neutral.
Electrons are in the nucleus.
Electrons are outside the nucleus.
17. Practice Questions
Section A: Short answer
- Define an atom and state why atoms are electrically neutral.
- Write any three postulates of Dalton's atomic theory and one that was later disproved.
- State the three observations of Rutherford's alpha-scattering experiment and what each proves.
- List two limitations of Rutherford's model.
- What are the postulates of Bohr's model concerning energy levels?
- Differentiate between an orbit and an orbital.
- What is meant by valency? Find the valency of sodium and oxygen.
- Define isotopes, isobars and isotones with one example of each.
- Why do isotopes of an element have the same chemical properties?
- State Pauli's exclusion principle and Hund's rule.
Section B: Numerical and application
Calculate the numbers of protons, neutrons and electrons in ⁴⁰Ca, ³¹P and ²⁷Al³⁺.The element X has the electron arrangement 2, 8, 6. Identify the element, its valency and
the type of ion it forms. Bromine has isotopes ⁷⁹Br (50.7 percent) and ⁸¹Br (49.3 percent). Calculate the average
atomic mass. How many atoms are present in 4 g of helium (atomic mass 4 u)? A sample has a half-life of 3 hours. What fraction remains after 12 hours? Write the full electron configuration of iron (Z = 26) and chlorine (Z = 17), and draw an
orbital box diagram for oxygen.
Section C: Diagram tasks
Draw and label Bohr diagrams for lithium, nitrogen and chlorine. Sketch the gold foil experiment, labelling the source, foil, screen and three paths of alpha
particles. Draw an energy level diagram showing absorption and emission of light for hydrogen.
Draw the three isotopes of hydrogen showing their nucleons.
Answer Key (selected)
- Q7. Sodium (2,8,1) has valency 1; oxygen (2,6) has valency 8 − 6 = 2.
- Q11. ⁴⁰Ca: 20 p, 20 n, 20 e. ³¹P: 15 p, 16 n, 15 e. ²⁷Al³⁺: 13 p, 14 n, 10 e.
- Q12. Atomic number = 16, so it is sulphur. Valency 2. It gains two electrons t form S²⁻.
- Q13. (79 × 0.507) + (81 × 0.493) = 40.053 + 39.933 = 79.99 u (about 80u).
- Q14. 4 g of He = 1 mol = 6.022 × 10²³ atoms (helium is monatomic).
- Q15. 12 ÷ 3 = 4 half-lives, so (½)⁴ = 1/16 remains.
- Q16. Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s².
Cl: 1s² 2s² 2p⁶ 3s² 3p⁵.
18. Quick Revision Summary
- The atom is the smallest unit of an element that retains its chemical identity.
- Dalton proposed solid atoms; Thomson found electrons; Rutherford found the nucleus;
Bohr introduced energy levels; quantum mechanics introduced orbitals and probability. - Atoms contain protons (+1), neutrons (0) and electrons (−1). Protons and neutrons sit
in the nucleus and account for nearly all the mass. - Atomic number (Z) = number of protons. Mass number (A) = protons + neutrons.
- Isotopes share Z; isobars share A; isotones share the neutron number.
- Shell capacity is 2n²; the outer shell holds at most 8 electrons.
- Valence electrons decide chemical behaviour and valency.
- Electrons fill orbitals by the Aufbau principle, Pauli exclusion principle and Hund's
rule.- Ions form by gaining or losing electrons, never protons.
- Light is emitted when electrons fall from higher to lower energy levels, producing line
spectra. - Radioactive decay produces alpha, beta and gamma radiation. Half-life describes how
fast a sample decays. - Mole = 6.022 × 10²³ particles.
19. Key Terms Glossary
Term
Meaning
Atom
Smallest unit of an element with that element's chemical
properties
Nucleus
Dense central core containing protons and neutrons
Nucleon
A proton or a neutron
Electron
Negatively charged particle outside the nucleus
Atomic number
Number of protons
Mass number
Number of protons plus neutrons
Isotope
Atoms of one element with different numbers of neutrons
Shell
Energy level in which electrons are located
Orbital
Region of space where an electron is most likely found
Valence electron
Electron in the outermost shell
Valency
Combining capacity of an atom
Ion
Charged atom or group of atoms
Cation / Anion
Positive ion / negative ion
Ground state
Lowest energy arrangement of an atom
Excited state
Higher energy state after absorbing energy
Half-life
Time for half of a radioactive sample to decay
Mole
6.022 × 10²³ particles
Study tip: Redraw the timeline of atomic models, the Rutherford diagram and the first 20 electron arrangements from memory every day for one week. Then test yourself with the practice questions without looking at the answers.