Journey Inside the Atom | CBSE Class 9 Science Notes

This note covers early atomic ideas, the discovery of subatomic particles, Thomson’s, Rutherford’s and Bohr’s models, symbols of elements, atomic number, mass number, electron distribution, valency, isotopes, average atomic mass, isobars and applications of isotopes.

How did the idea of the atom develop?

From philosophical ideas to experiments

Atoms are the tiny building blocks of matter. Living beings and non-living objects are ultimately composed of atoms, which cannot be seen with the naked eye. Understanding their internal structure required scientists to ask whether these particles could be divided further.

More than 2,000 years ago, Acharya Kanada in India suggested that repeatedly dividing matter, called dravya, would eventually lead to particles that could no longer be divided. He called these particles parmanus. His ideas are recorded in the Sanskrit text Vaisesika Sutras.

In this description, a parmanu is infinitely small and cannot be perceived by the senses. Parmanus combine into dyads, groups of two, and triads, groups of three, and further combinations. These combinations form the material universe, including living bodies.

The description did not specify the proportions in which parmanus combine to make different substances. In ancient Greece, Leucippus and Democritus proposed a similar idea, calling indivisible particles atomos, a Greek word meaning indivisible.

What changed with Dalton?

The early concept of the atom began as an imaginary idea rather than an experimental observation. In 1808, John Dalton proposed an atomic theory based on scientific experiments. He described matter as consisting of indivisible atoms that could not be broken into smaller parts.

Dalton’s theory provided the first scientific description of how matter is made. Later discoveries required changes to the idea of indivisibility. An element is a substance identified by the proton count of its atoms. An atomic model represents the proposed structure of an atom; models were revised as experiments supplied new evidence.

The sequence of models shows how scientific explanations develop. Early models were not fully correct, but each helped frame new questions about what atoms contain, how their parts are arranged and why atoms of different elements differ.

How did electrons lead to Thomson’s model?

What did the cathode ray experiment show?

Radioactivity is the emission of invisible energy and particles, called radiation, by certain elements. Such discoveries challenged the idea of indivisible atoms. In 1897, J. J. Thomson investigated the conduction of electric current through gases at very low pressure.

He applied a high voltage across two electrodes, the electrical terminals inside a glass tube. Rays travelled from the cathode, the negative electrode, towards the anode, the positive electrode. These were called cathode rays.

Studying the rays in electric and magnetic fields showed that they were streams of negatively charged particles much lighter than atoms. These particles were later called electrons. Their emission from atoms showed that atoms contain smaller, or subatomic, components.

The nature of cathode rays did not depend on the cathode material or the gas in the tube. This showed that electrons are components of all atoms. The electron’s relative charge is conventionally taken as −1, meaning one unit of negative charge.

A glass tube contains a cathode marked negative and an anode marked positive. The diagram labels a high voltage power supply and a vacuum pump, with the ray path running through the tube.

See Fig. 8.1 in your NCERT textbook

How did Thomson explain neutrality?

A neutral atom has balanced positive and negative charges. Thomson proposed a sphere of positive charge with electrons distributed throughout it. This became the plum pudding model, comparing electrons to plums embedded in a pudding.

A watermelon provides another comparison: its red pulp represents the positively charged matter and its seeds represent electrons. The model explained balanced charges, but its spread-out positive charge later failed to explain the gold foil experiment.

Small blue dots are distributed within a red circular representation of a positive sphere. Labels identify an electron and the positive sphere.

See Fig. 8.2 in your NCERT textbook

Note: Atoms do not show any colour. Colours in atomic diagrams are for illustration; the watermelon comparison describes the arrangement in a model.

What did the gold foil experiment reveal?

What was expected and what was observed?

In 1911, Geiger and Marsden, working under Ernest Rutherford, directed a narrow beam of alpha particles at extremely thin gold foil. The symbol α means alpha. These positively charged particles are emitted by certain radioactive elements.

An alpha particle is a helium nucleus containing two positively charged protons and two uncharged neutrons. A nucleus is the small central region of an atom. Thomson’s evenly spread positive charge suggested that the particles should pass straight through or be deflected only slightly.

Most particles passed through undeflected, some were sharply deflected, and a few bounced back. Scattering means deflection from a straight path, so this investigation is also called the alpha-ray scattering experiment. Thomson’s model could not explain the observed pattern.

Parallel arrows labelled as a beam of alpha particles approach the thin gold foil. Many paths continue straight, while other paths turn through different angles, including a path directed backwards.

See Fig. 8.4 in your NCERT textbook

What did Rutherford propose?

Rutherford’s nuclear model placed all the positive charge and most of the mass in a dense nucleus. Electrons revolve around it, somewhat like planets around the Sun. This comparison gives the name planetary model.

The approximate atomic diameter is 10⁻¹⁰ metres, while the nuclear diameter is about 10⁻¹⁵ metres. The nucleus is therefore about 10⁵, or one lakh, times smaller than the atom. Here a diameter is the distance across the object through its centre.

Worked example 1. A sheet of paper is 0.1 millimetres thick, equivalent to 10⁻⁴ metres. Estimate the number of atoms across its thickness if each atom has a diameter of about 10⁻¹⁰ metres.

Answer: The sign ≈ means approximately equal to. Number of atoms ≈ thickness ÷ atomic diameter = 10⁻⁴ ÷ 10⁻¹⁰ = 10⁶. About one million atoms would need to be stacked across the thickness.

How did Bohr explain atomic stability?

Why did Rutherford’s model need revision?

Rutherford’s model explained the scattering results but could not explain atomic stability, the continued existence of atoms without collapse. A particle moving in a circle continually changes direction and therefore accelerates. In the earlier explanation, an accelerating electron should lose energy.

As its energy decreased, the electron should spiral towards the positively charged nucleus and eventually fall into it. Atoms would collapse if this happened. Since matter remains intact, a model was needed that could explain why moving electrons do not undergo this collapse.

What are stationary states?

In 1913, Niels Bohr proposed fixed circular paths called stationary states, orbits or shells. Each shell gives an electron a definite energy, so shells are also called energy levels.

1. Electrons can revolve in allowed shells, but not in the spaces between them.
2. While an electron moves in a fixed shell, it does not lose energy.
3. The innermost shell has the least energy; shells farther from the nucleus have higher energy.
4. An electron can change shells by absorbing or releasing a fixed amount of energy equal to the difference between the two levels.

The shells are labelled K, L, M and N, corresponding to shell numbers n = 1, 2, 3 and 4. Here n denotes the shell number counted outwards from the nucleus. Each shell can hold only a certain number of electrons.

Four concentric circles surround a central nucleus. From inside to outside, they are labelled K-shell (n = 1), L-shell (n = 2), M-shell (n = 3) and N-shell (n = 4).

See Fig. 8.7 in your NCERT textbook

Bohr introduced stationary states as a postulate, an assumption of the model. In such a state, an electron’s energy remains constant even while it moves. This explained stability and many experimental observations, although Bohr’s model also later proved to have limitations.

How do electrons, protons and neutrons compare?

Which particles contribute charge and mass?

Protons provide the positive charge of the nucleus. Rutherford discovered and named them. A proton is much heavier than an electron, but their charges are equal in magnitude and opposite in sign. Equal numbers of protons and electrons therefore balance each other’s charges.

For example, a neutral helium atom has two protons and two electrons; a neutral sodium atom has eleven of each. Changing the neutron count does not directly add electrical charge because a neutron is uncharged.

In 1932, James Chadwick discovered the neutron, whose mass is nearly equal to a proton’s mass. The discovery explained additional nuclear mass. Helium has two protons, yet its atomic mass is about four times that of hydrogen, rather than double.

In the particle symbols, e⁻ means electron, p⁺ means proton and n⁰ means neutron. The superscripts indicate negative, positive and zero charge respectively. The particle symbol n⁰ should be distinguished from the shell number n.

Why is electron mass usually ignored here?

Most atomic mass comes from protons and neutrons packed in the nucleus. Electrons are so light in comparison that their mass can be ignored in these calculations. This means negligible mass for the calculation, rather than no mass.

Lighter atoms often have equal numbers of protons and neutrons, such as carbon with six of each and oxygen with eight of each. Heavier nuclei have many more neutrons: the examples of iron and uranium contain 26 and 92 protons, with 30 and 146 neutrons respectively.

How are chemical symbols written and understood?

Why use a common system?

A chemical symbol is the recognised representation of an element. John Dalton introduced pictorial symbols in 1803. In 1813, Berzelius suggested deriving symbols from Latin names, helping establish alphabetic chemical symbols.

The International Union of Pure and Applied Chemistry, abbreviated IUPAC, approves element names and symbols. Shared symbols allow scientists to communicate across languages without repeatedly writing full element names or inventing separate systems.

Many symbols use the first letter or the first two letters of an element’s name. The first letter is always a capital; a second letter, if present, is lower case. Thus hydrogen is H, aluminium is Al and cobalt is Co.

Some symbols use the first letter and another letter that is not second in the name. Chlorine is Cl and zinc is Zn. The letter case forms part of the symbol: AL is not the correct symbol for aluminium, and CO is not the correct symbol for cobalt.

Why do some symbols differ from English names?

Other symbols come from Latin, Greek or German names. Iron is Fe from Latin ferrum, mercury is Hg from Greek hydrargyros, and tungsten is W from German wolfram. Learning these origins explains why the letters may differ from the English name.

How do atomic number and mass number describe an atom?

What do Z and A count?

Definition: The atomic number, represented by Z, is the number of protons in an atom’s nucleus. It identifies the element. The mass number, represented by A, is the total number of protons and neutrons in the nucleus.

Protons and neutrons together are called nucleons. For the following counting formulas, P denotes the number of protons, N the number of neutrons and E the number of electrons. These capital letters represent counts, not chemical symbols in these formulas.

Z = P

A = P + N

N = A − Z

E = Z for a neutral atom.

Hydrogen has atomic number 1 and helium has atomic number 2. Their neutral atoms therefore have one and two electrons respectively. The proton count determines identity; the neutron count contributes to the mass number.

How is the notation read?

In standard atomic notation, the mass number is written to the upper left of the element symbol and the atomic number to its lower left. For example, ¹²₆C denotes carbon with mass number 12 and atomic number 6.

Read the two numbers separately before calculating. The upper number counts all nucleons, while the lower number counts protons. Subtracting the lower number from the upper gives the neutron count. Electron numbers follow from neutrality, not by adding electrons to the mass number.

Worked example 2. A neutral atom has atomic number 26 and 56 nucleons. Find its proton, electron and neutron counts.

Answer: P = Z = 26 and E = Z = 26. Since 56 nucleons means A = 56, N = A − Z = 56 − 26 = 30. The atom has 26 protons, 26 electrons and 30 neutrons.

Worked example 3. A nucleus contains 20 protons and has mass number 41. Find its neutron count.

Answer: N = A − P = 41 − 20 = 21 neutrons. The mass number includes both kinds of nucleons, so the proton count must be subtracted.

How are electrons distributed among shells?

What are the Bohr and Bury rules?

Electronic configuration is the distribution of electrons among an atom’s shells. Bohr and Bury described rules for this distribution. Shells fill stepwise, beginning with the shell closest to the nucleus and moving outwards through K, L, M and N.

Let Cₙ mean the maximum electron capacity of shell number n. Then Cₙ = 2n², where n² means n multiplied by itself. The K-shell can hold 2 × 1² = 2 electrons, the L-shell 2 × 2² = 8 and the M-shell 2 × 3² = 18.

The outermost shell can accommodate at most eight electrons, with the first shell limited to two. These two statements must be read together: a shell’s maximum capacity does not mean that it already contains that many electrons in every atom.

Note: The M-shell capacity is 18, but the first eighteen elements have at most eight electrons in that shell. Argon’s electronic configuration is 2, 8, 8.

What pattern appears in the first eighteen elements?

Hydrogen has one electron in K; helium has two there. Lithium begins the L-shell after K is complete. Sodium begins the M-shell after the arrangement 2, 8. Read each distribution from K outwards and check that its sum equals the neutral atom’s electron count.

The following table brings together particle counts and shell distributions. Read each row as an example of an atom with the stated neutron count. The K, L, M and N columns describe shells; a hyphen marks an entry without a numerical value.

Schematic atoms from hydrogen to argon show electrons as dots on circular shells around central dots. The diagrams display the successive filling of K, L and M shells.

See Fig. 8.11 in your NCERT textbook

Worked example 4. Write the electronic configuration of a neutral magnesium atom with atomic number 12.

Answer: It has 12 electrons. Put 2 in K, then 8 in L. The remaining 12 − 2 − 8 = 2 occupy M, giving the configuration 2, 8, 2.

How does electronic configuration determine valency?

What is the combining capacity of an atom?

The combining capacity of an atom can be expressed as the number of hydrogen or chlorine atoms with which it combines to form a compound. Hydrogen and chlorine each have a combining capacity of one.

In water, written H₂O, one oxygen atom combines with two hydrogen atoms, giving oxygen a combining capacity of two. In a chemical formula, a subscript indicates the number of atoms of the element immediately before it; an omitted subscript means one.

NH₃, ammonia, contains one nitrogen atom with three hydrogen atoms, indicating nitrogen’s combining capacity of three. MgCl₂, magnesium chloride, has one magnesium atom with two chlorine atoms, indicating magnesium’s combining capacity of two.

How do valence electrons help?

The valence shell is the outermost shell containing electrons, and its electrons are valence electrons. An octet is an arrangement of eight electrons in the outermost shell. Elements with complete octets, or two valence electrons in helium, are largely unreactive and more stable.

Atoms with incomplete valence shells are usually more reactive. They lose, gain or share electrons to complete the octet. Valency is the number of electrons gained, lost or shared to achieve this stable arrangement.

Generally, an element with fewer than four valence electrons tends to lose electrons; one with more than four tends to gain them. Carbon has four valence electrons and cannot easily gain or lose them, so it can share four electrons.

Valency and valence-electron count are different quantities. Oxygen has six valence electrons but valency two. Some compounds appear to violate the usual valency rule, so the rule should not be treated as an exceptionless description of every compound.

What are isotopes and why are their chemical properties similar?

How can atoms of one element have different masses?

Definition: Isotopes are atoms of the same element with the same atomic number but different mass numbers. They have equal proton counts but different neutron counts.

The discovery of isotopes changed Dalton’s proposal that all atoms of an element have the same mass. The element’s identity remains the same because the proton count is unchanged. Adding different numbers of neutrons changes the mass number.

Hydrogen has three naturally occurring isotopes: protium, ¹₁H, at approximately 99.98%; deuterium, ²₁H, at approximately 0.015%; and tritium, ³₁H, in traces. Each contains one proton. Their neutron counts are zero, one and two respectively.

A trace means a very small amount. Each neutral hydrogen isotope has one electron because its single proton must be balanced by one negative electron. Different neutron counts do not require different electron counts.

Three circular atomic diagrams are labelled protium, deuterium and tritium. Each has one electron on its shell, while the central clusters contain one, two and three particles respectively.

See Fig. 8.12 in your NCERT textbook

Which properties remain alike?

Isotopes have similar chemical properties because their neutral atoms have the same electron count and electronic configuration. Chemical properties depend mainly on valence electrons. Physical properties, such as melting and boiling points, differ between isotopes.

Carbon has the isotopes ¹²₆C, ¹³₆C and ¹⁴₆C. Each has six protons and six electrons, but their neutron counts differ. Carbon-12 is the most abundant carbon isotope in nature. In a name such as carbon-12, the number after the hyphen is the mass number.

Worked example 5. Two neutral atoms each have 11 protons. One has 12 neutrons and the other 13. Compare their identities and mass numbers.

Answer: Both have atomic number 11 and belong to the same element. Their mass numbers are 11 + 12 = 23 and 11 + 13 = 24. They are isotopes.

How is average atomic mass calculated?

Why is a weighted average needed?

The unified atomic mass unit, symbol u, is a small unit used to express atomic masses. Kilograms and grams are inconveniently large for atoms. Relative abundance describes the proportion of an isotope present in a naturally occurring mixture.

Chlorine occurs in two isotopic forms with masses 35 u and 37 u in the ratio 3:1. Chlorine-35 constitutes about 75% and chlorine-37 about 25%. Here % means per hundred, so 75% corresponds to the fraction 75/100.

A simple arithmetic mean adds the two isotope masses and divides by two. That gives $\frac{35 + 37}{2}\,\mathrm{u} = 36\,\mathrm{u}$, but ignores their unequal abundances. A weighted average gives each isotope a contribution proportional to how common it is.

1. Write down each isotope’s mass and its percentage abundance.
2. Divide each percentage abundance by 100 to obtain a fraction.
3. Multiply each isotope’s mass by its abundance fraction.
4. Add the contributions to obtain the weighted average atomic mass.

Worked example 6. Calculate chlorine’s weighted average atomic mass using isotope masses 35 u and 37 u and abundances 75% and 25% respectively.

Answer: Weighted average $= (35 \times \frac{75}{100} + 37 \times \frac{25}{100})\,\mathrm{u} = (\frac{105}{4} + \frac{37}{4})\,\mathrm{u} = \frac{142}{4}\,\mathrm{u} = 35.5\,\mathrm{u}$.

Does each atom have the average mass?

The answer 35.5 u does not mean that an individual chlorine atom has that fractional mass. It represents the mixture. For example, one million chlorine atoms would contain 7.5 lakh chlorine-35 atoms and 2.5 lakh chlorine-37 atoms in this ratio.

The distinction is between the mass of an individual isotope and the average across many atoms. The more abundant isotope contributes more to the weighted result. The simple average fails here because it treats the two isotope masses as equally represented.

How do isobars differ from isotopes, and where are isotopes used?

What stays equal in isobars?

Isobars are atoms of different elements with the same mass number but different atomic numbers. Calcium, potassium and argon provide examples with mass number 40 and atomic numbers 20, 19 and 18 respectively.

The total number of nucleons is equal in these atoms, but the number of protons differs. They therefore remain different elements. For isotopes, the unchanged quantity is atomic number; for isobars, it is mass number.

What are some applications of isotopes?

Some isotopes have special properties useful in different fields. Uranium-235 is used as fuel in nuclear reactors to generate electricity. Cobalt-60 is a radioactive isotope used in radiation treatment for cancer.

Iodine-131 is used to treat goitre, an enlargement of the thyroid gland, and thyroid cancer. Carbon-14 helps determine the ages of ancient fossils and artefacts in archaeology and geology. These are applications of particular isotopes.

Is the exploration of atomic structure complete?

Bohr’s model was an important step, but electrons do not follow well-defined paths like its fixed orbits. The modern description uses electron clouds, regions around the nucleus where electrons are most likely to be found.

These regions allow predictions about where an electron is most likely to be, rather than its exact position. Atomic models therefore continue the same pattern of development: evidence improves explanations, and an earlier useful model can be replaced by a more complete description.

Glossary

• Atom — A tiny building block of matter containing smaller particles arranged within its structure.
• Electron — A negatively charged subatomic particle with a much smaller mass than a proton.
• Proton — A positively charged subatomic particle located in the nucleus of an atom.
• Neutron — An uncharged nuclear particle with a mass nearly equal to that of a proton.
• Nucleus — The extremely small, dense central region containing all positive charge and most atomic mass.
• Stationary state — An allowed shell in Bohr’s model where a moving electron does not lose energy.
• Atomic number — The number of protons in an atom’s nucleus, identifying the element.
• Mass number — The total number of protons and neutrons present in an atom’s nucleus.
• Nucleons — The protons and neutrons counted together as particles in the nucleus.
• Electronic configuration — The distribution of an atom’s electrons among its different shells or energy levels.
• Valence electrons — The electrons present in the outermost occupied shell of an atom.
• Valency — The combining capacity expressed through electrons gained, lost or shared to achieve a stable configuration.
• Isotopes — Atoms of the same element with equal atomic numbers but different mass numbers.
• Isobars — Atoms of different elements with equal mass numbers but different atomic numbers.
• Weighted average atomic mass — The average obtained by accounting for both isotope masses and their relative abundances.

Common errors and misconceptions

• Misconception: Dalton’s indivisible atom remains a complete description. Correct: Discoveries of electrons, protons and neutrons showed that atoms contain smaller particles.
• Misconception: Most alpha particles bounced back from gold foil. Correct: Most passed through undeflected; some were sharply deflected, and a few bounced back.
• Misconception: Rutherford explained atomic stability completely. Correct: His model could not explain why electrons would avoid losing energy and collapsing into the nucleus.
• Misconception: Electrons have no mass, so they are ignored. Correct: Their mass is almost negligible compared with nuclear particles and can be ignored in these calculations.
• Misconception: Mass number is the total of protons, neutrons and electrons. Correct: It counts protons and neutrons, the particles called nucleons.
• Misconception: The M-shell must contain 18 electrons in argon. Correct: Its maximum capacity is 18, but argon’s configuration is 2, 8, 8.
• Misconception: Valency always equals the number of valence electrons. Correct: Oxygen has six valence electrons but valency two because it can gain two electrons to complete its octet.
• Misconception: Every chlorine atom has mass 35.5 u. Correct: This is the weighted average for its naturally occurring isotopic mixture.

Exam-style questions with model answers

Q1. State Thomson’s model of a neutral atom and explain how it balances electrical charge. [2 marks]

1. Thomson described an atom as a positively charged sphere containing negatively charged electrons distributed throughout it.
2. The total positive charge balances the total negative charge of the electrons, so the atom as a whole is electrically neutral.

Q2. In the gold foil experiment, positively charged alpha particles were directed at extremely thin gold foil. Most passed through undeflected, some were sharply deflected and a few bounced back. Explain three conclusions about atomic structure supported by these observations. [3 marks]

1. Most of the atom must be empty space because most alpha particles could cross the thin foil without being deflected.
2. The atom’s positive charge is concentrated in a small central region, rather than spread uniformly throughout the atom, as indicated by sharp deflections.
3. The centre is extremely small and dense and contains most atomic mass, accounting for the few particles that bounced back.

Q3. Explain the stability problem in Rutherford’s model and how Bohr’s postulates addressed it. Include allowed shells, energy loss and transitions between shells. [5 marks]

1. In Rutherford’s model, an electron moving in a circular path continually changes direction and therefore accelerates around the positively charged nucleus.
2. An accelerating charged electron should lose energy in this explanation, spiral inwards and eventually fall into the nucleus, causing the atom to collapse.
3. Bohr proposed fixed allowed shells called stationary states. Each has a definite energy, and electrons cannot revolve in the spaces between allowed shells.
4. He postulated that an electron does not lose energy while moving in a stationary state. Its energy remains constant, addressing the problem of collapse.
5. An electron can change shells by absorbing or releasing a fixed amount of energy equal to the difference between the two energy levels.

Q4. A neutral magnesium atom has atomic number 12 and mass number 24. The K and L shells hold a maximum of 2 and 8 electrons respectively. Find its numbers of protons, neutrons and electrons, and write its electronic configuration. [4 marks]

1. The proton count is 12 because the atomic number is defined as the number of protons in the nucleus.
2. The neutron count is mass number minus atomic number, so it is 24 − 12 = 12.
3. The electron count is 12 because a neutral atom has equal numbers of electrons and protons.
4. The electronic configuration is 2, 8, 2: two electrons occupy K, eight occupy L and the remaining two occupy M.

Q5. Sodium, oxygen and carbon have electronic configurations 2, 8, 1; 2, 6; and 2, 4 respectively. Explain the valency of each using loss, gain or sharing of electrons. [3 marks]

1. Sodium has one electron in its valence shell. It can lose this electron to leave a complete octet in the outermost occupied shell, giving valency one.
2. Oxygen has six valence electrons and can gain two more to complete an octet, so its valency is two.
3. Carbon has four valence electrons and cannot easily gain or lose them. It can share four electrons to complete its octet, giving valency four.

Q6. Chlorine occurs as isotopes of mass 35 u and 37 u with abundances 75% and 25% respectively, where u is the unified atomic mass unit. Calculate its weighted average atomic mass, explain why an unweighted mean is unsuitable, and interpret the result. [5 marks]

1. The abundance fractions are 75/100 and 25/100. A weighted average must use these unequal proportions because the two isotopes are not equally common.
2. The contribution from chlorine-35 is 35 × 75/100 = 105/4 u. Its larger abundance gives it greater weight in the calculation.
3. The contribution from chlorine-37 is 37 × 25/100 = 37/4 u. Add both contributions to obtain 142/4 = 35.5 u.
4. The unweighted mean, (35 + 37) ÷ 2 = 36 u, is unsuitable because it ignores the stated difference in isotope abundances.
5. The value 35.5 u describes the average mass across the isotopic mixture. It does not mean that an individual chlorine atom has that fractional mass.

Q7. Two neutral atoms each contain 11 protons, but have 12 and 13 neutrons respectively. Calculate their mass numbers and identify their relationship. [2 marks]

1. The mass numbers are 11 + 12 = 23 and 11 + 13 = 24, because mass number counts protons and neutrons.
2. They are isotopes of the same element: both have atomic number 11, but their neutron counts and mass numbers differ.

Q8. Carbon-12 and carbon-14 both have atomic number 6. Calcium-40, potassium-40 and argon-40 have atomic numbers 20, 19 and 18 respectively. Classify both sets as isotopes or isobars, explaining the role of atomic number and mass number. The number in each name is its mass number. [4 marks]

1. Carbon-12 and carbon-14 have the same atomic number, 6, so their nuclei contain equal numbers of protons and they belong to the same element.
2. Their mass numbers differ, so the carbon atoms are isotopes: atoms of one element with equal atomic numbers but unequal mass numbers.
3. Calcium-40, potassium-40 and argon-40 have different atomic numbers, so their proton counts differ and they are different elements.
4. All three have mass number 40, so they are isobars. Their equal total nucleon counts do not imply equal proton counts.

Key takeaways

• Atomic models developed through new experimental evidence, from indivisible atoms to descriptions containing subatomic particles and energy levels.
• Thomson placed electrons within a positive sphere; Rutherford concentrated positive charge and most atomic mass in a tiny nucleus.
• Most alpha particles crossed gold foil undeflected, supporting the conclusion that most atomic space is empty.
• Bohr proposed allowed shells where electrons retain fixed energy while moving, providing an explanation for atomic stability.
• Atomic number counts protons; mass number counts protons and neutrons; a neutral atom has equal proton and electron counts.
• Electronic configuration describes shell populations, while valency concerns electrons gained, lost or shared to achieve a stable configuration.
• Isotopes share atomic number but differ in mass number; isobars share mass number but have different atomic numbers.
• Average atomic mass depends on isotope masses and relative abundances, so a weighted average is required for unequal isotope proportions.

Test yourself

Why did cathode rays suggest that electrons occur in all atoms?

Their nature was independent of both the cathode material and the gas inside the tube.

What does the symbol n represent in the shell-capacity rule 2n²?

It represents the shell number counted from the nucleus: K is 1, L is 2 and M is 3.

A neutral atom has atomic number 17 and 18 neutrons. What are its mass number and electron count?

Its mass number is 17 + 18 = 35, and it has 17 electrons to balance its 17 protons.

How do valence electrons and valency differ for oxygen with configuration 2, 6?

Oxygen has six valence electrons but valency two because it can gain two electrons to complete its octet.

Why do neutral isotopes have similar chemical properties?

They have the same electron count and electronic configuration; chemical properties depend mainly on valence electrons.

Why does equal mass number not necessarily mean equal atomic number?

Mass number counts protons and neutrons together. Isobars have equal totals but different proton counts and therefore different atomic numbers.

What is the correct symbol for cobalt, Co or CO?

The correct symbol is Co: the first letter is a capital and the second is lower case.

Do modern atomic descriptions place electrons on exact Bohr orbits?

No. Electron clouds describe regions where electrons are most likely to be found, rather than exact fixed paths.