Atomic Structure — Science Sage

History of Atomic Models

Our understanding of the atom has evolved dramatically over two centuries as new experimental evidence emerged:

Scientist	Year	Model	Key Evidence / Experiment
John Dalton	1808	Solid sphere ("billiard ball")	Laws of definite and multiple proportions; atoms are indivisible, identical within an element
J.J. Thomson	1904	Plum pudding model	Cathode ray experiment; discovered electrons (negatively charged particles embedded in a positive sphere)
Ernest Rutherford	1911	Nuclear model	Gold foil experiment: most α-particles pass through; some deflect greatly → small, dense, positive nucleus; mostly empty space
Niels Bohr	1913	Planetary model	Hydrogen line spectrum; electrons orbit in fixed energy levels; emit/absorb photons when jumping between levels
Quantum model	1920s+	Electron cloud / orbital model	Wave-particle duality (de Broglie); Heisenberg uncertainty principle; electrons in probability clouds (orbitals)

Rutherford's Gold Foil Experiment: A beam of α-particles was fired at a thin gold foil. Most passed straight through (atom is mostly empty space), but a small fraction bounced back. This disproved Thomson's model and proved the existence of a small, dense, positively charged nucleus.

Structure of the Atom

Particle	Charge	Relative Mass	Location
Proton	+1	1 u	Nucleus
Neutron	0	1 u	Nucleus
Electron	−1	≈ 1/1836 u (negligible)	Shells around nucleus

The nucleus contains protons and neutrons (collectively called nucleons). It is extremely small and dense — if an atom were the size of a sports stadium, the nucleus would be a marble at the centre. Electrons occupy the vast empty space around the nucleus.

Atomic Number, Mass Number & Nuclide Notation

Atomic number (Z): number of protons in the nucleus; defines the element; equals the number of electrons in a neutral atom.
Mass number (A): total number of protons + neutrons (nucleons).
Number of neutrons: N = A − Z

ᴬ_Z X e.g. ¹²₆C means Carbon with A=12, Z=6, N=6

For a neutral atom: number of electrons = number of protons = Z

Isotopes

Isotopes are atoms of the same element (same Z, same number of protons) that have different numbers of neutrons (different A).

Isotopes have identical chemical properties (same electron configuration, same valence electrons).
Isotopes have different physical properties: different mass, different density, some are radioactive.

Isotope	Protons	Neutrons	Name
¹H	1	0	Protium (common hydrogen)
²H	1	1	Deuterium ("heavy hydrogen")
³H	1	2	Tritium (radioactive)
¹²C	6	6	Carbon-12 (standard)
¹⁴C	6	8	Carbon-14 (radioactive — used in dating)

Relative atomic mass (Ar): the weighted average mass of all naturally occurring isotopes of an element, relative to ¹/₁₂ the mass of ¹²C:

Ar = Σ (isotope mass × fractional abundance)

Electron Configuration

Electrons occupy shells (energy levels) around the nucleus. The maximum number of electrons per shell is 2n² (n = shell number), but for elements with Z ≤ 20, shell 3 holds a maximum of 8 electrons before shell 4 starts filling.

Shell (n)	Max electrons (2n²)	In practice for Z≤20
1	2	2
2	8	8
3	18	8 (then shell 4 starts)
4	32	2 (for Ca, Z=20)

Electron configuration notation lists electrons per shell separated by dots or dashes. Examples:

H (Z=1): 1 He (Z=2): 2
Li (Z=3): 2,1 Na (Z=11): 2,8,1
Mg (Z=12): 2,8,2 Cl (Z=17): 2,8,7
Ca (Z=20): 2,8,8,2

Valence electrons are the electrons in the outermost shell. They determine how an atom bonds and reacts. Elements in the same group of the periodic table have the same number of valence electrons.

Atomic Spectra & Flame Tests

When electrons absorb energy, they jump to higher energy levels (excited state). When they fall back to lower levels, they emit photons of specific wavelengths — producing a line emission spectrum unique to each element.

The hydrogen emission spectrum (Balmer series, visible light) shows distinct lines:

Red line ≈ 656 nm (n=3→2)
Blue-green line ≈ 486 nm (n=4→2)
Violet lines ≈ 434, 410 nm (n=5→2, n=6→2)

Flame tests use characteristic emission colours to identify metal ions:

Metal ion	Flame colour
Lithium (Li⁺)	Crimson / carmine red
Sodium (Na⁺)	Bright yellow / orange-yellow
Potassium (K⁺)	Lilac / violet
Calcium (Ca²⁺)	Brick red / orange-red
Copper (Cu²⁺)	Blue-green / verdigris

IEB Extension — Subshells & Orbital Notation

Within each shell, electrons occupy subshells (s, p, d, f) and within subshells, they occupy orbitals (each holding max 2 electrons):

s subshell: 1 orbital, max 2 electrons
p subshell: 3 orbitals, max 6 electrons
d subshell: 5 orbitals, max 10 electrons

H: 1s¹ He: 1s² Li: 1s² 2s¹ C: 1s² 2s² 2p² Ne: 1s² 2s² 2p⁶ Na: 1s² 2s² 2p⁶ 3s¹

Hund's Rule: Electrons fill each orbital in a subshell singly before any orbital is doubly occupied (electrons prefer parallel spins to minimise repulsion).

Pauli Exclusion Principle: No two electrons in the same atom can have the same four quantum numbers; each orbital holds at most 2 electrons with opposite spins (↑↓).

Atomic Model Explorer

Mode

Select Model

Question 1 — Atomic Models

Rutherford's gold foil experiment led him to conclude that the atom is mostly empty space with a small, dense, positive nucleus. Which observation most directly supported this conclusion?

Correct! Most particles passing through → mostly empty space. The tiny fraction bouncing straight back → collision with a small, very dense, positively charged nucleus. This disproved Thomson's "plum pudding" model.

The key observation was that MOST α-particles passed straight through (atom is mostly empty space), but a very small number bounced back at large angles (small, dense nucleus). If the positive charge were spread out (Thomson's model), none would deflect sharply.

Question 2 — Proton, Neutron & Electron Counts

An atom is represented as ³⁵₁₇Cl. How many protons, neutrons, and electrons does a neutral atom of this element have?

Correct! Z=17 → protons = 17. A=35 → neutrons = A−Z = 35−17 = 18. Neutral atom: electrons = protons = 17.

In ³⁵₁₇Cl: Z (bottom) = 17 = protons = electrons (neutral). A (top) = 35 = protons + neutrons → neutrons = 35−17 = 18.

Question 3 — Isotopes

Chlorine has two main isotopes: ³⁵Cl (75% abundant) and ³⁷Cl (25% abundant). What is the relative atomic mass of chlorine?

Correct! Ar = (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.5 u. This is why the periodic table shows Cl as 35.5.

Ar = (35 × 0.75) + (37 × 0.25) = 26.25 + 9.25 = 35.5 u. The relative atomic mass is a weighted average of all isotopes by their natural abundance.

Question 4 — Electron Configuration

What is the correct electron configuration for phosphorus (P, Z = 15)?

Correct! P has 15 electrons. Fill shell 1: 2, shell 2: 8, shell 3: 5. Total: 2+8+5 = 15. ✓ Phosphorus has 5 valence electrons (Group 15).

Fill shells in order: shell 1 = 2 electrons (full), shell 2 = 8 electrons (full), remaining = 15−10 = 5 in shell 3. Configuration: 2, 8, 5.

Question 5 — Valence Electrons

Sodium (Na, Z=11) and potassium (K, Z=19) are both in Group 1. What do they have in common in terms of electron configuration, and why does this explain their similar chemistry?

Correct! Na = 2,8,1 and K = 2,8,8,1 — both have exactly 1 valence electron. Chemical reactivity is determined by the outermost shell. Both readily lose this 1 electron to form +1 ions, explaining similar reactivity.

Na (2,8,1) and K (2,8,8,1) both have 1 valence electron. Chemistry is determined by the outermost electrons. Both lose 1 electron easily → both form +1 ions and react vigorously with water.

Question 6 — Atomic Spectra

A student holds a piece of wire with an unknown salt in a gas flame and observes a bright yellow-orange colour. Which metal ion is most likely present?

Correct! Sodium produces a characteristic bright yellow-orange flame. This is caused by electrons in Na atoms being excited and emitting photons at ≈589 nm (yellow) as they return to ground state.

Flame colours: Na⁺ = yellow, K⁺ = lilac, Li⁺ = crimson, Cu²⁺ = blue-green, Ca²⁺ = brick red. The bright yellow-orange is the signature of sodium.

Question 7 — Isotopic Abundance from Ar (Analysis)

Copper occurs naturally as two isotopes: ⁶³Cu (isotopic mass 62.93 u) and ⁶⁵Cu (isotopic mass 64.93 u). The periodic table gives copper's relative atomic mass as 63.55 u. What percentage of naturally occurring copper atoms are ⁶³Cu?

Correct! Let x = fraction of ⁶³Cu, so (1−x) = fraction of ⁶⁵Cu. Ar = 62.93x + 64.93(1−x) = 63.55. Expanding: 64.93 − 2x = 63.55, so 2x = 1.38, x = 0.69. So 69% of copper atoms are ⁶³Cu (and 31% are ⁶⁵Cu) — this matches copper's real isotopic composition.

Set up a weighted-average equation with x = fraction of ⁶³Cu: 62.93x + 64.93(1−x) = 63.55. This simplifies to 64.93 − 2x = 63.55, so x = 1.38/2 = 0.69, i.e. 69% ⁶³Cu and 31% ⁶⁵Cu. Unlike a simple Ar calculation, here you must rearrange to find an unknown abundance rather than compute Ar directly.

IEB Extension Practice IEB ONLY

IEB Question 1 — Orbital Notation

Which of the following is the correct full orbital (subshell) notation for sulfur (S, Z = 16)?

Correct! S has 16 electrons: 1s²(2) + 2s²(2) + 2p⁶(6) + 3s²(2) + 3p⁴(4) = 16. The 3p subshell has 4 electrons across 3 orbitals: ↑↓ ↑ ↑ (Hund's rule applies).

Fill subshells in order: 1s²(2) 2s²(2) 2p⁶(6) 3s²(2) = 12 electrons so far. S has 16, so 4 more go into 3p → 3p⁴. Full notation: 1s² 2s² 2p⁶ 3s² 3p⁴.

IEB Question 2 — Hund's Rule

Carbon (Z = 6) has the configuration 1s² 2s² 2p². According to Hund's Rule, how are the two 2p electrons distributed across the three 2p orbitals?

Correct! Hund's Rule: electrons occupy orbitals of equal energy singly and with parallel spins before pairing up. For carbon's 2p²: [↑][↑][ ] — one electron in each of two orbitals. This minimises electron-electron repulsion.

Hund's Rule states electrons fill degenerate (equal energy) orbitals one at a time with parallel spins before pairing. So carbon's 2p² = [↑][↑][ ], not [↑↓][ ][ ]. Pairing electrons in the same orbital costs extra energy due to repulsion.

Show all working for calculations. Draw large, neat, labelled diagrams for Bohr models. Use nuclide notation (ᴬ_Z X) where required.

Question 1 — Nuclide Notation Calculations

For each nuclide below, determine the number of protons, neutrons, and electrons (assume neutral atoms). Also write the full name of the element.

(a) ²³₁₁Na (b) ⁵⁶₂₆Fe (c) ¹⁹₉F (d) ²⁷₁₃Al (e) ²⁰₈Pb

Question 2 — Bohr Diagrams

Draw complete, labelled Bohr diagrams for each of the following elements. Your diagrams must show the nucleus (with proton and neutron count), and all electron shells with the correct number of electrons in each shell.

(a) Oxygen (O, Z=8, A=16) (b) Magnesium (Mg, Z=12, A=24) (c) Chlorine (Cl, Z=17, A=35) (d) Argon (Ar, Z=18, A=40)

For each element, also state: (i) the number of valence electrons, and (ii) the group number in the periodic table.

Question 3 — Isotope Relative Atomic Mass Calculation

Boron (B, Z=5) occurs naturally as two isotopes: ¹⁰B (abundance 19.9%) and ¹¹B (abundance 80.1%).
(a) State the number of protons, neutrons, and electrons in each isotope (assume neutral atoms).
(b) Calculate the relative atomic mass of boron. Show all working.
(c) Explain why the two isotopes have identical chemical properties but different physical properties.
(d) The periodic table lists boron's atomic mass as 10.81. Does your calculated value agree? Account for any small difference.

Question 4 — Flame Test Identification

A laboratory technician tests three unknown solutions (A, B, and C) using a flame test. The results are:
• Solution A: lilac/violet flame
• Solution B: brick red flame
• Solution C: blue-green flame

(a) Identify the metal ion most likely present in each solution.
(b) Explain, using the concept of electron energy levels, why different metals produce different coloured flames.
(c) Why is the sodium test unreliable if the platinum wire is contaminated with sodium?

Question 5 — History of Atomic Models Essay

Write a structured account (approximately 15–20 lines) tracing the development of the atomic model from Dalton (1808) to the modern quantum model. For each model:
• Name the scientist and date
• Describe the model briefly
• State the experimental evidence that supported OR disproved it
• Explain what limitation led to the next model

Conclude by explaining why scientific models change over time.

Question 6 — Reading a Mass Spectrum Data Table

A mass spectrometer analyses a sample of pure magnesium and records the following data for its three naturally occurring isotopes:

Isotope	Isotopic mass (u)	Relative abundance (%)
²⁴Mg	23.99	78.99
²⁵Mg	24.99	10.00
²⁶Mg	25.98	11.01

(a) State the number of protons and neutrons in an atom of each isotope.
(b) Use the data in the table to calculate the relative atomic mass (Ar) of magnesium. Show your working to 2 decimal places.
(c) Compare your answer to (b) with the value listed on the periodic table (24.31). Comment on the agreement.
(d) Explain why ²⁴Mg contributes far more to the final Ar value than ²⁶Mg does, even though their isotopic masses are similar.