Study sheet: Fundamentals of Atomic Structure and Periodic Trends

📋 Course Outline

1. Atomic theory and models
2. Subatomic particles and isotopes
3. Electron configuration and shells
4. Ionization energy trends
5. Periodic table blocks
6. Configuration exceptions and noble gas notation
7. Ions, free radicals and periodic trends

📖 1. Atomic theory and models

🔑 Key Concepts & Definitions

• Atomos : Atoms are defined as indivisible particles proposed by Democritus for matter’s basic units.
• Dalton’s atomic theory : Dalton’s theory says elements have identical atoms, different elements have different atoms, and atoms combine in whole-number ratios.
• Thomson’s atomic model : Thomson’s model treats the atom as containing sub-atomic particles, implying the atom is divisible.
• Rutherford’s nuclear model : Rutherford’s model explains the atom as mostly empty space with a tiny dense positive nucleus.
• Bohr’s atomic model : Bohr’s model states electrons occupy fixed energy levels and can change levels by emitting or absorbing energy.

📝 Essential Points

• Thomson’s model included electrons and suggested the atom could be divided.
• Rutherford’s gold foil experiment led to the discovery of a dense positive nucleus.
• Bohr introduced fixed energy levels and quantum ideas to explain the hydrogen emission spectrum.
• Atomic stability in Bohr’s model comes from electrons moving in fixed energy levels without losing energy continuously.
• Each model’s limits include failure to explain spectral lines under Thomson and Rutherford, and incomplete quantum treatment for Bohr with multi-electron atoms.

💡 Memory Hook

Rutherford finds the “nucleus” by bouncing alpha particles; Bohr adds “steps” (quantised levels) to make spectra fit.

📖 2. Subatomic particles and isotopes

🔑 Key Concepts & Definitions

• Cathode ray tube : A cathode ray tube experiment uses electric fields to observe how charged beams deflect, revealing subatomic particle properties.
• Atomic number (Z) : Atomic number is the count of protons in an atom.
• Mass number (A) : Mass number is the total nucleons in an atom, equal to protons plus neutrons.
• Isotopes : Isotopes are atoms of the same element with different numbers of neutrons.
• Nucleon number : Nucleon number is the total number of protons and neutrons in a nucleus.

📝 Essential Points

• In the electric field deflection test, protons (+) deflect toward the negative plate and away from the positive plate.
• Electrons (−) deflect toward the positive plate and away from the negative plate, while neutrons are not deflected.
• Relative mass values given are proton ≈ neutron and proton is 1836 times heavier than an electron.
• Formulas given are atomic number Z = number of protons and mass number A = protons + neutrons.
• For neutrons: number of neutrons = A − Z.

💡 Memory Hook

Isotopes share Z, differ in A by changing neutrons (neutrons = A − Z).

📖 3. Electron configuration and shells

🔑 Key Concepts & Definitions

• Principal quantum number (n) : The principal quantum number n labels the principal shells and corresponds to the electron energy level.
• Shell : A shell is a principal energy level at a specific distance range from the nucleus where electrons can be found.
• Sub-shell : A sub-shell is a region within a principal shell with a specific energy level and type such as s, p, d, or f.
• Atomic orbital : An atomic orbital is a space region around the nucleus that can hold a maximum of two electrons.
• Hund’s Rule : Hund’s Rule states electrons occupy separate orbitals of the same sub-shell before pairing up.

📝 Essential Points

• Orbital capacity values given are: s sub-shell has 1 orbital, p has 3 orbitals, and d has 5 orbitals.
• Each orbital holds 2 electrons, and opposite spins are required for the second electron in the same orbital.
• Aufbau Principle is used to fill orbitals from lowest energy to higher energy in order.
• Pauli Exclusion Principle is that no two electrons in an atom share the same set of four quantum numbers (n, l, m, s).
• Orbital shapes given are s spherical, p dumbbell with three orientations, and d typically more complex or clover-like.

💡 Memory Hook

Pauli: 2 per box; Hund: spread first; Aufbau: fill lowest energy first.

📖 4. Ionization energy trends

🔑 Key Concepts & Definitions

• First ionisation energy (IE₁) : First ionisation energy is the energy needed to remove 1 mole of electrons from neutral gaseous atoms to form 1 mole of gaseous ions.
• Successive ionisation energies : Successive ionisation energies are the energies required to remove the next electron after previous electrons have already been removed.
• Principal quantum shell : A principal quantum shell is associated with electron energy level n and can change which ionisation step causes a large jump.
• Shielding effect : Shielding effect occurs when inner electrons reduce the attraction felt by outer electrons, affecting ionisation energy.

📝 Essential Points

• Given example for calcium: Ca(g) → Ca⁺(g) + 1e⁻ has IE₁ = 590 kJ/mol.
• Given example for calcium: Ca⁺(g) → Ca²⁺(g) + 1e⁻ has IE₂ = 1150 kJ/mol.
• Given example for calcium: Ca²⁺(g) → Ca³⁺(g) + 1e⁻ has IE₃ = 4940 kJ/mol.
• An expected general pattern is IE₁ < IE₂ < IE₃ because electrons removed later are held more tightly by the nucleus.
• A large jump between successive ionisation energies indicates the electron removed is from a principal quantum shell closer to the nucleus.

💡 Memory Hook

Big jump = next shell inward; small steps = same shell being stripped.

📖 5. Periodic table blocks

🔑 Key Concepts & Definitions

• s-block elements : s-block elements are atoms whose last electron enters the s-orbital.
• p-block elements : p-block elements are atoms whose last electron enters the p-orbital.
• d-block elements : d-block elements are atoms whose last electron enters the d-orbital.
• f-block elements : f-block elements are atoms whose last electron enters the f-orbital, found at the table’s bottom.

📝 Essential Points

• s-block elements include group 1 and group 2 elements, including H and He.
• p-block elements are stated to be elements of group 13 to 18.
• d-block elements are stated to be elements of group 3 to 12 and usually have partially filled d-orbitals.
• f-block elements are commonly known as inner transition or rare earth elements and lie at the bottom of the table.
• The “last electron enters” criterion determines the element’s block from its electron configuration.

💡 Memory Hook

Last electron decides the block: s→s-block, p→p-block, d→d-block, f→f-block.

📖 6. Configuration exceptions and noble gas notation

🔑 Key Concepts & Definitions

• Full box notation : Box notation represents orbitals as boxes and electrons as arrows, with arrow direction showing spin.
• Noble gas core notation : Noble gas notation replaces the filled inner shells with the nearest noble gas core written in brackets.
• Exception (Cr and Cu) : Cr and Cu do not follow the simple expected s-and-d filling order in their ground-state electron configurations.
• Half-filled stability : The half-filled arrangement is described as energetically more stable than incompletely filled arrangements for d-orbitals.

📝 Essential Points

• For Cr(24), the given exception configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴ 4s² and the next-step adjustment gives 3d⁵ 4s¹ in their practice set.
• For Cu(29), the given exception configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁹ 4s².
• The energetically more stable arrangements listed are half-filled (3d⁵ 4s¹) and fully filled (3d¹⁰ 4s¹).
• The stated reason for the exception is increased stability of fully filled or half-filled subshells compared with incomplete ones.
• In d-block ions formed from Cr-type behavior, the 4s electrons are removed before the 3d electrons.

💡 Memory Hook

Cr and Cu “prefer” 3d⁵4s¹ (half) and 3d¹⁰4s¹ (full), even if that means shifting from the simple fill pattern.

📖 7. Ions, free radicals and periodic trends

🔑 Key Concepts & Definitions

• Free radical : A free radical is a species with one or more unpaired electrons.
• Free radical dot notation : Dot notation shows an unpaired electron as a dot next to the element symbol.
• Ionic radius : Ionic radius is the effective size of an ion compared across periods and groups using periodic patterns.
• Cation : A cation forms when an atom loses electrons and has more positive charge than neutral.
• Anion : An anion forms when an atom gains electrons and has more negative charge than neutral.

📝 Essential Points

• Free radicals in the notes are represented with unpaired electrons shown as dots, for example Cl· and H3C·.
• For ion electron configurations in the notes, outer-sub-shell electrons are generally removed first when metals form positive ions.
• For Cr-type d-block ions, the note states the 4s sub-shell fills before 3d, but when electrons are lost the 4s electrons leave first.
• Examples given include Cl + 1e⁻ → Cl⁻ and Na → Na⁺ + 1e⁻.
• The notes connect periodic patterns with atomic and ionic radii and also ask to deduce trends that accompany ionisation energy patterns.

💡 Memory Hook

Unpaired electron = dot (free radical); metal ion loses outer electrons first.

📅 Key Dates

1. 400 BCEDemocritus proposed atoms (Atomos, indivisible/un-cuttable).
2. Early 1800sDalton developed the first scientific atomic theory.
3. 1836Mass comparison stated proton is 1836 times heavier than an electron.
4. 1913Bohr model introduced fixed circular orbits with definite energy.
5. 1924de Broglie wave nature of electrons provided a basis for probability.
6. 1925Pauli introduced the Pauli Exclusion Principle.
7. 1926Schrödinger’s quantum mechanical model described electrons by a wave equation.
8. 1927Hund introduced electron filling of orbitals to minimize spin-pair repulsion.

📊 Synthesis Tables

Comparison of atomic models

Model	Main idea	Key limitation
Thomson model	Atom contains sub-atomic particles (electrons) and is divisible	Couldn’t explain atomic spectral lines and failed to explain proton concentration in a nucleus.
Rutherford model	Atom is mostly empty space with a dense positive nucleus	Failed to explain spectral lines produced by atoms.
Bohr model	Electrons move in fixed energy levels and jump with quantised energy	Failed to explain spectral lines for multi-electron atoms and treated electrons only as particles.

⚠️ Common Pitfalls & Confusions

1. Students may mix up charge deflection directions, remembering incorrectly whether electrons move toward the positive or negative plate.
2. Students may confuse atomic number Z with mass number A, swapping neutrons = A − Z.
3. Students may think IE is the same for each step, forgetting that IE generally increases (IE₁ < IE₂ < IE₃).
4. Students may misapply block rules using the wrong “last electron enters” subshell rather than n, s, p, d, or f entry.
5. Students may apply the general filling order to Cr and Cu without allowing the given half-filled or fully filled stability exceptions.

✅ Exam Checklist

1. Explain Democritus’ atomos idea and Dalton’s main claims about identical atoms and whole-number ratios.
2. Describe how Thomson, Rutherford, and Bohr atomic models differ in their main idea and one stated limitation each.
3. Use cathode ray tube electric-field behavior to state which way protons, electrons, and neutrons deflect.
4. Calculate isotopic nucleus quantities using Z = protons, A = protons + neutrons, and neutrons = A − Z.
5. Determine electron numbers in a neutral atom, cation, and anion using the given “add or remove electrons” relationships.
6. Define first ionisation energy (IE₁) precisely in terms of removing 1 mole of electrons in the gaseous state.
7. Write and interpret successive ionisation energy equations as removing the next electron from the previous ion.
8. Predict where a large jump in successive IE occurs and link it to removal from a shell closer to the nucleus.
9. State the four electron-energy filling rules: Aufbau principle, Pauli exclusion, Hund’s rule, and orbital capacity (2 per orbital).
10. Identify orbital and sub-shell capacities (s has 1 orbital, p has 3, d has 5) and maximum electrons per orbital and sub-shell.
11. Represent electron configurations using full, shorthand with [noble gas], and box notation as specified (up to atomic number 36 for practice).
12. State the block of an element from where the last electron enters (s, p, d, or f) and the stated group ranges for s, p, and d blocks.
13. Use the given Cr and Cu exceptions to state the more stable half-filled and fully filled arrangements (3d⁵4s¹ and 3d¹⁰4s¹).
14. Define a free radical as having one or more unpaired electrons and recognize dot notation for the unpaired electron.