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Core Idea 1: Matter — Section 1.1 Atomic Structure

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Atomic Orbitals & Configuration Rules

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The Concept of an Atomic Orbital

Electrons do not travel along fixed, localized circular tracks or circular orbits.
Instead, they traverse regions of space surrounding the positive core called atomic orbitals.
An orbital represents a probability boundary context—specifically, a region holding a 98% probability chance of locating an electron.
Orbitals differ distinctly in 3D geometry shapes: s, p, d, and f.

Shell Capacities: n, n², and 2n²

A principal quantum shell groups orbitals located roughly the same distance out from the core nucleus.
For any principal quantum number n:
• Total sub-shells available = n
• Total unique individual orbitals = n²
• Maximum total electron assignment capacity = 2n²

Principal Shell (n)	Total Orbitals (n²)	Orbital Type Breakdown	Max Electron Capacity (2n²)
n = 1	1	1s (one)	2
n = 2	4	2s (one), 2p (three)	8
n = 3	9	3s (one), 3p (three), 3d (five)	18

The Trio of Electron Assignment Rules

1. Aufbau Principle: Electrons must progressively fill lower-energy ground states first. Note: The 4s sub-shell has lower energy than 3d and fills first!

2. Pauli Exclusion Principle: Each individual orbital slot can hold a maximum limit of two electrons, and they must maintain **opposite (anti-parallel) spins** to minimize electrical repulsion.

3. Hund’s Rule of Multiplicity: Degenerate orbitals (orbitals of identical energy like p_x, p_y, p_z) must be occupied **singly first** with parallel spins before pairing begins.

Key Teaching Takeaways

Orbital Visualizations: Ensure students can confidently draw s and p profiles complete with labeled Cartesian 3D axes (x, y, z).
Energy Convergence: Emphasize the convergent nature of successive sub-shells, highlighting the structural overlap where 4s sits below 3d during filling configurations.
Writing Exercises: Test students thoroughly on ‘electrons-in-boxes’ representations to spot spin violations or premature single orbital pairings.

Section 1.1: Atomic & Electronic Structure (Comprehensive)

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Orbitals, Shell Dynamics & Full Configuration Rules

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Atomic Orbitals & Spatial Shapes

Electrons are not localized in fixed planetary orbits around the nucleus[cite: 59, 60].
An atomic orbital is a mathematical region of space around a free nucleus where the probability of finding a particular electron is greatest (~98% chance)[cite: 61].
s Orbitals: Spherical shape centered on the nucleus[cite: 64].
p Orbitals: Dumb-bell-shaped; three degenerate types ($p_x, p_y, p_z$) pointing along the x-, y-, and z-axes respectively. Each consists of two identical lobes[cite: 64].

The Five Degenerate d-Orbitals

There are five separate types of d-orbitals which are all degenerate (share identical baseline energy levels):

d_xy, d_yz, d_zx Orbitals: Each features four distinct geometric lobes situated cleanly between the respective xy, yz, and zx coordinate planes[cite: 64].
d_x²-y² Orbital: Contains four matching lobes positioned directly along the x- and y-axes[cite: 64].
d_z² Orbital: Uniquely shaped compared to the other four; consists of two prominent vertical lobes positioned along the z-axis combined with a central toroidal 'ring' running along the middle plane[cite: 65].

Note: In syllabus drawing questions, Cartesian x-, y-, and z-axes must always be explicitly illustrated to demonstrate 3D spatial orientation[cite: 65].

Quantum Shell Rules & Capacity Limits

A principal quantum shell comprises groups of sub-shells situated at a comparable radial distance from the nucleus[cite: 66]. In the n^th shell, there are exactly n sub-shells, n² individual orbitals, and a maximum population cap of 2n² electrons.

Shell Index (n)	Orbitals in Shell (n²)	Sub-shell Types	Orbital Breakdown per Type	Max Electron Capacity
n = 1	1	s	1s (one)	2
n = 2	4	s, p	2s (one), 2p (three)	8
n = 3	9	s, p, d	3s (one), 3p (three), 3d (five)	18
n = 4	16	s, p, d, f	4s (one), 4p (three), 4d (five), 4f (seven)	32

The Three Strict Core Configuration Rules

1. Aufbau Principle (Building Up): Incoming electrons must always fill the orbital with the lowest available energy state first[cite: 75, 76]. Due to relative radial shielding, the 4s orbital has lower energy than 3d orbitals and must fill first[cite: 74].

2. Pauli Exclusion Principle: An individual orbital can hold a maximum of two electrons[cite: 76]. These paired electrons are only stable when spinning in opposite directions (↑↓)—their resulting magnetic attraction offsets their mutual electrostatic charge repulsion[cite: 76, 80].

3. Hund's Rule of Multiplicity: When filling degenerate sub-shells (like identical energy 2p or 3d slots), orbitals must be occupied singly first with parallel spins before pairing occurs[cite: 77]. Electrons only begin to pair up once the sub-level is exactly half-filled[cite: 78].

Syllabus Exceptions: d-Block Anomalies

Chromium (₂₄Cr) Exception: Configuration is [Ar] 3d⁵ 4s¹ instead of [Ar] 3d⁴ 4s²[cite: 86, 93]. This occurs because half-filled 3d and 4s sub-shells exhibit superior structural stability[cite: 93].
Copper (₂₉Cu) Exception: Configuration is [Ar] 3d¹⁰ 4s¹ instead of [Ar] 3d⁹ 4s² [cite: 86, 94] because a completely filled 3d sub-shell maximizes stabilization energy.
Transition Metal Cation Ionisation: Crucially, 4s electrons are removed first during positive ion formation. Once 3d orbitals are occupied, they repel the 4s orbital to a slightly higher energy level.
• ₂₆Fe: [Ar] 3d⁶ 4s² ➤ ₂₆Fe²⁺: [Ar] 3d⁶ ➤ ₂₆Fe³⁺: [Ar] 3d⁵

Ground State vs. Excited State Dynamics

Ground State: An atom is in its ground state when all its electrons reside in the available energy orbitals of lowest potential level. Most elements sit in this stable baseline at room temperature[cite: 95].
Excited State: Occurs when one or more electrons absorb standalone energy packets and are actively promoted to a higher energy level[cite: 101].
Excited configurations violate the Aufbau timeline sequence but still comply with individual orbital spin safety thresholds[cite: 95, 101].

                        Carbon Ground State: 1s² 2s² 2px¹ 2py¹ 2pz⁰ [cite: 102]

                        Carbon Excited State (C*): 1s² 2s¹ 2px¹ 2py¹ 2pz¹ [cite: 102]

Section 1.1: Atomic & Electronic Structure

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Orbitals, Shell Dynamics & Configuration Rules

A-Level Chemistry Interactive Lecture Module

Atomic Orbitals & Spatial Shapes

Electrons are not localized in fixed planetary orbits around the nucleus[cite: 181, 182].
An atomic orbital is a mathematical region of space around a free nucleus where the probability of finding a particular electron is greatest (~98% chance)[cite: 182, 183].
s Orbitals: Completely spherical shape centered symmetrically on the nucleus[cite: 187].
p Orbitals: Dumb-bell-shaped; three degenerate types ($p_x, p_y, p_z$) pointing directly along the x-, y-, and z-axes respectively, featuring two identical lobes[cite: 188, 196, 197, 198].

The Five Degenerate d-Orbitals

There are five distinct types of d-orbitals which are completely degenerate (possess identical energy profiles)[cite: 199, 211]:

d_xy, d_yz, d_zx Orbitals: Each features four distinct geometric lobes situated cleanly between the respective coordinate planes[cite: 207, 208].
d_x²-y² Orbital: Contains four matching lobes positioned explicitly along the x- and y-axes[cite: 209].
d_z² Orbital: Shape differs from the other four; consists of two vertical lobes along the z-axis combined with a unique central horizontal 'ring'[cite: 210].

Note: In syllabus drawings, 3D Cartesian axes must always be shown to accurately illustrate directional properties[cite: 212].

Quantum Shell Rules & Capacity Limits

A principal quantum shell groups orbitals located roughly the same distance out from the core nucleus[cite: 218]. In the n^th shell, there are exactly n sub-shells, n² individual orbitals, and a maximum population cap of 2n² electrons[cite: 222].

Shell Index (n)	Orbitals in Shell (n²)	Sub-shell Types	Orbital Breakdown per Type	Max Electron Capacity
n = 1	1	s	1s (one)	2
n = 2	4	s, p	2s (one), 2p (three)	8
n = 3	9	s, p, d	3s (one), 3p (three), 3d (five)	18
n = 4	16	s, p, d, f	4s (one), 4p (three), 4d (five), 4f (seven)	32

The Three Strict Core Configuration Rules

1. Aufbau Principle: Incoming electrons must progressively occupy the lowest available energy orbital first[cite: 247]. Due to energy convergence patterns, the 4s orbital holds lower baseline energy than 3d orbitals and must fill first[cite: 244, 246].

2. Pauli Exclusion Principle: Each individual orbital can hold a maximum of two electrons[cite: 248]. These paired electrons remain stable only when spinning in opposite directions (↑↓) to create a balancing magnetic attraction[cite: 248, 258].

3. Hund's Rule of Multiplicity: When filling a degenerate sub-shell, orbitals must be occupied singly first (maintaining parallel electron spins) before pairing begins[cite: 249].

Syllabus Exceptions: d-Block Anomalies

Chromium (₂₄Cr) Exception: Ground state is [Ar] 3d⁵ 4s¹ instead of [Ar] 3d⁴ 4s²[cite: 341, 363, 364]. Half-filled 3d and 4s sub-shells yield enhanced structural stability[cite: 363].
Copper (₂₉Cu) Exception: Ground state is [Ar] 3d¹⁰ 4s¹[cite: 347, 365]. A completely filled 3d sub-shell maximizes baseline electronic stabilization[cite: 364].
Transition Metal Ionisation: Crucially, 4s electrons are removed first during positive ion formation[cite: 366]. Once 3d orbitals fill, they repel 4s electrons to a slightly higher energy state[cite: 367, 368].
• ₂₆Fe: [Ar] 3d⁶ 4s² ➤ ₂₆Fe²⁺: [Ar] 3d⁶ ➤ ₂₆Fe³⁺: [Ar] 3d⁵ [cite: 369, 371, 372]

Ground State vs. Excited State Dynamics

Ground State: Configuration where all internal electrons reside strictly within the lowest available structural energy levels[cite: 374].
Excited State: Occurs when one or more electrons absorb external energy and are actively promoted to a higher energy level[cite: 387].
Excited configurations deviate directly from the Aufbau baseline sequence, but still strictly observe individual orbital pairing rules[cite: 390].

                        Carbon Ground State: 1s² 2s² 2px¹ 2py¹ 2pz⁰ [cite: 389]

                        Carbon Excited State (C*): 1s² 2s¹ 2px¹ 2py¹ 2pz¹ [cite: 390]

Section 1.1: Atomic Structure — The Nucleus & Field Trajectories

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Anatomy of the Atom & Field Deflections

A-Level Chemistry Core Lecture Module

Fundamental Sub-Atomic Particles

Nucleons: Protons and neutrons are tightly packed together inside the central nucleus, making it incredibly dense[cite: 29, 38].
Extranuclear Region: Electrons move rapidly through the vast, empty space surrounding the core[cite: 29, 40].
Mass vs Volume: The nucleus contains nearly all the mass of the atom but takes up very little of its actual volume[cite: 39]. The diameter of an atom is about 10⁵ times larger than its nucleus[cite: 44].

Particle	Relative Mass (a.m.u.)	Relative Charge
Proton (p)	1	+1
Neutron (n)	1	0
Electron (e⁻)	1 / 1840	-1

Behavior in an Electric Field

Direction of Deflection: Protons (positive) are attracted toward the negative plate[cite: 54, 58]. Electrons (negative) bend toward the positive plate[cite: 54, 58]. Neutrons carry no charge and pass straight through without any deflection[cite: 61, 61].
Magnitude of Deflection: Lighter particles curve far more than heavier ones[cite: 64]. The angle of displacement is governed by:

Magnitude of Deflection \propto charge / mass

Because an electron's mass is extremely small compared to a proton's, electrons show a vastly larger deflection angle[cite: 66].

Deducing Relative Deflection Angles

Problem Context: A beam of ¹H⁺ nuclei (protons) passes through an electric field and deflects at an angle of 4.0°[cite: 86]. Calculate the expected deflection angles for the following nuclear particles under identical field conditions:

1. Deuteron Nuclei (²H⁺)

Charge/Mass Ratio = 1 / 2 = 0.5 [cite: 88] Expected Angle = 0.5 \times 4.0° = 2.0° (Bends toward the negative plate) [cite: 90]

2. Helium Nuclei (⁴He²⁺ / Alpha Particles)

Charge/Mass Ratio = 2 / 4 = 0.5 [cite: 76] Expected Angle = 0.5 \times 4.0° = 2.0° (Deflects exactly the same amount as ²H⁺) [cite: 79]

Core Summary for Students

Atomic Composition: Atoms contain vast regions of empty space[cite: 40]. The small, highly dense central nucleus is responsible for the overall mass and positive charge concentration[cite: 38, 39].
Field Rules: The charge of a sub-atomic particle determines its absolute direction of path deflection, while the overall charge-to-mass ratio controls the final magnitude of that deflection curve[cite: 62, 63].
Analytical Strategy: Always verify whether a question refers to neutral atoms or isolated nuclei before using mass metrics to compute field displacement behaviors[cite: 20].

Section 1.1: Atomic Structure — Notations, Isotopes & Terms

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Nuclides, Isotopes & Chemical Classifications

A-Level Chemistry Interactive Lecture Module

Nuclide Notation Framework

A nuclide represents a unique atomic nucleus with a specific count of protons and neutrons[cite: 100].
Standard baseline representation follows the ^A_ZX system[cite: 101]:

A = Nucleon Number (Mass Number): The combined total number of protons and neutrons present in the nucleus[cite: 101].

Z = Proton Number (Atomic Number): The total number of protons, which fundamentally dictates the core identity of the element[cite: 102, 110].

Subtracting the proton index directly yields the absolute neutron balance within the species: (A - Z)[cite: 103].

Tracking Sub-Atomic Particle Balances

Neutral Atoms: The count of negative electrons perfectly balances the internal proton count (Electrons = Protons)[cite: 107].
Anions (Negative Ions): Formed by adding extra electrons to a neutral atom (Electrons > Protons)[cite: 108].
Cations (Positive Ions): Formed by removing valence electrons from the atomic shells (Electrons < Protons)[cite: 109].

Species Type	Protons (Z)	Neutrons (A - Z)	Electrons
¹⁶₈O (Atom)	8 [cite: 105]	16 - 8 = 8 [cite: 105]	8 [cite: 105]
¹⁶₈O²⁻ (Anion)	8 [cite: 105]	16 - 8 = 8 [cite: 105]	8 + 2 = 10 [cite: 105]
²³₁₁Na⁺ (Cation)	11 [cite: 105]	23 - 11 = 12 [cite: 105]	11 - 1 = 10 [cite: 105]

Relational Terminology distinctions

Isotopes: Atoms of the identical element containing the same proton count but differing neutron totals[cite: 122]. They share identical chemical actions (controlled by electrons) but vary in mass-dependent physical parameters like density and rates of diffusion[cite: 126].

Isoelectronic: Structures, atoms, or ions that contain the exact same net electron count (e.g., H₂O, Na⁺, and Al³⁺ each possess 10 total electrons)[cite: 128, 129].

Isotonic: Specific nuclear configurations that possess the identical net number of neutrons[cite: 130].

Core Summary for Students

Chemical Uniformity: Because isotopes possess identical electron counts and spatial valence arrangements, they display completely identical reactivity in chemical processes[cite: 126].
Ion Conversions: Remind students that net changes in structural configuration state charges impact the electron index count exclusively[cite: 108, 109]. The internal proton context of the dense nucleus stays completely unchanged[cite: 108, 109].
Exam Precision: Ensure students can quickly match definitions without confusing isotonic (neutrons) with isoelectronic (electrons) during complex composition evaluations[cite: 128, 130].

Section 1.1: Atomic Structure — Orbitals & Configuration Rules

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Orbitals, Shell Dynamics & Configuration Rules

A-Level Chemistry Interactive Lecture Module

Atomic Orbitals & Spatial Shapes

Electrons do not travel in fixed orbits around the nucleus; they are not localized.
An atomic orbital is a region of space round the nucleus in which the probability of finding a particular electron is the greatest (~98% chance).
s Orbitals: Spherical in shape and centered symmetrically around the nucleus.
p Orbitals: Dumb-bell-shaped; three degenerate types ($p_x, p_y, p_z$) consisting of two lobes along the x-, y-, and z-axes respectively.

The Five Degenerate d-Orbitals

Electrons can occupy s, p, d, and f orbitals. There are five distinct degenerate types of d-orbitals:

d_xy, d_yz, d_zx Orbitals: Each consists of four lobes of the same size and shape lying on the xy, yz, and zx planes respectively (between the axes).
d_x²-y² Orbital: Consists of four lobes located directly along the x- and y-axes.
d_z² Orbital: Uniquely shaped compared to the other four; it consists of two lobes along the z-axis with a 'ring' in the middle.

Note: When drawing shapes of orbitals, the x-, y-, and z-axes must be shown to illustrate their 3-D directional properties.

Quantum Shell Rules & Capacity Limits

A shell is a group of orbitals that are about the same distance out from the nucleus. In the n^th shell, there are exactly n sub-shells, n² individual orbitals, and a maximum population cap of 2n² electrons.

Shell Index (n)	Orbitals in Shell (n²)	Sub-shell Types	Orbital Breakdown per Type	Max Electron Capacity
n = 1	1	s	1s (one)	2
n = 2	4	s, p	2s (one), 2p (three)	8
n = 3	9	s, p, d	3s (one), 3p (three), 3d (five)	18
n = 4	16	s, p, d, f	4s (one), 4p (three), 4d (five), 4f (seven)	32

The Three Rules of Electronic Configurations

1. Aufbau Principle: An added electron will always occupy the orbital with the lowest energy first. Orbitals are filled in strict order of increasing energy starting from 1s. Note: The 4s orbital has a slightly lower energy than the 3d orbitals and fills first.

2. Pauli Exclusion Principle: Each individual orbital can hold a maximum of two electrons. Paired electrons can only be stable when they spin in opposite directions (↑↓) so that their resulting magnetic attraction counterbalances identical charge repulsion.

3. Hund's Rule of Multiplicity: When filling a sub-shell, each orbital must be occupied singly first (keeping electron spins parallel/same) before they are occupied in pairs.

Syllabus Exceptions: d-Block Anomalies

Chromium (₂₄Cr) Exception: Ground state is [Ar] 3d⁵ 4s¹ instead of [Ar] 3d⁴ 4s². This is because half-filled 4s and 3d sub-shells are more stable.
Copper (₂₉Cu) Exception: Ground state is [Ar] 3d¹⁰ 4s¹ instead of [Ar] 3d⁹ 4s² because a completely filled 3d sub-shell is more stable.
Transition Metal Cation Ionisation: Crucially, 4s electrons are removed first in the formation of positive ions. Once the 3d orbitals are occupied by electrons, these repel the 4s orbital to a slightly higher energy level.
• ₂₆Fe: [Ar] 3d⁶ 4s² ➤ ₂₆Fe²⁺: [Ar] 3d⁶ ➤ ₂₆Fe³⁺: [Ar] 3d⁵

Ground State vs. Excited State Dynamics

Ground State: An atom is in the ground state when the electrons are in the orbitals of the lowest available energy level. Most atoms are in this state at room temperature.
Excited State: An atom is in an excited state when one or more electrons absorb energy and are actively promoted to a higher energy level.
Excited configurations deviate from the standard Aufbau sequence but still follow individual orbital max spin thresholds.

                        Carbon Ground State: 1s² 2s² 2px¹ 2py¹ 2pz⁰

                        Carbon Excited State (C*): 1s² 2s¹ 2px¹ 2py¹ 2pz¹

Section 1.1: Atomic Structure — Ionisation Energies

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Ionisation Energies & Successive Data Interpretation

Section 1.1: Atomic Structure — Ionisation Energies

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Ionisation Energies & Successive Data Interpretation

Section 1.1: Atomic Structure — Ionisation Energies

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Ionisation Energies & Successive Data Interpretation

A-Level Chemistry Interactive Lecture Module

Defining Ionisation Energies

Ionisation energies provide clear experimental proof that electrons occupy discrete, quantized energy levels. [cite: 478]
First Ionisation Energy (1ˢᵗ I.E.): The energy required to remove one electron from each atom in a mole of gaseous atoms to produce one mole of gaseous unipositive cations. [cite: 479]
Successive Levels: Values grow progressively larger with each removal because the remaining electrons experience a stronger attraction from the constant nuclear charge. [cite: 492]

Syllabus Equation Standards (State symbols mandatory): [cite: 490]

1ˢᵗ I.E. : X(g) ➔ X⁺(g) + e⁻ [cite: 482, 483] 2ⁿᵈ I.E. : X⁺(g) ➔ X²⁺(g) + e⁻ [cite: 484, 485] 3ʳᵈ I.E. : X²⁺(g) ➔ X³⁺(g) + e⁻ [cite: 486, 487]

Factors Influencing Ionisation Energy

The absolute magnitude of an element's ionisation energy depends on three interrelated physical factors: [cite: 494]

1. Size of Positive Nuclear Charge: As the proton count increases, the electrostatic attraction for the outermost electron increases, requiring more energy to remove it (I.E. increases). [cite: 494, 507, 508]

2. Distance from Nucleus (Atomic/Ionic Radius): As atomic size increases, the valence outer shells sit further away from the positive core, decreasing attraction and lowering the energy threshold (I.E. decreases). [cite: 509, 510, 511]

3. Screening (Shielding) Effect: Inner shells of core electrons repel the outermost valence electrons, screening them from the positive nucleus. Greater shielding weakens core attraction (I.E. decreases). [cite: 512, 513, 514, 515]

Deducing Groups from Successive Jumps

Plotting successive log values reveals massive, sharp energy jumps that indicate a change in the principal quantum shell. [cite: 537, 564, 585]
A large jump occurs because an electron is stripped from an inner shell that is closer to the nucleus and experiences much weaker shielding. [cite: 564, 566]

Worked Example: Identifying an Unknown Element [cite: 588]

Given data (kJ mol⁻¹): 790, 1600, 3200, 4400, (MASSIVE JUMP) ➔ 16100, 19800, 23800 [cite: 588, 590]

Analysis: The massive jump between the 4ᵗʰ and 5ᵗʰ values reveals that the first 4 electrons are easily removed from the outer valence shell, while the 5ᵗʰ belongs to a core inner shell. [cite: 592]
Conclusion: The element has 4 valence electrons, putting it in Group 14. [cite: 593]

Evidence for Sub-Shell Structures

Analyzing successive entries within a single shell reveals smaller, steady rises followed by sharp internal discrepancies. [cite: 540, 571]
For example, tracking the inner shell of Potassium shows a steady rise for the first 6 electrons, followed by a noticeable internal jump for the final 2. [cite: 571]
Structural Proof: This clear shift indicates that principal shell 3 is further divided into sub-shells: a higher energy 3p sub-shell holding 6 electrons, and a lower energy 3s sub-shell holding 2. [cite: 572, 573]

Real-World Physics: Street Lamps

Orange-tinted street lamps contain solid sodium alongside a small amount of neon gas. Light is generated when these gaseous atoms are ionised within an applied electric field. [cite: 598]
The Red Shift Delay: When first switched on, the lamp emits a distinct red glow. This occurs because neon is already a gas and ionises immediately, while solid sodium must absorb heat to vaporise before it can undergo ionisation. [cite: 599, 601, 602]
The Orange Domination: After a short period, the light transitions to a bright orange color. This shift happens because sodium's 1ˢᵗ I.E. is much lower than neon's (494 kJ mol⁻¹ vs 2080 kJ mol⁻¹), allowing it to ionise far more efficiently once vaporised. [cite: 599, 603, 604]