CAMBRIDGE INTERNATIONAL AS LEVEL

AS Chemistry (9701)

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1.1 Subatomic Particles and Nuclear Symbolism

Atoms consist of a small, dense positively charged nucleus containing protons and neutrons, surrounded by negatively charged electrons orbiting in energy levels . Most of the atomic mass is concentrated in the nucleus . Electrostatic forces of attraction hold the negative electrons and the positive nucleus together .

Subatomic Particle Properties

Particle Relative Charge Mass (atomic mass units)
Proton (p) +1 1
Neutron (n) 0 1
Electron (e-) -1 1 / 1840

Standard Isotopic Notation

Proton Number (Z): The total number of protons in the nucleus (defines the element) .

Nucleon Number (A): The total number of protons plus neutrons in the nucleus .

Neutron Count: Calculated as A - Z .

1.2 Atomic & Ionic Radius Trends

  • Across a Period (Left to Right): Atomic radii decrease because nuclear charge increases while inner-shell shielding remains constant, pulling outer electrons closer to the nucleus .
  • Down a Group (Top to Bottom): Atomic radii increase because extra quantum shells are added, placing outer electrons further away .
  • Cations (+): Smaller than parent atoms because lost electrons lead to increased nuclear attraction per remaining electron .
  • Anions (-): Larger than parent atoms due to electron-electron repulsion and weaker net nuclear attraction per electron .

1.3 Isotopes

Isotopes are atoms of the same element containing identical numbers of protons and electrons, but different numbers of neutrons . They exhibit identical chemical properties due to matching outer-shell electron configurations, but differ slightly in physical properties such as mass and density .

Interactive Atomic Builder

2.1 Quantum Shells and Subshells

Electrons reside in principal quantum shells numbered by n . Higher quantum numbers indicate higher energy and greater average distance from the nucleus .

Electron Capacities per Shell (2n2)

  • n = 1: Max 2 electrons (1s)
  • n = 2: Max 8 electrons (2s, 2p)
  • n = 3: Max 18 electrons (3s, 3p, 3d)
  • n = 4: Max 32 electrons (4s, 4p, 4d, 4f)

2.2 Orbital Geometry & Shapes

An atomic orbital is a region of space holding up to 2 electrons with opposite spins . Nodes represent points of zero electron probability; more nodes correlate with higher energy .

  • s orbitals: Spherical shape .
  • p orbitals: Dumbbell shapes aligned along three axes (px, py, pz) .
s Orbital (Spherical) p_x Orbital (Dumbbell) p_y Orbital (Dumbbell)

2.3 Energy Level Hierarchy

Subshells fill in order of increasing energy: 1s < 2s < 2p < 3s < 3p < 4s < 3d . Note that the 4s subshell is lower in energy than 3d, filling before 3d .

Anomalous Electronic Configurations

Copper (Cu, Z=29): [Ar] 3d10 4s1 (a fully filled d-subshell is particularly stable) .

Free Radicals: Species containing one or more unpaired electrons (e.g., Chlorine radical Cl•: 1s2 2s2 2p6 3s2 3p5) .

Subatomic Particles in Electric Fields

Observe trajectory deflection when charged particles travel through an electric field .

  • Electrons (e-): Deflect heavily and rapidly toward the positive plate due to their tiny mass .
  • Protons (p+): Deflect slightly toward the negative plate due to their much larger mass (~1840x electron mass) .
  • Neutrons (n0): Travel in a straight line with zero deflection due to neutrality .

4.1 Ionization Energy Definition & Factors

First Ionization Energy (IE): The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions .

Key Factors Influencing Ionization Energy

  • Nuclear Charge: Higher proton numbers create stronger electrostatic attraction for outer electrons → higher IE .
  • Atomic Radius: Greater distance between nucleus and valence electrons weakens attraction → lower IE .
  • Shielding Effect: Inner quantum shells block full nuclear charge from outer shell → lower IE .
  • Spin-Pair Repulsion: Paired electrons within an orbital experience mutual repulsion, making them easier to remove → lower IE .

4.2 Trend Across Period 2 & 3

Ionization energy generally increases across a period but displays characteristic dips at subshell changes (s2 → p1) and spin-pairing points (p3 → p4) .

Atomic Number (Z) 1st IE (kJ/mol) H He Li Be B C N O F Ne Na Mg

4.3 Successive Ionization Energies

Successive IE data provides evidence for shell structure. Large jumps in energy occur when an electron is removed from a new inner quantum shell closer to the nucleus .

Element Group 1st IE 2nd IE 3rd IE 4th IE Jump Location & Significance
Sodium (Na) Group 1 494 4560 6940 9540 Jump after 1st electron → 1 valence electron (3s1)
Magnesium (Mg) Group 2 736 1450 7740 10500 Jump after 2nd electron → 2 valence electrons (3s2)
Aluminium (Al) Group 3 577 1820 2740 11600 Jump after 3rd electron → 3 valence electrons (3s2 3p1)

Lesson Summary

  • Subatomic Mass Distribution: Nearly all atomic mass is concentrated in the central nucleus containing protons (+1) and neutrons (0) .
  • Isotopes: Same atomic number (Z), different nucleon number (A) . Same chemical properties, minor physical differences .
  • Subshell Order: Orbitals fill in increasing energy sequence: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 .
  • Field Behavior: Light electrons deflect strongly toward positive plates; heavy protons deflect weakly toward negative plates; neutrons pass undeflected .
  • Ionization Jumps: Sharp leaps in successive IE indicate breaking into a new inner quantum shell .

Check Your Understanding

1. Why is there a large jump between the 3rd and 4th ionization energies of Aluminium (Al)?




2. How does a beam of protons behave relative to electrons when fired into an electric field at equal velocity?