1.1 Particles in the atom and atomic radius

1. Atomic Structure Overview

   • Atoms are primarily empty space surrounding a tiny, dense nucleus.
   • The nucleus contains positively charged protons and neutral neutrons.
   • Negatively charged electrons are found in discrete energy levels (shells) in the empty space around the nucleus.

2. Sub-atomic Particles

   • Proton (p):
     • Relative charge: $+1$
     • Relative mass: $1$
   • Neutron (n):
     • Relative charge: $0$
     • Relative mass: $1$
   • Electron (e^-$):
     • Relative charge: -1
     • Relative mass: \frac{1}{1836} (negligible)

3. Key Terminology

   • Atomic Number (Z) / Proton Number: The number of protons in an atom's nucleus. It defines the element.
   • Mass Number (A) / Nucleon Number: The total number of protons and neutrons in an atom's nucleus.

4. Distribution of Mass and Charge

   • Mass: Concentrated in the nucleus (protons and neutrons account for nearly all the mass).
   • Charge: The nucleus is positively charged (due to protons). The electrons occupy the space around the nucleus, giving the atom an overall neutral charge (in atoms, protons = electrons).

5. Behaviour in an Electric Field

   • Beams of particles (protons, neutrons, electrons) moving at the same velocity in an electric field:
     • Protons: Deflected towards the negative plate (due to positive charge).
     • Neutrons: Not deflected (due to no charge).
     • Electrons: Deflected towards the positive plate (due to negative charge). Electrons are deflected more significantly than protons due to their much smaller mass-to-charge ratio.

6. Determining Sub-atomic Particles

   • For an atom:
     • Number of protons = Atomic Number (Z)
     • Number of electrons = Number of protons (for a neutral atom)
     • Number of neutrons = Mass Number (A) - Atomic Number (Z)
   • For an ion:
     • Number of protons = Atomic Number (Z)
     • Number of electrons = Number of protons - charge (for positive ions) or Number of protons + charge (for negative ions)
     • Number of neutrons = Mass Number (A) - Atomic Number (Z)

7. Atomic and Ionic Radius Trends

   • Across a Period (left to right):
     • Atomic radius decreases. Nuclear charge increases, pulling outer electrons closer. Shielding by inner electrons remains largely constant.
     • Ionic radius generally decreases for positive ions (cations) and generally decreases for negative ions (anions) separately, but expands significantly when transitioning from cations to anions (e.g., from Group 1 to Group 17, the anions are much larger than the cations of the same period).
   • Down a Group (top to bottom):
     • Atomic radius increases. More electron shells are added, increasing the distance between the nucleus and outer electrons, despite increased nuclear charge. Shielding also increases.
     • Ionic radius increases. Similar reason to atomic radius due to the addition of electron shells.

1.2 Isotopes

1. Definition of Isotope

   • Isotopes are atoms of the same element that have the same number of protons (atomic number) but a different number of neutrons (and thus a different mass number).

2. Isotope Notation

   • The notation for isotopes is \text{ }_y^x A
     • x: Mass number (A) or Nucleon number (protons + neutrons)
     • y: Atomic number (Z) or Proton number
     • A: Chemical symbol of the element
   • Example: Carbon-12 is \text{ }6^{12} C$, Carbon-14 is$\text{ }6^{14} C

3. Chemical Properties of Isotopes

   • Isotopes of the same element have the same chemical properties.
   • Explanation: Chemical properties are determined by the number and arrangement of electrons, especially the valence electrons. Since isotopes have the same number of protons, they also have the same number of electrons (in neutral atoms) and thus the same electronic configuration.

4. Physical Properties of Isotopes

   • Isotopes of the same element have different physical properties.
   • Explanation (limited to mass and density):
     • Mass: Isotopes have different numbers of neutrons, leading to different mass numbers, so their actual masses differ.
     • Density: Since density is mass per unit volume (\text{density} = \frac{\text{mass}}{\text{volume}}), and isotopes have different masses but similar atomic volumes, their densities will differ.

1.3 Electrons, energy levels and atomic orbitals

(Ground state only; elements hydrogen to krypton assessed)

1. Key Terms in Electronic Structure

   • Shells (Principal Energy Levels): Main energy levels where electrons are found. Designated by the principal quantum number (n = 1, 2, 3, …).
   • Sub-shells: Divisions within shells, each with a specific shape ($s, p, d, f$).
   • Orbitals: Regions within sub-shells where there is a high probability of finding an electron. Each orbital can hold a maximum of two electrons with opposite spins.
   • Principal Quantum Number (n$)</strong>: Indicates the main energy level of an electron ($n = 1, 2, 3, …$). Higher$n means higher energy and further from the nucleus.
   • Ground State: The lowest possible energy state of an atom or ion, where electrons occupy the lowest available energy levels.

2. Sub-shells and Electron Capacity

   • s sub-shell: 1 orbital, maximum 2 electrons (e.g., 1s, 2s)
   • p sub-shell: 3 orbitals, maximum 6 electrons (e.g., 2p, 3p)
   • d sub-shell: 5 orbitals, maximum 10 electrons (e.g., 3d, 4d)

3. Order of Increasing Energy of Sub-shells

   • The general order for the first three shells and 4s/4p:
     1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p
   • (Note: 4s$fills before$3d$for most elements due to lower energy levels initially, though$3d becomes lower in energy for transition metals once filled).

4. Electronic Configurations

   • Shows the distribution of electrons in shells, sub-shells, and orbitals.
   • Full electronic configuration: Lists all occupied sub-shells from lowest to highest energy.
     • Example for Fe (Z=26): 1s^22s^22p^63s^23p^63d^64s^2
   • Shorthand (condensed) electronic configuration: Uses the symbol of the preceding noble gas to represent the filled inner shells.
     • Example for Fe: [Ar] 3d^64s^2

5. Explanation of Electronic Configurations

   • Energy of electrons: Electrons fill orbitals in order of increasing energy (Aufbau principle).
   • Inter-electron repulsion: Electrons within the same orbital repel each other, which is minimized by electrons occupying separate orbitals within a sub-shell before pairing up (Hund's Rule) and by having opposite spins (Pauli Exclusion Principle).

6. Determining Electronic Configuration

   • Given atomic/proton number and charge, apply the Aufbau principle, Hund's Rule, and Pauli Exclusion Principle to fill electrons into sub-shells.
   • For ions, adjust the number of electrons (remove for cations, add for anions) from the highest energy orbitals (or 4s before 3d for transition metal cations).

7. Electrons in Boxes Notation

   • Visual representation of electrons in orbitals, using boxes for orbitals and arrows for electrons (up and down for spin).
   • Example for Fe ([Ar] 3d^64s^2):
     • [Ar]
     • \text{4s: } \uparrow \downarrow
     • \text{3d: } \uparrow \downarrow \quad \uparrow \quad \uparrow \quad \uparrow \quad \uparrow

8. Shapes of s and p Orbitals

   • s orbital: Spherical shape.
   • p orbital: Dumbbell shape, oriented along x, y, z axes ($px, py, p_z$).

9. Free Radical

   • A free radical is a species (atom, ion, or molecule) that contains one or more unpaired electrons. These are highly reactive.

1.4 Ionisation energy

(Ground state only; elements hydrogen to krypton assessed)

1. Definition of First Ionisation Energy (IE)

   • First Ionisation Energy (IE_1$)</strong>: The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous$1+ ions.

2. Equations for Ionisation Energies

   • First Ionisation Energy (IE_1):
     X(g) \rightarrow X^+(g) + e^- (where X is an atom)
   • Second Ionisation Energy (IE_2):
     X^+(g) \rightarrow X^{2+}(g) + e^-$$ (energy to remove the second electron)
   • Subsequent Ionisation Energies: Each successive IE removes an electron from an increasingly positively charged ion, thus requiring more energy.

3. Trends in Ionisation Energies

   • Across a Period (left to right):
     • General trend: Ionisation energy increases.
     • Reason: Nuclear charge increases, and atomic radius decreases. The outer electrons are more strongly attracted to the nucleus, requiring more energy to remove.
     • Exceptions: Decreases from Group 2 to 13 (e.g., Be to B) due to the first electron in a p orbital being slightly higher in energy and experiencing some shielding from s electrons. Decreases from Group 15 to 16 (e.g., N to O) due to spin-pair repulsion in the p sub-shell (first electron removed from a paired electron in p orbital).
   • Down a Group (top to bottom):
     • General trend: Ionisation energy decreases.
     • Reason: Atomic radius increases (more electron shells), and shielding by inner electrons increases. The attraction between the nucleus and the outer electrons decreases, making them easier to remove.

4. Variation in Successive Ionisation Energies

   • Successive ionisation energies always increase because each subsequent electron is removed from an ion with an increasing positive charge, leading to a stronger electrostatic attraction from the nucleus.
   • Large jumps (steps) occur when an electron is removed from a new, inner electron shell. This indicates moving from a valence electron to a core electron which is much closer to the nucleus and less shielded.

5. Ionisation Energies and Nuclear Attraction

   • Ionisation energies are a direct measure of the attraction between the positive nucleus and the negatively charged outer electron being removed.

6. Factors Influencing Ionisation Energies

   • Nuclear Charge: Higher nuclear charge increases attraction for electrons, increasing IE.
   • Atomic/Ionic Radius: Smaller radius means electrons are closer to the nucleus, increasing attraction and IE.
   • Shielding by Inner Shells and Sub-shells: Inner electrons repel outer electrons, reducing the effective nuclear charge felt by outer electrons (shielding effect). More shielding decreases IE.
   • Spin-pair Repulsion (within sub-shells): When two electrons occupy the same orbital, their mutual repulsion makes it slightly easier to remove one of them, leading to a slightly lower IE (e.g., Group 16 vs. Group 15).

7. Deducing Electronic Configurations from Successive IE Data

   • Analyze the pattern of successive IE values. Large jumps indicate the removal of an electron from a new, inner electron shell.
   • The number of electrons removed before a significant jump indicates the number of valence electrons.
   • This helps determine the group an element belongs to and its electronic configuration (e.g., 1 large jump after first electron suggests Group 1, 2 electrons before large jump suggests Group 2, etc.).

8. Deducing Position in Periodic Table from Successive IE Data

   • By determining the number of valence electrons (from the significant jump in successive IEs), one can deduce the element's group number. For example, if there is a big jump after the 3rd IE, the element has 3 valence electrons, belonging to Group 13.
   • The number of main shells (principal quantum numbers) can sometimes be inferred from the total number of electrons and the IE values, helping to estimate the period.