1.1. ATOMIC STRUCTURE

How are mass and charge distributed in an atom?

• Mass is concentrated in the nucleus (protons and neutrons).

• Electrons contribute very little to mass but create a cloud of negative charge.

• Electrostatic attraction between the positive nucleus and negative electrons holds the atom together.

How do protons, neutrons, and electrons behave in an electric field?

• Electrons: Deflected strongly toward the positive plate (light and negatively charged).

• Protons: Deflected weakly toward the negative plate (heavier and positively charged).

• Neutrons: Not deflected (neutral charge).

How can you determine the numbers of protons, neutrons, and electrons in atoms and ions?

• Protons: Equal to the atomic number (Z).

• Neutrons: Mass number (A) - Atomic number (Z).

• Electrons:

  • Neutral atom: Same as protons.

  • Positive ion: Fewer electrons than protons.

  • Negative ion: More electrons than protons.

What is atomic radius, and how does it vary across a period and down a group?

• Definition: Atomic radius is half the distance between the nuclei of two covalently bonded atoms of the same type.

• Across a Period:

  • Atomic radius decreases due to increased positive nuclear charge pulling electrons closer.

  • Extra electrons are added to the same principal quantum shell, enhancing nuclear attraction.

• Down a Group:

  • Atomic radius increases due to additional electron shells.

  • Inner electrons shield outer electrons from the positive nucleus, weakening nuclear pull.

How does the electron shell theory explain atomic radius trends?

• Across a Period:

  • Increased nuclear charge pulls electrons closer, reducing atomic size.

• Down a Group:

  • More electron shells increase shielding, reducing nuclear attraction and increasing atomic size.

• Special Case:

  • Atomic radius increases sharply between noble gases and alkali metals due to an extra principal quantum shell.

What is ionic radius, and how does it vary with charge?

• Definition: Ionic radius is the size of an ion.

• Negative Ions:

  • Formed by gaining electrons.

  • Increased electron repulsion weakens nuclear pull, increasing ionic radius.

  • Greater negative charge → Larger ionic radius.

• Positive Ions:

  • Formed by losing electrons.

  • Fewer electrons experience stronger nuclear attraction, decreasing ionic radius.

  • Greater positive charge → Smaller ionic radius.

How does the electron shell theory explain ionic radius trends?

• Negative Ions:

  • Extra electrons increase repulsion, weakening nuclear pull and enlarging the ion.

• Positive Ions:

  • Fewer electrons experience stronger attraction, shrinking the ion.

What are isotopes?

• Isotopes are atoms of the same element with the same number of protons and electrons but a different number of neutrons.

• Example:

  • Carbon-12 (6 neutrons) and Carbon-14 (8 neutrons) are isotopes of carbon.

• Isotopes are named using the chemical symbol followed by a dash and the mass number (e.g., Carbon-12).

How are isotopes represented using notation?

x = Mass number (protons + neutrons). y = Atomic number (protons). A = Chemical symbol.

Example:

• Chlorine-35: Mass number = 35, Atomic number = 17, Chemical symbol = Cl.

• Chlorine-37: Mass number = 37, Atomic number = 17, Chemical symbol = Cl.

Why do isotopes of the same element have the same chemical properties?

• Isotopes have the same number of electrons in their outer shells.

• Electrons determine chemical reactions, so isotopes of the same element behave identically in chemical processes.

Why do isotopes of the same element have different physical properties?

• Isotopes differ in the number of neutrons, which affects their mass and density.

• Neutrons are neutral and only add mass to the atom.

• Example:

  • Chlorine-35 has a lower mass and density than Chlorine-37 due to fewer neutrons

What are shells, sub-shells, and orbitals in an atom?

• Shells: The main energy levels around the nucleus, numbered using principal quantum numbers (n).

  • n = 1 (up to 2 electrons), n = 2 (up to 8 electrons), n = 3 (up to 18 electrons), n = 4 (up to 32 electrons).

• Sub-shells: Divisions of shells, labeled as s, p, d, and f. Their energy increases as s < p < d < f.

  • Example: n = 1 only has an s sub-shell, while n = 3 has s, p, and d sub-shells.

• Orbitals: Specific regions in sub-shells where electrons are found.

  • s: 1 orbital (2 electrons), p: 3 orbitals (6 electrons), d: 5 orbitals (10 electrons), f: 7 orbitals (14 electrons).

What is the principal quantum number (n) and how does it relate to shells?

• The principal quantum number (n) indicates the energy level of a shell.

  • Lower n = closer to the nucleus and lower energy.

  • Higher n = farther from the nucleus and higher energy.

• It also determines the maximum number of electrons in a shell:

  • n = 1 (2 electrons), n = 2 (8 electrons), n = 3 (18 electrons), n = 4 (32 electrons).

What is the ground state of an atom?

• The ground state is the most stable electronic configuration with the lowest energy.

• Electrons fill sub-shells in order of increasing energy, starting with the 1s orbital.

• Example: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p.

How many orbitals and electrons are in each type of sub-shell (s, p, d)?

• s sub-shell: 1 orbital, 2 electrons.

• p sub-shell: 3 orbitals, 6 electrons.

• d sub-shell: 5 orbitals, 10 electrons.

• f sub-shell (not at AS level): 7 orbitals, 14 electrons.

What is the order of increasing energy in sub-shells?

• Energy generally increases as follows: s < p < d < f.

• The 4s sub-shell is lower in energy than the 3d sub-shell, so it is filled first.

• Example: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p.

How do you describe an atom's electronic configuration?

• It lists the number of electrons in each shell and sub-shell.

• Example for potassium (19 electrons): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.

• Box notation uses arrows in boxes to show individual electron spin.

What determines the energy of electrons and inter-electron repulsion?

• Energy of electrons: Electrons fill the orbitals in order of increasing energy (e.g., 1s < 2s < 2p < 3s < 3p < 4s < 3d).

• Inter-electron repulsion:

  • Electrons have spin, either clockwise or anticlockwise, and repulsion occurs between electrons with similar spin (spin-pair repulsion).

  • To minimize repulsion, electrons occupy separate orbitals within the same sub-shell first. Pairing occurs only when no empty orbitals are available.

Shape of s Orbitals

• The s orbitals are spherical in shape

• The size of the s orbitals increases with increasing shell number

  • E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1)

Shape of p Orbitals

• The p orbitals are dumbbell-shaped

• Every shell has three p orbitals except for the first one (n = 1)

• The p orbitals occupy the x, y and z-axis and point at right angles to each other so are oriented perpendicular to one another

• The lobes of the p orbitals become larger and longer with increasing shell number

How can you determine the electronic configuration of atoms and ions using full and shorthand notations?

• Full notation: Lists all electrons from 1s onwards, e.g., Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s².

• Shorthand notation: Uses the nearest preceding noble gas to simplify, e.g., Fe: [Ar] 3d⁶ 4s².

• Ions:

  • Negative ions: Add electrons to the outer sub-shell.

  • Positive ions: Remove electrons from the outer sub-shell. Transition metals lose electrons from the 4s sub-shell first, not the 3d sub-shell.

What is electrons in boxes notation, and how is it used?

• Electrons in boxes: Represents orbitals as boxes arranged by increasing energy, with electrons shown as arrows indicating spin.

• Example: For Fe ([Ar] 3d⁶ 4s²):

  • 4s: ↑↓

  • 3d: ↑↓ ↑↓ ↑↓

What is a free radical, and how is it represented?

• Definition: A species with one or more unpaired electrons.

• Representation: Unpaired electron shown as a dot. Example: Chlorine radical (1s² 2s² 2p⁶ 3s² 3p⁵).

  • Two p orbitals have paired electrons, and one p orbital has the unpaired electron.

• Formation: Free radicals are formed via homolytic fission, where a covalent bond splits evenly, sharing electrons between two atoms.

What is the first ionisation energy (IE1), and how is it defined?

• Definition: First ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions under standard conditions (298 K, 101 kPa).

• Equation for IE1 (Example – Calcium):

  • Ca (g)→Ca+(g)+e⁻

  • IE1=+590 kJ mol⁻¹

• Key Points:

  • Ionisation energy is always positive as energy is required to overcome the attraction between the electron and the positive nucleus (endothermic process).

  • Units: kJ mol⁻¹

How do you write equations for second and subsequent ionisation energies?

• Second Ionisation Energy (IE2): Energy required to remove one mole of electrons from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.

  • Equation: Ca+(g)→Ca²⁺(g)+e⁻

  • IE2=+1145.4 kJ mol⁻¹

• Subsequent Ionisation Energies: Follows the same pattern, with increasing energy required for each successive electron removal due to the increasing positive charge on the ion.

What are the trends in ionisation energies across a period and down a group?

• Across a Period (Left to Right):

  • Increase in IE1 due to:

    • Increasing nuclear charge: Greater attraction between nucleus and electrons.

    • Decreasing atomic radius: Outer shell is closer to the nucleus.

    • Constant shielding: Electrons are added to the same shell.

  • Exceptions: Slight decreases due to sub-shell effects (e.g., spin-pair repulsion in Oxygen vs. Nitrogen).

• Down a Group (Top to Bottom):

  • Decrease in IE1 due to:

    • Increased atomic radius: Outer electrons are farther from the nucleus.

    • Increased shielding: Inner electrons reduce nuclear attraction.

    • These factors outweigh the increasing nuclear charge.

What causes variations in successive ionisation energies of an element?

• Successive Ionisation Energies Increase:

  • Removing electrons from a positive ion requires more energy.

  • Less shielding and a higher proton-to-electron ratio increase nuclear attraction.

• Sharp Increase: Occurs when electrons are removed from an inner, more stable shell closer to the nucleus.

Why are ionisation energies due to the attraction between the nucleus and the outer electron?

• Outer electrons are attracted by the positive charge of the nucleus.

• Factors affecting nuclear attraction:

  • Nuclear charge: More protons = stronger attraction.

  • Atomic radius: Larger distance = weaker attraction.

  • Shielding: Inner electrons reduce the effective nuclear charge felt by outer electrons.

What factors influence the ionisation energy of an element?

• Nuclear Charge: More protons = stronger attraction to outer electrons = higher IE.

• Atomic/Ionic Radius: Greater distance from nucleus = weaker attraction = lower IE.

• Shielding by Inner Shells/Sub-shells: More inner electrons = greater shielding = lower IE.

• Spin-Pair Repulsion: Electrons in the same orbital repel each other, making it easier to remove one = lower IE.

How can successive ionisation energy data be used to deduce the electronic configuration of elements?

• Successive Ionisation Energies:

  • Successive ionisation energy values increase as more electrons are removed.

  • Large increases in ionisation energy occur when removing an electron from a full shell (closer to the nucleus).

• Deducing Electronic Configuration:

  • Identify where the sharp increase in energy occurs.

  • The number of electrons removed before the sharp increase corresponds to the number of electrons in the outer shell.

• Example (Calcium):

  • Successive ionisation energies: 590, 1145, 4940 kJ mol⁻¹.

  • Sharp increase after removing 2 electrons indicates the element has 2 outer electrons → group 2 element.

  • Electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².

How can successive ionisation energy data be used to deduce the position of an element in the Periodic Table?

• Outer Electrons & Group Number:

  • The number of electrons in the outer shell (found from the sharp increase) determines the element's group.

  • Example: Two outer electrons → group 2.

• Number of Electron Shells & Period Number:

  • The number of occupied shells corresponds to the period.

  • Successive ionisation energy data can show the energy levels of electrons, indicating the number of shells.

• Example:

  • Successive ionisation data: 800, 2427, 3660, 25000 kJ mol⁻¹.

  • Sharp increase after 3 electrons → group 3 (outer configuration: 3 electrons).

  • Ionisation energies indicate 3 shells → period 3.

  • Element: Aluminium (Al).

What factors influence successive ionisation energy data?

• Nuclear Charge:

  • Higher proton number → stronger attraction to electrons → higher ionisation energy.

• Shielding:

  • Inner electrons reduce the effective nuclear charge felt by outer electrons → lower ionisation energy.

• Atomic/Ionic Radius:

  • Larger radius → greater distance between nucleus and outer electron → lower ionisation energy.

• Sub-Shell & Spin-Pair Repulsion:

  • Electrons in the same orbital repel each other → lower ionisation energy due to repulsion.