FUNCHEM.3 The Periodic Table: Trends and Properties
Fundamentals of Medicinal and Pharmaceutical Chemistry: The Periodic Table
Recommended Reading
General Chemistry - The Essential Concepts by Chang and Goldsby 7e
Sections 2.4, 8.3, 8.4, 8.5, 9.5
Learning Outcomes
Upon completion of this material, you should be able to:
Recognise the arrangement of elements in the periodic table according to increasing atomic number.
Differentiate between periods and groups, metals, non-metals and metalloids, alkali metals, alkaline earth metals, transition metals, halogens, inert gases, lanthanides and actinides in the periodic table.
Define 'effective nuclear charge'.
Define 'atomic radius', 'ionic radius', 'ionisation energy', 'electron affinity' and 'electronegativity' and describe how each varies on going across a period and down a group in the Periodic Table.
The Periodic Table: The Alphabet of Chemistry
Nature simplifies the study of elements by grouping 118 elements into families, displayed by the Periodic Table.
Historical Context: Mendeleev's Periodic Table
Dmitri Mendeleev (1834-1907) published the first arrangement of elements in 1869.
He organised elements in order of increasing atomic mass.
He had no knowledge of atomic structure or the reasons for elemental behaviour.
He successfully predicted the existence and properties of several undiscovered elements by leaving gaps in his table.
The Modern Periodic Table
Elements are arranged according to increasing atomic number.
Groups: Vertical columns exhibiting similar chemical properties.
Periods: Horizontal rows.
Elements are represented by chemical symbols.
Classification of Elements
Metals: Good conductors of heat and electricity.
Non-metals: Usually poor conductors of heat and electricity.
Metalloids: (Not explicitly defined but implied by the differentiation)
Elements show great variety in chemical properties, but similarities exist within groups.
Reactivity: Group I (alkali metals) and Group VII (halogens) elements are generally the most reactive.
Valence Electrons
Valence electrons are the outermost electrons in an atom.
They are crucial for chemical reactions, participating in sharing and exchange between atoms.
The attraction between the positive nucleus and the negative valence electrons depends on:
Number of protons in the nucleus (actual nuclear charge, Z).
Shielding effect of other electrons closer to the nucleus.
Distance from the nucleus to the outermost shell.
Effective Nuclear Charge, Z_{eff}
Definition: The nuclear charge felt by an electron when both the actual nuclear charge (Z) and the repulsive (shielding) effects (\sigma) of other electrons are taken into account.
Formula: Z_{eff} = Z - \sigma
\sigma (sigma) is the shielding constant (or screening constant).
\sigma is always greater than zero but smaller than Z (0 < \sigma < Z).
Trends:
Across a Period (left to right):
Z_{eff} increases.
The actual nuclear charge, Z, increases.
Added electrons are valence electrons and do not shield each other effectively.
This results in a greater effective nuclear charge felt by the valence electrons.
Down a Group (top to bottom):
Z_{eff} increases as calculated per Slater's rules, which account for the effective shielding of electrons in each orbital "shell".
Atomic Radius
Definition: Half the distance between the nuclei of two adjacent atoms of the same element.
Trends:
Across a Period (left to right):
Atomic number (Z) increases, leading to more protons in the nucleus.
Electrons are added to the same energy level.
Effective nuclear charge (Z_{eff}) increases, causing a stronger attraction between the nucleus and valence electrons.
Atomic radius decreases.
Down a Group (top to bottom):
Electrons enter another energy level (a new electron shell).
This results in a greater distance between the nucleus and the outermost shell.
Atomic radius increases.
Ionic Radius
Definition: The radius of a cation or an anion.
Formation of an Anion (X⁻):
When an atom gains one or more electrons.
The nuclear charge remains the same.
The additional electron(s) increase electron-electron repulsion.
The size increases, meaning the ionic radius increases (e.g., Cl atom (99 ext{ pm}) vs. Cl⁻ ion (181 ext{ pm})).
Formation of a Cation (X⁺):
When an atom loses one or more electrons.
The nuclear charge remains the same.
The loss of electron(s) reduces electron-electron repulsion.
The size decreases, meaning the ionic radius decreases (e.g., Na atom (186 ext{ pm}) vs. Na⁺ ion (99 ext{ pm})).
Ionic Radii of Isoelectronic Species (Same Number of Electrons):
Cations: Compare Na⁺ (11 protons, 10 electrons, 99 ext{ pm}) and Mg²⁺ (12 protons, 10 electrons, 65 ext{ pm}).
Mg²⁺ has a bigger nuclear charge (12 vs. 11 protons).
This results in a stronger attraction and a smaller size.
Radius of a dipositive ion (X^{2+}) < Radius of a unipositive ion (X^{+}).
Anions: Compare O²⁻ (8 protons, 10 electrons) and F⁻ (9 protons, 10 electrons).
F⁻ has a bigger nuclear charge (9 vs. 8 protons).
This results in a stronger attraction and a smaller size for F⁻ compared to O²⁻.
Radius of a dinegative ion (X^{2-}) > Radius of a uninegative ion (X^{-}).
Practical Implications: Ionic size is crucial for biological processes, such as ions moving through specific channels in cell membranes (e.g., channels allowing Na⁺ but not larger K⁺).
Ionisation Energy (IE)
Definition: The minimum energy (in kJ/mol) required to remove an electron from a gaseous atom in its ground state.
Energy is needed to overcome the attractive force between the positive nucleus and the negative electrons.
First Ionisation Energy (IE_1): The energy required to remove the first electron.
M(g) \rightarrow M^+(g) + e^-
Second Ionisation Energy (IE_2): The energy required to remove the second electron from a unipositive ion.
M^+(g) \rightarrow M^{2+}(g) + e^-
Third Ionisation Energy (IE_3): The energy required to remove the third electron from a dipositive ion.
M^{2+}(g) \rightarrow M^{3+}(g) + e^-
Trends:
Across a Period (left to right):
Nuclear charge (Z) increases.
Electrons are added to the same energy level.
Effective nuclear charge (Z_{eff}) increases, leading to a stronger attraction to the outermost electron.
Ionisation energy increases.
Down a Group (top to bottom):
Electrons enter another energy level, increasing the distance between the nucleus and the outermost shell.
The increased distance weakens the nuclear attraction to the outermost electron.
Ionisation energy decreases.
Exceptions to Trends:
Beryllium (Be) vs. Boron (B): IE_1 of Be (1s^2 2s^2) is higher than B (1s^2 2s^2 2p^1) because removing a 2p electron from B is easier than removing an electron from the stable, filled 2s orbital of Be.
Nitrogen (N) vs. Oxygen (O): IE_1 of N (1s^2 2s^2 2p^3; half-filled, stable p-orbital) is higher than O (1s^2 2s^2 2p^4) because removing a paired electron from Oxygen (which alleviates electron-electron repulsion and leads to a more stable half-filled 2p configuration) is easier than removing an electron from the stable half-filled 2p orbital of Nitrogen.
Overall: Alkali metals have the lowest ionisation energies, while noble gases have the highest.
Electron Affinity (EA)
Definition: The negative of the energy change (\Delta H) that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
X(g) + e^- \rightarrow X^-(g)
Sign Convention:
When energy is released (exothermic process, \Delta H is negative), the electron affinity is considered positive.
Example: F(g) + e⁻
ightarrow F⁻(g), \Delta H = -328 ext{ kJ mol}^{-1}, Electron affinity is +328 ext{ kJ mol}^{-1}.
A more positive electron affinity indicates a greater tendency for an atom to accept an electron.
Electronegativity
Definition: The ability of an atom to attract towards itself the electrons in a chemical bond.
Examples:
In an H:H bond (like H_2), electrons are shared equally.
In an H:F bond (like HF), Fluorine (F) is more electronegative than Hydrogen (H).
F pulls the shared electrons closer to itself.
This creates a slight negative charge on F (\delta^-) and a slight positive charge on H (\delta^+).
This separation of charge creates a dipole moment towards F.
Why Fluorine (F) is the Most Electronegative Element:
It has a high nuclear charge (9 protons).
It exhibits little shielding due to its small atomic radius, meaning its outer electrons are close to the nucleus.
It has both a high electron affinity (readily picks up electrons) and a high ionisation energy (does not easily lose electrons), making it extremely effective at attracting shared electrons.
Electronegativity Order (Top Four): F > O > N > C (Mnemonic: "Front Office Never Closes")
Trends: (generally)
Across a Period (left to right): Electronegativity increases.
Down a Group (top to bottom): Electronegativity decreases.
Summary of Periodic Trends
Property | Across a Period (Left to Right) | Down a Group (Top to Bottom) |
|---|---|---|
Atomic Radius | Decreases | Increases |
First I.E. | Increases | Decreases |
Electronegativity | Increases | Decreases |
Upon completion of this material, you should be able to:
Recognise the arrangement of elements in the periodic table according to increasing atomic number.
Elements in the modern periodic table are arranged according to increasing atomic number.
Differentiate between periods and groups, metals, non-metals and metalloids, alkali metals, alkaline earth metals, transition metals, halogens, inert gases, lanthanides and actinides in the periodic table.
Groups: Vertical columns of elements that exhibit similar chemical properties.
Periods: Horizontal rows of elements.
Metals: Good conductors of heat and electricity.
Non-metals: Usually poor conductors of heat and electricity.
Metalloids: Although not explicitly defined in the provided material, metalloids are generally understood to have properties intermediate between metals and non-metals.
Alkali metals: Group I elements, generally the most reactive elements.
Halogens: Group VII elements, generally the most reactive elements.
(Alkaline earth metals, transition metals, inert gases, lanthanides, and actinides are not explicitly defined or differentiated in the provided text, beyond the mention of groups and general classifications.)
Define 'effective nuclear charge'.
Effective nuclear charge (Z {eff}) is the nuclear charge felt by an electron when both the actual nuclear charge (Z) and the repulsive (shielding) effects (
\sigma
) of other electrons are taken into account. It is calculated as Z
{eff} = Z - \sigma.
Define 'atomic radius', 'ionic radius', 'ionisation energy', 'electron affinity' and 'electronegativity' and describe how each varies on going across a period and down a group in the Periodic Table.
Atomic Radius: Half the distance between the nuclei of two adjacent atoms of the same element.
Across a Period (left to right): Decreases (due to increasing Z
{eff}
).Down a Group (top to bottom): Increases (due to electrons entering a new energy level).
Ionic Radius: The radius of a cation or an anion.
Anion formation (gain electron): Size increases due to increased electron-electron repulsion.
Cation formation (lose electron): Size decreases due to reduced electron-electron repulsion.
Isoelectronic species: Radius decreases with increasing nuclear charge (e.g., X^{2+} < X^{+} and X^{2-} > X^{-}).
Ionisation Energy (IE): The minimum energy required to remove an electron from a gaseous atom in its ground state.
Across a Period (left to right): Increases (due to increasing Z
{eff}
and stronger attraction to the outermost electron).Down a Group (top to bottom): Decreases (due to increased distance between the nucleus and the outermost shell).
Electron Affinity (EA): The negative of the energy change ( \Delta H ) that occurs when an electron is accepted by an atom in the gaseous state to form an anion.
A more positive electron affinity indicates a greater tendency for an atom to accept an electron. Trends are not explicitly detailed in the summary for this property.
Electronegativity: The ability of an atom to attract towards itself the electrons in a chemical bond.
Across a Period (left to right): Increases.
Down a Group (top to bottom): Decreases.