Periodic Table: Groups and Periods

The Periodic Table

• Mendeleev introduced the modern periodic table.
• Elements are arranged by atomic number and are categorized into groups, periods, and blocks.
• Group elements exhibit similar chemical behaviors.
• In most instances, the number of valence electrons matches the group number.
• Elements in the same group display similar physical and chemical properties.
• Alkali metals (Li, Na, K, Rb, Cs) and halogens (F, Cl, Br, I) serve as examples of element families.
• Group 4 or 14 elements (C, Si) do not form a family due to differing n-numbers.

Groups

• Vertical arrangements of elements.
• Elements in a group have the same number of valence electrons and similar atomic configurations.
• Li= [He] 2s¹, ₁₁Na = [Ne] 3s¹, and 19 K= [Ar] 4s1 are examples of elements with one valence electron in their s-orbital, having one electron outside a noble gas configuration.

Periods

• Horizontal arrangements of elements.
• Atomic number increases by one from one element to the next.
• There are 7 periods, the length of which varies.
• The shortest period has 2 elements and the longest has 32.
• The period corresponds to the main energy level that the valence electron occupies.
• P1 (1s) contains 2 elements, P2 (2s, 2p) contains 8 elements, P3 (3s, 3p) contains 8 elements.
• P4 (4s, 3d, 4p) contains 18 elements, P5 (5s, 4d, 5p) contains 18 elements, and P6 (6s, 4f, 5d, 6p) contains 32 elements.

Blocks

• Groups of elements share the same valence electron orbital.
• Elements in groups G1 and G2 are s-block elements.
• Elements from group 3 to group 0 are p-block elements since their valence electrons are found in the p-orbital.
• D-block elements (transition elements) possess valence electrons that fill d-orbitals, such as the 1st transition series (4th period: Sc - Zn).
• F block elements, consist of valence electrons filling the f-orbital, including lanthanides [Lanthanum, La (57) - Lutetium, Lu (71)] and actinides [Ac (Actinium) 89 - Lr (Lawrencium) 103].

Periodic Properties

• Physical and chemical attributes of atoms that are recurring.
• Examples include atomic/ionic radius, ionization energy, electron affinity, and electronegativity.
• Outer electrons involved in bond formation make up Valence electrons.
• Inner electrons are Core electrons.
• The net attraction of protons in the nucleus for valence electrons is Effective nuclear charge.
• Shielding effect is the decreased attraction of protons for valence electrons by the core electrons.
• Effective nuclear charge increases across a period.
• Shielding effect increases down a group.

Atomic Radius

• The size of an atom relies on the strength of the nucleus' proton attraction to valence electrons.
• The attraction for valence electrons is greater as you move across a period and the size of the atom gets smaller.
• Atomic size increases from halogen to rare gases due to repulsion among electrons.
• Denoted as 1/2d=r

Ionic Radius

• Cations are smaller than their parent atoms.
• This is because they have fewer electrons than protons, which leads to more effective nuclear charge.
• Cationic radius decreases across a period but grows down a group.
• Anions are larger than their parent atoms because there are more electrons than protons, which reduces the effective nuclear charge.

Ionization Energy (Ei)

• It is the energy to remove the most loosely bound electrons from a gaseous atom.
• Determining factors include atomic radius, shielding effect, stability of electronic configuration, and effective nuclear charge.
• Ionization energy decreases down a group; atomic radius and screening rise as penetration falls.
• Ionization energy increases across a period.

Ionization Energy Stability Factors

• The 2p electrons are at a higher energy level than 2s electrons
• The higher the energy level, the easier it is to remove electrons from 2p.
• The s-electrons have greater penetrative power than the p-electrons.

Electron Affinity

• Energy change when an electron is added to a gaseous atom.
• Adding the second electron is a positive value, like O(g) + e- → O2-(g), Eea= +798 KJ/mol.
• The 2nd electron in the oxygen was added against the repulsive force of negatively charged oxygen, requiring energy input.
• Decrease in electron affinity down a group.
• Decrease in attraction for added electrons, increase in atomic size, and decrease in effective nuclear charge.
• Discrepancy: Fluorine (F) has unexpectedly lower negative electron affinity due to its small size and strong repulsive field.
• Additional factors include atomic radius, the stability of the electronic configuration, and the charge carried by the anion.

Electronegativity

• This is the ability of an electron in a molecule to attract bonded/shared electrons to itself.
• Its main determinant is atomic radius.
• The lower the atomic radius the higher the electronegativity.
• Fluorine has the highest electronegativity, quantified at 4.0 on Pauling's scale.
• Francium is the lowest, measured at 0.7.
• Halogens are the most electronegative elements and alkali metals are the the least.
• Standard enthalpy of formation is the energy change when one mole of a compound is formed from elements in their standard states.

Learn about the periodic table, focusing on the arrangement of elements into groups and periods. Discover how elements within the same group share similar chemical properties and valence electron configurations. Explore examples of alkali metals, halogens, and the characteristics of elements in Group 4 or 14.