Periodicity Flashcards
What is the periodic table?
A chart organizing elements based on increasing atomic number (Z) and similar chemical properties in vertical groups.
What do elements in the same group have in common?
Same number of valence electrons, leading to similar chemical properties.
What do elements in the same period have in common?
Same number of electron shells but increasing protons and valence electrons.
What are periodic table blocks?
s-block: Groups 1 & 2 (Alkali & Alkaline Earth metals); p-block: Groups 13-18 (Non-metals, Metalloids, Noble Gases); d-block: Transition metals (Groups 3-12); f-block: Lanthanides & Actinides (inner transition metals).
What is atomic radius?
Distance from the nucleus to the outermost electron shell.
Why does atomic radius decrease across a period?
Z_eff increases as protons pull electrons closer; shielding stays the same, so attraction strengthens; stronger electrostatic pull reduces size.
Why does atomic radius increase down a group?
More electron shells increase size; shielding increases, reducing nuclear attraction; Z_eff remains nearly constant.
Which element has the largest atomic radius?
Francium (Fr), due to low Z_eff and many electron shells.
Which element has the smallest atomic radius?
Helium (He), due to highest Z_eff for its period and no shielding.
What is ionisation energy?
The energy needed to remove an electron from an atom in the gas phase.
Why does ionisation energy increase across a period?
Z_eff increases, pulling electrons closer; smaller atomic radius makes electrons harder to remove.
Why does ionisation energy decrease down a group?
Larger atomic radius reduces nuclear attraction; more shielding weakens nuclear pull.
Why does ionisation energy drop from Group 2 to Group 13?
Group 2 has full s-orbitals (ns²), which are stable; Group 13 has one p-electron (ns² np¹), which is easier to remove.
Why does ionisation energy drop from Group 15 to Group 16?
Group 15 has a half-filled p-orbital (ns² np³), which is stable; Group 16 has a paired electron (ns² np⁴), increasing repulsion, making removal easier.
What is electronegativity?
An atom’s ability to attract bonding electrons.
Why does electronegativity increase across a period?
Z_eff increases, strengthening nuclear attraction; smaller atomic radius means electrons are closer to the nucleus.
Why does electronegativity decrease down a group?
Larger atomic radius reduces nuclear attraction; more shielding weakens the pull on bonding electrons.
Which element has the highest electronegativity?
Fluorine (F) = 3.98 (Pauling scale).
Which element has the lowest electronegativity?
Francium (Fr) = 0.7, due to large size & weak Z_eff.
How does metal reactivity change across a period?
Decreases, as Z_eff increases, making electron loss harder.
How does metal reactivity change down a group?
Increases, as larger atoms lose electrons more easily.
How does non-metal reactivity change across a period?
Increases, as non-metals attract electrons more easily.
How does non-metal reactivity change down a group?
Decreases, due to weaker attraction for electrons.
How do alkali metals react with water?
Violently, forming hydroxide + hydrogen gas.
How does alkali metal reactivity change down the group?
Increases, due to lower ionisation energy.
How do alkaline earth metals react with water?
Less violently, forming hydroxides + hydrogen gas.
What is electron configuration?
The arrangement of electrons in atomic orbitals (e.g., 1s² 2s² 2p⁶).
What is the Aufbau Principle?
Electrons fill the lowest energy orbitals first.
What is Hund’s Rule?
Electrons fill degenerate orbitals singly before pairing.
What is the Pauli Exclusion Principle?
No two electrons in an atom can have the same set of quantum numbers.
What is the significance of spdf notation?
It shows subshell arrangement: s (2), p (6), d (10), f (14).
Why do transition metals have variable oxidation states?
d-orbitals allow different electron losses, forming multiple ions.
Why are noble gases unreactive?
Full valence shells, making them stable.
Why does Group 1 have low ionisation energy?
Low Z_eff & large atomic size, making electron loss easy.
Why do halogens form -1 ions?
High electronegativity, easily gaining 1 electron.
Why do noble gases have no electronegativity values?
They don’t form bonds, so they don’t attract electrons.
What happens to atomic radius across a period?
Decreases.
What happens to atomic radius down a group?
Increases.
What happens to ionisation energy across a period?
Increases.
What happens to ionisation energy down a group?
Decreases.
What happens to electronegativity across a period?
Increases.
What happens to electronegativity down a group?
Decreases.
What happens to metallic character across a period?
Decreases.
What happens to metallic character down a group?
Increases.
