periodicity Flashcards
(38 cards)
Define the terms “period” and “group” in the context of the periodic table.
Periods: Horizontal rows. All elements in a period have the same number of electron shells.
Groups: Vertical columns. All elements in a group have the same number of outer shell electrons, giving them similar chemical properties.
How is the periodic table arranged?
The periodic table is arranged by atomic (proton) number.
Describe how the periodic table is divided into blocks.
The periodic table is divided into:
s-block: Groups 1–2 + Helium.
p-block: Groups 3–0 (13–18).
d-block: Transition metals (middle section).
f-block: Lanthanides and actinides (bottom).
How is an element’s block determined?
An element’s block is determined by its proton number and its outermost sub-shell being filled:
s-block: Last electron enters an s sub-shell
p-block: Last electron enters a p sub-shell
d-block: Last electron enters a d sub-shell
f-block: Last electron enters an f sub-shell
Use phosphorus (Group 5, Period 3) as an example to explain how electron configuration can be deduced from the periodic table.
You can deduce an atom’s electron configuration using its period and group.
E.g. Phosphorus (P), Group 5, Period 3:
Configuration: 1s² 2s² 2p⁶ 3s² 3p³
Define the term “periodicity” in chemistry.
Periodicity is defined as the Trends in physical and chemical properties of elements across a period (left to right).
Describe the trend in atomic radius across a period and explain the reason for it.
Atomic radius decreases across a period.
Because Proton number increases, increases nuclear charge and electrons get pulled closer.
Shielding remains constant, as electrons are added to the same shell.
Define first ionisation energy.
First ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous 1⁺ ions.
State and explain the trend in first ionisation energy across Period 3 (Na → Ar).
The trend generally increases.
Reason: because Proton number increases causing stronger attraction to outer electrons as a result more energy is needed to remove them.
What are the exceptions to the general ionisation energy trend across a period and why?
• There is a drop from Mg to Al because Al’s outer electron is in 3p, which is higher in energy and shielded by 3s², making it easier to remove.
• There is a drop from P to S because in S, pairing in the 3p orbital causes electron repulsion, making it easier to remove one of the paired electrons.
Describe and explain the trend in melting points across Period 3 from sodium to aluminium.
Melting point increases.
Reason: because there is Stronger metallic bonding due to:
Increase positive charge of ions (Na⁺, Mg²⁺, Al³⁺),
Increase in number of delocalised electrons,
Decrease in atomic radius.
Explain why silicon has a very high melting point.
Silicon:
Very high melting point due to
Macromolecular structure (giant covalent), tetrahedral.
And
Strong covalent bonds throughout therefore lots of energy needed to break.
Describe the melting point trend for phosphorus (P₄), sulfur (S₈), and chlorine (Cl₂), and explain why sulfur has the highest among them.
These are molecular substances.
Held by van der Waals forces:
Weak intermolecular forces → low melting points.
S₈ > P₄ > Cl₂ due to increasing molecule size causing stronger van der Waals.
Why does argon have a very low melting point?
Argon is monatomic, with the weakest van der Waals forces → very low melting point.
What should students memorise about the periodic table for the exam?
You won’t be given a block-labelled periodic table in the exam — memorise block positions.
What notation is used for electron configuration and what is shielding?
Understand the 1s² 2s² 2p⁶… notation for electron configuration.
Shielding is when Inner electrons reduce nuclear attraction felt by outer electrons (but remains similar across a period).
Describe how the modern periodic table is arranged.
Arranged by increasing atomic (proton) number.
Divided into periods (rows) and groups (columns).
Explain the difference between the periodic table developed by Mendeleev and the modern periodic table.
Developed originally by Mendeleev, modern version is ordered by proton number, not relative atomic mass.
What do all elements in the same period have in common?
Elements in the same period have the same number of electron shells.
Period number = number of occupied shells.
What do all elements in the same group have in common?
Elements in the same group have the same number of electrons in their outer shell.
Group number = number of outer electrons.
Therefore, elements in the same group have similar chemical properties.
Describe the periodic table in terms of its blocks.
Periodic table is divided into:
• s-block: group 1 & 2 (+ He); outermost electrons in s orbital.
• p-block: groups 3–0 (13–18); outer electrons in p orbital.
• d-block: transition metals; outer electrons in d orbital.
• f-block: lanthanides & actinides; outer electrons in f orbital.
How can the electron configuration of an element be determined from the periodic table?
Use the position in the table (period, group, block) to determine configuration.
Fill orbitals in order:
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s →
Give the electron configuration of phosphorus.
Phosphorus (P) – Period 3, Group 5:
Electron configuration: 1s² 2s² 2p⁶ 3s² 3p³
Give the electron configuration of vanadium.
Vanadium (V) – Period 4, d-block:
Electron configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d³
(or [Ar] 4s² 3d³)
