3.1.1 - Periodicity Flashcards
(63 cards)
How are elements arranged in the periodic table?
Elements are arranged in increasing atomic number in the periodic table.
Why do elements in the same group have similar chemical properties?
The atoms of elements in a group have similar outer shell electron configurations, resulting in similar chemical properties
What are the elements in Period 2?
Li, Be, B, C, N, O, F, Ne
What are the elements in Period 3?
Na, Mg, Al, Si, S, Cl, Ar
How are elements classified into s, p, or d blocks?
According to the type of orbital their highest-energy electron occupies.
What does periodicity mean?
A repeating pattern across periods in properties like atomic radius, melting points, and ionisation energies.
What is the trend in atomic radius across a period?
It decreases due to increased nuclear charge attracting electrons more strongly in the same shell.
What is the trend in atomic radius down a group?
It increases due to more shells being added.
Why do atoms get smaller across a period despite more electrons?
Because the increased proton number pulls electrons in more strongly without additional shielding.
What is the definition of first ionisation energy?
The energy needed to remove one electron from each atom in one mole of gaseous atoms.
What is the equation for first ionisation energy?
H(g) → H⁺(g) + e⁻
What is the equation for second ionisation energy?
H⁺(g) → H2⁺(g) + e⁻
Why must the first ionisation energy equation always be written in the gas phase?
Because ionisation energies are defined for gaseous atoms, ensuring consistent comparison.
Why does it not matter if the atom normally forms a +1 ion or not in ionisation energy questions?
Ionisation energy is a theoretical concept based on removing one electron from a gaseous atom — not about real ion formation.
Why is it useful to study the pattern of ionisation energies across a period?
It provides insight into the electronic structure and shell arrangement.
What are the 3 main factors affecting ionisation energy?
- Nuclear charge (more protons = stronger attraction)
- Distance from nucleus
- Electron shielding from inner shells.
What is the trend in ionisation energies across the period?
- Nuclear charge increases
- Electron Shielding stays the same
- Atomic Radius decreases = stronger attraction between nucleus and electrons.
OVERALL: Energy needed to remove an electron increases.
What is the trend in ionisation energies down the group?
- Number of shells increases = distance increases, so weaker forces of attraction.
- More shells = increased shielding, weaker attraction
- Increase in nuclear charge is outweighed by these factors.
THEREFORE, first ionisation energies decrease.
How does the pattern of first ionisation energy support periodicity?
It shows a repeating trend across periods, reflecting similar changes in nuclear charge, shielding, and atomic radius.
Why is helium’s first ionisation energy the highest?
Its only electron is closest to the nucleus and experiences no shielding.
Why is Na’s first ionisation energy lower than Ne’s?
Na’s outer electron is in the 3s shell, further from the nucleus and more shielded.
Why is there a small drop in ionisation energy from Mg to Al?
Al’s outer electron is in a 3p orbital, which is higher in energy and more shielded than Mg’s 3s electrons.
Why is there a drop from P to S?
In S, the 3p orbital starts to pair electrons, causing repulsion that makes the second electron easier to remove.
Why do the explanations for Mg→Al and P→S drops differ from general trends?
Because they involve sub-shell changes and electron repulsion — not just nuclear attraction and shielding.
