Periodicity Flashcards

Periodicity

A-Level AQA Chemistry

This pack of flashcards covers the concepts of periodicity, electron configuration, and chemical bonding in elements. It explains atomic radius trends, ionization energy, metallic bonding strength, and variations in bonding types across different elements, including the transition from metallic to covalent and van der Waals forces.

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Describe periodicity.	Periodicity is the repeating pattern of physical or chemical properties going across the periods.
Define s, p, and d blocks.	Elements are classified as S, p or d block, according to which orbitals the highest energy electrons are in.
How does atomic radius change across a period?	Atomic radii decrease from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding.
What is the trend in atomic radius in period 2?	The atomic radii decrease from left to right in period 2.
Describe the trend in 1st ionization energy.	There is a general trend across to increase in 1st ionization energy. This is due to the increasing number of protons as the electrons are being added to the same shell.
What is the reason for the small drop in ionization energy between phosphorous and sulfur?	Sulfur's outer electron is being paired up with another electron in the same 3p orbital.
Explain why the second electron is easier to remove when added to an orbital.	When the second electron is added to an orbital, there is a slight repulsion between the two negatively charged electrons, making the second electron easier to remove.
What is the relationship between the number of electrons in the outer shell and the strength of metallic bonding?	The more electrons there are in the outer shell that are released to the sea of electrons, the stronger the metallic bonding.
Define metallic bonding.	Metallic bonding is a type of chemical bonding that occurs between metal atoms, where the outer electrons are delocalized and form a 'sea' of electrons that surround the positively charged metal ions.
Describe the relationship between the size of an ion and its bonding strength.	A smaller sized ion with a greater positive charge makes the bonding stronger.
What is the general trend in energy needed to break bonds?	Higher energy is needed to break bonds.
What type of bonding does Na, Mg, Al, Si, P, S, Cl, and Ar exhibit?	Na, Mg, Al, and Si exhibit metallic bonding, P and S exhibit macromolecular bonding, Cl2(g), S8(s), and P4(s) exhibit simple molecular bonding, and Ar exhibits monoatomic weak van der Wa bonding.
Why does S8 have a higher melting point than P4?	S8 has more electrons, resulting in stronger van der Waals forces between molecules.

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