periodicity

what is periodicity? (1)

repeating trends of physical or chemical properties

what are periods in the periodic table? (1)

horizontal rows of elements in the periodic table

what are groups in the periodic table? (2)

• vertical columns in the periodic table

• where all the elements have the same number of electrons in their outermost principal energy level (PEL) and similar properties

where are the blocks of elements found? (4)

how does reactivity change in the s-block elements as we move down a group? (1)

elements get more reactive as we move down a group

how does reactivity change in non-metals as we move up a group? (1)

elements tend to get more reactive as we move up a group

how reactive are the transition metals in the d-block? (1)

usually unreactive

what type of structures do elements in groups 1, 2, and 3 have? (1)

elements in groups 1, 2, and 3 are all metals and have giant metallic structures

what type of structure does silicon (Si) in group 4 form? (1)

a macromolecular structure with 4 covalent bonds

what type of structures do elements in groups 5, 6, and 7 form? (1)

they are non-metals so form simple molecular structures

what is the structure of argon in Group 0, and why is it inert? (2)

• has a simple molecular structure

• with a full outer PEL of electrons (making it inert)

why do molecular structures have low melting and boiling points? (1)

they have weak Van der Waals forces between molecules that need to be broken

why do metallic structures have high melting and boiling points? (1)

due to strong electrostatic attraction between positive ions and delocalised electrons

why do macromolecular structures like silicon have very high melting and boiling points? (2)

due to strong covalent bonds which require a lot of energy to break

what type of bonding do Na, Mg, and Al possess? (1)

metallic bonding

how does the charge on the metal ion change across Na, Mg, and Al? (2)

the charge on the metal ion increases from 1+ to 3+

what happens to the number of delocalised electrons across Na, Mg, and Al? (1)

the number of delocalised electrons increases

how does the strength of metallic bonding change across Na, Mg, and Al? (2)

the strength of metallic bonding increases, making the metals harder to melt

why do Na, Mg, and Al have increasing melting and boiling points? (2)

Na, Mg, and Al have increasing melting and boiling points due to stronger metallic bonding

what type of structure does silicon (Si) have? (1)

macromolecular structure (similar to diamond)

how are silicon atoms bonded in its macromolecular structure? (2)

• each silicon atom is bonded to 4 others in a tetrahedral structure

• forming a giant 3D structure

why does silicon have a very high melting and boiling point? (2)

• silicon has a very high melting and boiling point due to strong covalent bonds

• which require a lot of energy to break

what types of structures do P, S, Cl, and Ar form? (1)

they form simple molecular structures

how do P, S, and Cl exist in nature compared to Ar? (2)

P, S, and Cl exist as simple molecules, while Ar exists as separate atoms

why do P, S, Cl, and Ar have low melting and boiling points? (2)

• they have weak Van der Waals forces between molecules

• requiring less energy to break and melt/boil the compounds

why does sulfur have a higher melting point than phosphorus, chlorine, or argon? (3)

• phosphorus exists as P₄ molecules, sulfur as S₈ molecules, chlorine as Cl₂ molecules, and argon as Ar atoms

• sulfur, being a larger molecule, has more Van der Waals forces between molecules

• more energy is required to break these forces, resulting in a higher melting point

what is the order of melting and boiling points for P, S, Cl, and Ar? (1)

S₈ > P₄ > Cl₂ > Ar

what happens to ionisation energy down a group, and why? (5)

• ionisation energy decreases

• the electron is removed from a higher principal energy level

• the electron is further from the nucleus

• there is more shielding

• weaker attraction between the nucleus and the outer electron means less energy is required to remove it

what happens to ionisation energy across a period, and why? (4)

• ionisation energy increases

• the number of protons increases

• shielding is constant, and atomic radius decreases

• stronger attraction between the nucleus and the outer electron means more energy is required to remove it

why is there an exception for Group 3 ionisation energy across a period? (3)

• ionisation energy is lower

• the electron is removed from a higher energy p sub-level

• less energy is required to remove the electron

why is there an exception for Group 6 ionisation energy across a period? (3)

• ionisation energy is lower

• there is a pair of electrons in a p orbital

• extra repulsion means less energy is required to remove the electron