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A-level Chemistry: 3.2.1 Periodicity

Chemistry17 CardsCreated 2 months ago

These flashcards focus on the organization of the periodic table, including how it is arranged, periodic trends like atomic radius, and how to determine electron configurations of elements using their position on the table.

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The periodic table is arranged by ____ _____

Atomic number

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Key Terms

Term
Definition

The periodic table is arranged by ____ _____

Atomic number

Name the blocks

s-block, p-block, d-block, f-block

Elements within a period have same number of _____ ___

electron shells

Using the periodic table, write the electron configuration for phosphorus

[Ne] 3s² 3p³

Using the periodic table, write the electron configuration for cobalt

[Ar] 3d⁷ 4s²

Atomic Radius _____ across a Period

decreases

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A-level Chemistry: 3.2.1 Periodicity
TermDefinition

The periodic table is arranged by ____ _____

Atomic number

Name the blocks

s-block, p-block, d-block, f-block

Elements within a period have same number of _____ ___

electron shells

Using the periodic table, write the electron configuration for phosphorus

[Ne] 3s² 3p³

Using the periodic table, write the electron configuration for cobalt

[Ar] 3d⁷ 4s²

Atomic Radius _____ across a Period

decreases

Why does atomic radius decreases across a period?

  • As number of proton increases, positive charge of nucleus increases

  • Increased attraction between nucleus and electrons

  • Similar shielding in each period

    • As electrons added to same outer shell

Why don’t the extra electrons that element gains across period provide extra shielding?

  • Extra electrons are added to outer energy level

  • (Shielding is mainly provided by electrons in inner shells)

As you go across period 3, from sodium to magnesium to aluminium, why does melting and boiling points increase?

  • ∵ metal-metal bonds get stronger

  • Get stronger ∵ metal ions have increasing positive charge = increasing number of delocalised electrons and decreasing radius

Why does silicon have such a high melting point?

  • Silicon is macromolecular (tetrahedral structure) = many strong covalent bonds

  • Lot of energy needed to break these bonds

Why does phosphorus, sulfur and chlorine have low melting points?

  • They’re molecular substances so melting points depend upon strength of van der Waals forces between molecules

  • Van der Waals forces are weak and easily overcome

Why does sulfur have a higher melting point than phosphorus or chlorine? (2)


  • Sulfur is the biggest molecule (largest Mr)

  • Stronger van der Waals forces between molecules (they’re stronger with larger Mr)

Why does argon have a very low melting point?

∵ exists as individual atoms = weak van der Waals forces

Ionisation energy generally ______ across a period

increases

Why does ionisation energy generally increases across a period (4)

  • As number of proton increases, positive charge of nucleus increases

  • Increased attraction between nucleus and electrons

  • Similar shielding in each period

    • As electrons added to same outer shell

  • More energy needed to remove electron

Predict the element in Period 3 that has the highest second ionisation energy. Give a reason for your answer. (2)

  • Sodium / Na

  • Electron (removed) from the 2nd shell / 2p (orbital)

Explain why the ionisation energy of every element is endothermic (1)

Energy needed to overcome attraction between nucleus and electron