Revision:Periodicity - 03

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3.1 The Periodic Table

3.1.1

Elements increase in atomic number across each period, and down each group. The history is boring and pointless (like all history) - ignore it.

3.1.2

Group - the columns going down.

Period - the rows going across.

3.1.3

Group - number of valence electrons in the atom.

Period - number of main electron shells - s, p , d and f blocks as described above.

3.2 Physical Properties

3.2.1

Li --> Cs (down the alkali metals)

Atomic radius increases due to increased electron shielding.

Ionic radius increases due to increased electron shielding.

Ionisation energy decreases due to increased electron shielding.

Melting/boiling point decreases due to increased electron shielding --> decreased forces.

Electronegativity decreases due to increased shielding --> decreased attraction for outer electrons.

F --> I (Down the halogens):

Atomic radius increases due to increased electron shielding.

Ionic radius increases due to increased electron shielding.

Ionisation energy decreases due to increased electron shielding.

Melting/boiling point increases due to increased number of electrons->increased london dispersion forces.

Electronegativity decreases due to increased shielding -> decreased attraction for outer electrons.

Na --> Ar (across period 3):

Atomic radius decreases due to increased nuclear charge --> greater attraction for electrons.

Ionic radius decreases Na --> Al (due to increased nuclear charge) jumps Al --> Si (due to reversal of ionisation direction...increased electron-electron repulsion) decreases Si --> Ar (due to increased nuclear charge).

Ionisation energy increases due to increased nuclear charge.

Melting/boiling point increases Na --> Si (due to stronger metallic bonding - more delocalized electrons then network covalent) drops Si-P (due to network->molecular covalent) increases P --> S (due to increased LDF between molecules ie P4, S8). Drops to Cl, due to smaller molecules (Cl2) decreases to Ar (individual atoms --> fewer electrons --> smaller LDF).

Electronegativity increases due to increased nuclear charge --> greater attraction for electrons.

3.3 Chemical Properties

3.3.1

Reactions of elements in the same group are similar because they have identical outer shells (ie same number of valence electrons). Generalized reactions follow :

Alkali metal (group 1) with water:

Alkali metal (group 1) with Halogen:

(Na acts as a reducing agent - is oxidized, Cl₂ is reduced)

Halogen with water:

(Exception F₂ is such a strong oxidizer : )

Halogen + Halide ion

Halide ion with Silver ion

(a white precipitate)

(a cream precipitate)

(a yellow precipitate)

3.3.2

Elements on the left are metallic, right are non-metals, Al is a metalloid (semi-metal).

Oxides : Non-metals --> Acidic oxides , Metals --> Basic oxides, Metalloids --> Amphoteric (both acidic & basic) oxides.

Oxide	Adding H2O	Adding HCl	Adding NaOH	Nature
Na2O			No reaction	Basic Oxide
MgO			No reaction	Basic Oxide
Al2O3	Insoluble			Amphoteric Oxide
SiO2	Insoluble	No reaction		Acidic Oxide
P4O10 (or P4O6)		No reaction		Acidic Oxide
SO3 (or SO2)		No reaction		Acidic Oxide
Cl2O7		No reaction		Acidic Oxide

Halides (assuming Cl - could replace with Br, I, F etc): Ionic Chlorides --> dissolved in H₂O with little reaction, Covalent Chlorides --> dissolve + react to form HCl.

This isn't required....not like it's hard

(Exception : F2 is such a strong oxidizer: )

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