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periodicity Flashcards

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1
Q

Describe and explain the trend in atomic radius

across period 3

A

As you go across the period, atomic radius decreases:

● Number of protons in the nucleus/ nuclear charge increases.
● Number of electrons in the outer shell increases.
● Shielding remains the same.
● Nuclear attraction between the electrons and the nucleus
increases so electron shells are drawn closer to the nucleus,
decreasing the atomic radius.

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2
Q

Describe and explain the trend in ionic radius across

period 3

A

● From Na+ to Mg2+ to Al3+:
Ionic radius decreases because the number of electrons decreases so
there is greater attraction between outer shell electrons and the
nucleus meaning the electrons are drawn inwards.

● From P3- to S2- to Cl-:
Ionic radius increases because the number of electrons increases
which weakens the nuclear attraction meaning the electrons are not
drawn inwards as strongly.

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3
Q

Describe the trend in melting point

across period 3

A

● Melting point increases from sodium to silicon.
● There is a sharp decrease in melting point
between silicon and phosphorus.
● There is a slight increase in melting point between
phosphorus and sulfur.
● Melting point then decreases from sulfur to argon.

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4
Q

Why does melting point increase from sodium to

silicon?

A

Na → Mg → Al → Si

● Na, Mg and Al are all giant metallic structures.
● As you go from Na to Mg to Al, number of protons and electrons
increases. Atomic radius decreases.
● This leads to greater electrostatic attraction between nuclei and
electrons which requires more energy to overcome and melt the metal.
● Silicon has a giant covalent lattice structure which has strong
covalent bonds between atoms which require a lot of energy to overcome.

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5
Q

Why is there is a sharp decrease in melting point

between silicon and phosphorus?

A

● Silicon has a giant covalent lattice structure whereas
phosphorus has a simple covalent structure.
● The strong covalent bonds between the silicon atoms
require a lot of energy to overcome.

● The weak London forces between P4 molecules
require little energy to overcome.

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6
Q

Why is there is a slight increase in melting point between

phosphorus and sulfur?

A

● Sulfur has more atoms per molecule than phosphorus so
sulfur molecules contain more protons and electrons.
● As a result, the London forces between molecules are
stronger so more energy is required to overcome these
forces during melting.
● The increase is only small because sulfur is still a simple
molecular compound.

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7
Q

Why does melting point decreases from sulfur to argon?

A

● S8, Cl2 and Ar are all simple covalent substances.

● From S8 to Cl2 to Ar, the molecules are getting
smaller.
● This means that there are weaker intermolecular
(London) forces between molecules.
● As a result, less energy is required to overcome
these forces and melt the substance.

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8
Q

Describe and explain how electrical conductivity

varies across period 3

A

● Conductivity increases from sodium to magnesium to
aluminium because metallic bonding means that they
contain delocalised electrons that are free to move.
● Silicon is a semiconductor.
● Elements from phosphorus to argon are
non-conductors because they are simple molecular
substances (no delocalised electrons/ mobile charges).

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9
Q

How does ionisation energy vary across period 3?

A

In general, ionisation energy increases across a period
because:
● Nuclear charge and atomic radius increase, shielding
remains the same.
● Nuclear attraction increases.
● As a result, more energy is required to remove an
electron so ionisation energy increases.

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10
Q

What are the properties of ceramics?

A

● Strong
● High melting point
● Electrically insulating

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11
Q

Give some examples of ceramics

A

● Magnesium oxide (ionic)
● Aluminium oxide (ionic)
● Silicon dioxide (covalent)

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12
Q

How are the properties of ceramics based on their

structure?

A

● Strength: The ionic/ covalent bonds in ceramics are very
strong (giant structures).
● High melting points: Lots of energy is required to overcome
these strong ionic or covalent bonds to melt the substance.
● Electrically insulating: non-conductors. Covalent compounds
have no mobile electrons and when ionic compounds are
solid, the ions are fixed in a giant ionic lattice.

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13
Q

Describe how sodium reacts with oxygen (include an

equation)

A

2Na + ½O2 → Na2O
Sodium burns in oxygen with an orange
flame to produce sodium oxide, a white
solid.

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14
Q

Describe how magnesium reacts with oxygen

include an equation

A

Mg + ½O2 → MgO
Magnesium burns in oxygen with an
intense white flame to form magnesium
oxide, a white solid.

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15
Q

Describe how aluminium reacts with oxygen (include

an equation)

A 
4Al + 3O2 → 2Al2O3
Aluminium will burn in oxygen if 
powdered. Sprinkling this powder into a 
bunsen gives white sparkles and forms 
aluminium oxide, a white solid.
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16
Q

Describe how phosphorus reacts with oxygen

include an equation

A

P4 + 5O2 → P4O10
White phosphorus catches fire
spontaneously in air (burns with a white
flame). In excess oxygen, phosphorus (V)
oxide forms.

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17
Q

Describe how sulfur reacts with oxygen (include an

equation)

A

S + O2 → SO2
Sulfur burns in air on gentle heating with a pale blue flame. This

produces colourless SO2 gas.

To convert SO2 to SO3:
           ⇌
●  2SO2 + O2  2SO3
          ℃
●  400 - 450 , 1-2 atm, V2O5 catalyst.
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18
Q

Describe sodium reacts with chlorine (include an

equation)

A

Na + ½Cl2 → NaCl
Sodium burns in chlorine with a bright
orange flame to produce sodium
chloride, a white solid.

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19
Q

Describe how magnesium reacts with chlorine

include an equation

A

Mg + Cl2 → MgCl2
Magnesium burns in chlorine with an
intense white flame to form magnesium
chloride, a white solid.

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20
Q

Describe how aluminium reacts with chlorine

include an equation

A

2Al + 3Cl2 → 2AlCl3
Dry chlorine is passed over aluminium foil to form aluminium
chloride, a very pale yellow solid.
⇌
2AlCl3 Al2Cl6
℃
At around 180 - 190 (dependent upon pressure), AlCl3 is

converted to Al2Cl6 which then vaporises.

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21
Q

Describe how silicon reacts with chlorine (include an

equation)

A

Si + 2Cl2 → SiCl4
If chlorine is passed over powdered silicon and
heated, it reacts to form silicon tetrachloride, a
colourless liquid, which then vaporises (can be
condensed further along the apparatus).

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22
Q

Describe how phosphorus reacts with chlorine

include an equation

A

P4 + 10Cl2 → 4PCl5
White phosphorus burns spontaneously

in excess chlorine to form PCl5, an
off-white/ almost yellow solid.

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23
Q

Describe how sodium reacts with water (include an

equation)

A

2Na + 2H2O → 2NaOH + H2
A very exothermic reaction forms
hydrogen gas and a colourless solution
of sodium hydroxide.

24
Q

Describe how magnesium reacts with cold water

include an equation

A

Mg + 2H2O → Mg(OH)2 + H2
Magnesium hydroxide forms on the outside of
the metal strip. A few bubbles of hydrogen
float to the surface of container. The reaction
generally stops after this.

25
Q

Describe how magnesium reacts with

steam (include an equation)

A

Mg + H2O → MgO + H2
Magnesium burns with its typical white
flame.

26
Q

How do period 3 oxides (from sodium to sulfur) vary

in oxidation number?

A

Na2O, MgO, Al2O3, P4O10, SO2, SO3.
+1 +2 +3 +5 +4 +6
The general trend is that oxidation

number increase (apart from SO2) 
across period 3 oxides.
27
Q

How do period 3 chlorides (from sodium to

phosphorus) vary in oxidation number?

A

NaCl, MgCl , Al Cl , SiCl , PCl
+1 +2 2 +3 2 6 +4 4 +5 5
From sodium to phosphorus, the
oxidation number increases.

28
Q

Why does the oxidation number of period 3 oxides

and chlorides vary?

A

● Each element in period 3 has a different number of
electrons in its outer shell.
● Hence each element needs to gain/lose/share a different
number of electrons to have a full outer shell and form the
oxide/chloride.
● This leads to each element having a different oxidation
state.

29
Q

Describe how sodium oxide reacts with water

(include an equation)

A

Na2O + H2O → 2NaOH
● Exothermic.
● Forms a highly alkaline solution.

30
Q

Describe how magnesium oxide reacts with water

include an equation

A

MgO + H2O → Mg(OH)2
● Forms a slightly alkaline solution.

● Most of the Mg(OH)2 that is made is insoluble
and hence doesn’t dissolve in solution to
increase the pH.

31
Q

Describe how phosphorus(V) oxide reacts with water

(include an equation)

A

P4O10 + 6H2O → 4H3PO4
● Forms an acidic solution.
● Violent reaction

32
Q

Describe how sulfur dioxide reacts with

water (include an equation)

A

SO2 + H2O → H2SO3

● Forms an acidic solution

33
Q

Describe how sulfur trioxide reacts with water

(include an equation)

A

SO3 + H2O → H2SO4
● Forms an acidic solution
● Violent reaction

34
Q

How does sodium oxide react with hydrochloric

acid?

A

Na2O is a strong base. It reacts with an
acid to form a salt and water:

Na2O + 2HCl → 2NaCl + H2O

35
Q

How does magnesium oxide react with hydrochloric

acid?

A

MgO is a weaker base than Na2O though. It
reacts with warm dilute HCl to form a salt and
water:

MgO + 2HCl → MgCl2 + H2O

36
Q

What does amphoteric mean?

A

An amphoteric compound is able to act

as both an acid and a base.

37
Q

How is aluminium oxide amphoteric?

A 
Aluminium oxide (Al2O3) is amphoteric as 
it reacts with both acids and bases
38
Q

How does phosphorus (V) oxide react with NaOH?

A

There are many different reactions that
can occur between phosphorus (V) oxide
and NaOH, an example is:

P4O10 + 12NaOH → 4Na3PO4 + 6H2O

39
Q

How does sulfur dioxide react with NaOH?

A

● Sulfur dioxide is bubbled through sodium
hydroxide solution:

SO2 + 2NaOH → Na2SO3 + H2O
● If the sulfur dioxide is in excess:

Na2SO3 + H2O + SO2 → 2NaHSO3

40
Q

What oxides don’t react with water?

A

● Aluminium oxide - insoluble in water
● Silicon dioxide - breaking up its giant
covalent lattice structure is too difficult.

41
Q

Does silicon dioxide react with acids or bases?

A

Bases (e.g. sodium hydroxide)

42
Q

Describe how NaCl reacts with water (include an

equation)

A

NaCl dissolves in water to form a neutral
solution (pH 7).

NaCl(s)+ → Na- (aq) + Cl (aq)

43
Q

Describe how MgCl2 reacts with water (include an

equation)

A

MgCl2 dissolves in water to form a faintly acidic solution
(pH 6)
2+ -
MgCl2 + 6H2O → [Mg(H2O)6] + 2Cl
A small proportion of hydrogen ions are removed from
the hydrated magnesium ion, as it a weak acid:
2+ ⇌ + +
[Mg(H2O)6] + 2H2O [Mg(H2O)5(OH)] + H3O

44
Q

Describe how AlCl3 reacts with water (include an

equation)

A

AlCl3 + 6H2O → [Al(H2O)6] + 3Cl
Hydrated aluminium ions are a stronger acid than
hydrated magnesium ions so the position of
equilibrium lies further to the right:
3+ ⇌ 2+ +
[Al(H2O)6] + H2O [Al(H2O)5(OH)] + H3O

45
Q

Describe how SiCl4 reacts with water (include an

equation)

A

SiCl4 + 2H2O → SiO2 + 4HCl
Violent reaction, produces silicon dioxide and
fumes of hydrogen chloride gas.

46
Q

Write an equation for the reaction between PCl5 and

cold water

A

PCl5 + H2O → POCl3 + 2HCl

47
Q

How does PCl5 react with boiling water?

A

With water: PCl5 + H2O → POCl3 + 2HCl

If the water is boiling the POCl3 will continue to react:

POCl3 + 3H2O → H3PO4 + 3HCl
Overall boiling water equation:

PCl5 + 4H2O → H3PO4 + 5HCl

48
Q

Describe how the bonding in group 3 oxides and

chlorides varied across the period

A

Sodium and magnesium form ionic bonds with
oxygen and chlorine. Aluminium forms covalent
bonds with oxygen and either covalent or ionic
bonds with chlorine. Other period 3 elements
form simple covalent compounds.

49
Q

Why does the bonding in group 3 oxides

and chlorides vary across the period?

A

The difference in electronegativity between chlorine/
oxygen and the period 3 element decreases across the
period. There is sufficient difference in the
electronegativity of chlorine/ oxygen and sodium or
magnesium to form ions. After aluminium, the difference
in electronegativity is too small for ions to form.

50
Q

How can physical properties be used to predict the
type of chemical bonding in group 3 oxides and
chlorides?

A

A high melting point indicates a giant
molecular structure. This could either be
ionic (like NaCl and MgO) or covalent

(like SiO2).

51
Q

How can chemical properties be used to predict the
type of chemical bonding in group 3 oxides and
chlorides?

A

● Chlorides and water: ionic chlorides form a solution with a pH close to 7.
Covalent chlorides react to form an acidic solution and HCl gas.
● Oxides and water: covalent oxides form an acidic solution. Ionic oxides
may react to form an alkaline solution or they may not react.
● Acids and bases: ionic oxides are generally basic (react with acids).
Covalent oxides tend to be acidic (react with bases). Amphoteric oxides
such as aluminium oxide are usually ionic with some covalent character.
● Electrolysis: only molten ionic chlorides/ oxides undergo electrolysis.

52
Q

What is meant by periodicity?

A

The recurring variations or trends in the
properties of elements in the periodic
table.

53
Q

What group normally forms stable -1 ions?

A

Group 7:

F–, Cl–, Br– and I–

54
Q

What group normally forms stable +1 ions?

A

Group 1:

Li+, Na+, K+ and Rb+

55
Q

What group normally forms stable +2 ions?

A

Group 2:

Be2+, Mg2+, Ca2+ and Sr2+

56
Q

What structure do group 4 elements

normally have?

A

Giant covalent structure

57
Q

Which elements in the periodic table

form giant metallic structures?

A

The metals.

AS Chemistry (22 decks)

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