Period 3- Chemistry Unit 5

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  • Created by: Hannah
  • Created on: 22-06-11 19:16

Introduction

Period 3 = sodium to argon

Don't use first two periods as period 1 only has two elements and period 2 has the top elements of each group (small sizes and relatively high ionsiation energies) so are atypical. Period 3 is best for studying periodic trends.

Elements:

Na, Mg, Al= metals

Si = metalloid 

P4, S8, Cl2, Ar= non-metals (simple molecules)

From left to right across the period the metals tend to be less reactive. e.g. sodium is more reactive then magnesium as sodium produces a Na+ ion whilst magnesium makes an Mg2+ion. (It takes less energy to lose one electron than to lose two). 

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Reactions with water...

Sodium:

-Vigorous reaction with cold water (molten ball on surface fizzing/ H2 gas)

- 2Na (s) + 2H20 (g) ---> 2NaOH (aq) + H2 (g)

- Strong alkaline solution formed, sodium hydroxide, very soluble

Magnesium:

- Very slow reaction with cold water (cannot see reaction)

- Mg (s) + 2H2O (g) ---> Mg(OH)2 (s) + H2 (g)

- Produces a weak alkaline solution that is sparingly soluble

-Vigorous reaction with steam (more energy) to form magnesium oxide

- Mg (s) + H2O (g) ---> MgO (s) + H2 (g) 

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Reactions with oxygen...

Period 3 elements form oxides when they react with oxygen. They are usually oxidised t their highest states.

Sulfur is the exception- it forms SO2, in which it's only got a +4 oxidation state (high temp and a catalyst are needed to make SO3). 

All equations go: element + oxygen ---> oxide

Sodium:

-Vigorous reaction with ignited sodium, ionic sodium oxide formed:

- 2Na (s) + 1/2O2 (g) ---> Na2O (s)

Magnesium:

-Vigorous reaction with ignited magnesium, ionic magnesium oxide formed:

- Mg (s) + 1/2O2 (g) ---> MgO (s)

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Reactions with oxygen...

Aluminium:

-Sheets of it get slowly coated with a thin oxide layer. Powdered aluminium shows a vigorous reaction with sparks. Ionic aluminium oxide formed:

-2Al (s) + 1 1/2 O2 (g) ---> Al2O3 (s) 

Silicon:

-Vigorous reaction with silicon powder. Covalent giant molecular silicon dioxide formed:

- Si (s) + O2 (g) ---> SiO2 (s)

Phosphorus:

- Ignites spontaneously in oxygen, white solid produced of phosphorus (V) oxide. (P4 is a common allotrope (form) of phosphorus):

- P4 (s) + 5O2 ---> P4O10 (s)

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Reactions with oxygen...

Sulfur:

- Burns with a lilac flame to give a choking gas which fumes in most air. Covalent molecules of sulfur dioxide formed:

- S (s) + O2 (g) ---> SO2 (g)

The more reactive metals (Na, Mg) and the non-metals (P, S) react readily in air. While Al and Si react slowly.

Summary= (in order of Na, Mg, Al, Si, P and S)

Formula of oxide = Na2O, MgO, Al2O3, SiO2, P4O10, SO2

Reaction in air = vigorous, vigorous, slow, slow, spontaneously combusts, burns steadily

Flame= yellow, brilliant white, n/a, n/a, brilliant white, blue 

Can identify some of the oxides using flame tests

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Bonding and structure affect melting points of oxi

Na2O, MgO and Al2O3- the metal oxides

They have high melting points because they form giant ionic lattices. 

The strong forces of attraction between each ion mean it takes a lot of heat energy to break the bonds and melt them.

MgO has a higher melting point than Na2O because it forms 2+ ions, so bonds more strongly than the 1+ ions in Na2O.

Al2O3 has a lower melting point than you might expect because the 3+ ions distort the oxygen's electron cloud, making the bonds partially covalent.

SiO2 has a higher melting point than the other non-metal oxides because it has a giant macromolecular structure.

P4O10 and SO2 have relatively low melting points because they form simple molecular structures. The molecules are bound by weak intermolecular forces (dipole-dipole and van der Waals which takes little energy to overcome. 

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Ionic oxides are Alkaline, Covalent oxides are Aci

The ionic oxides of the metals Na and Mg dissolve in water to form hydroxides.

Solutions are both alkaline, but sodium hydroxide is more soluble in water, so it forms a more alkaline solution than magnesium hydroxide. 

Sodium:  Na2O (s) + H2O (l) ---> 2NaOH (aq) .... pH = 12-14

Magnesium:   MgO (s) + H2O (l) ---> Mg(OH)2 (aq) ....pH = 9-10

The simple covalent oxides of the non-metals P and S form acidic solutions. ALl of the acids are strong and so the pH of their solutions is about 0-2.

P4O10 (s) + 6H2O (l) ---> 4H3PO4 (aq) = Phosphoric (V) acid

SO2 (g) + H2O (l) ---> H2SO3 (aq) = Sulfurous acid (or sulfuric (IV) acid)

SO3 (l) + H2O (l) ---> H2SO4 (aq) = Sulfuric (VI) acid

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Silicon dioxide and Aluminium oxide

The giant covalent structure of silicon dioxide means that it is insoluble in water.

However, it will react with bases to form salts, so it is classed as acidic. 

Aluminium oxide, which is partially ionic and partially covalently bonded is also insoluble in water.

But it react with acids and bases to form salts. It can act as an acid or base- is amphoteric. 

Oxidation states, just so you know:

Na = +1, Mg = +2, Al = +3, Si = +4, P = +5, S = +6

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Acid + Base ---> Salt + Water

1. Basic oxides neutralise acids:

Na2O (s) + 2HCl (aq) ---> 2NaCl (aq) + H2O (l)

MgO (s) + H2SO4 (aq) ---> MgSO4 (aq) + H2O (l) 

2. Acidic oxides neutralise bases:

SiO2 (s) + 2NaOH (aq) ---> Na2SiO3 (aq) + H2O (l)

P4O10 (s) + 12NaOH (aq) ---> 4Na3PO4 (aq) + 6H2O (l)

SO2 (g) + 2NaOH (aq) ---> Na2SO3 (aq) + H2O (l)

SO3 (g) + 2NaOH (aq) --->Na2SO4 (aq) + H2O (l)

3. Amphoteric oxides neutralise acids and bases:

Al2O3 (s) + 3H2SO4 (aq) ---> Al2(SO4)3 (aq) + 3H2O (l)

Al2O3 (s) + 2NaOH (aq) + 3H2O (l) ---> 2NaAl(OH)4 (aq)

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Useful stuff...

Ionic bonding - Bond formed after two or more atoms give up or gain electrons, and become ions. Occurs between metals and non-metals. (Electrostatic attraction between oppositely charged ions). The positive ions are formed from metals (usually) and negative ions are usually formed from non-metals.

Covalent bond- Strong bond between two non- metal atoms. It consists of a shared pair of electrons. 

Trend 1: Atomic radius of the elements decrease across the period from left to right. Proton number increases and an electron is added to the same electron shell. Increased nuclear charge attracts electrons.

Trend 2: 1st Ionisation Energy. (removing an electron from a gaseous atom). Increases across the period from left to right. The further away from the positively charged nucleus that an electron is, the less strongly it is attracted, so it can be removed more easily. S atomic radius decreases, the ionisation energy increases.

Trend 3: Electronegativity (ability of an atom to attract electrons to itself). This increases across the period from left to right. Typically metals have low electronegativity (little ability to attract e-s) and non-metals have high electronegativity. 

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