8.2.2 Testing Period 3 Oxides

Aim:

To investigate the reactions of Period 3 elements with oxygen, form their oxides, and test the pH of the resulting aqueous solutions.

Introduction:

Period 3 elements, ranging from sodium ($Na$) to sulfur ($S$), react with oxygen to form oxides.

These oxides show varying properties depending on whether the element is metallic or non-metallic. The metallic elements ($Na$, $Mg$, $Al$) form ionic oxides, while the non-metallic elements ($Si$, $P$, $S$) form covalent molecular oxides.

The oxidation states of these elements in their oxides vary, and their reactions with water result in solutions of different pH values, providing insight into their acidic, basic, or amphoteric nature.

Materials and Equipment:

• Sodium ($Na$), Magnesium ($Mg$), Aluminium ($Al$), Silicon ($Si$), Phosphorus ($P$), and Sulfur ($S$)
• Oxygen supply (combustion jar or oxygen cylinder)
• Deflagrating spoon and tongs
• Bunsen burner
• pH probe or pH paper/universal indicator
• Water (to dissolve oxides)
• Heatproof mat, crucible, and safety equipment (gloves, goggles)

Method:

Part 1: Forming Period 3 Oxides

1. Reacting the metals:

• Sodium ($Na$): Using a deflagrating spoon, heat a small piece of sodium and lower it into a jar of oxygen. Observe the reaction, which involves a bright yellow flame.
• Magnesium (Mg): Heat a magnesium ribbon using a Bunsen burner until it ignites with a bright white flame. Place it into a combustion jar filled with oxygen for a vigorous reaction.
• Aluminium (Al): Hold a strip of aluminium using tongs and heat it strongly in a Bunsen burner flame. It glows brightly, forming a white oxide ($Al₂O₃$).

2. Reacting the non-metals:

• Silicon ($Si$): Heat a small sample of silicon in oxygen to produce silicon dioxide ($SiO₂$), although this reaction requires a very high temperature.
• Phosphorus ($P$): Handle white phosphorus with care as it spontaneously ignites in the air, producing dense white fumes of phosphorus pentoxide ($P₄O₁₀$).
• For red phosphorus, heat it in oxygen where it reacts less vigorously to form the same product.
• Sulfur ($S$): Heat sulfur in a deflagrating spoon until it melts, then lower it into the oxygen jar.
• Sulfur burns with a blue flame to form sulfur dioxide ($SO₂$).
• Further oxidation to sulfur trioxide ($SO₃$) occurs slowly without a catalyst.

Part 2: Testing the pH of the Oxides

1. After each element reacts with oxygen and forms its oxide, dissolve the resulting oxide in water.
2. Use a pH probe or universal indicator paper to test the pH of the aqueous solution for each oxide.

Observations and Reactions:

Sodium ($Na$):

• Observation: Burns with a bright yellow flame.
• Equation:

$$4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$$

• pH of aqueous oxide: Strongly basic (around 13-14), due to the formation of sodium hydroxide ($NaOH$) in water.

Magnesium ($Mg$):

• Observation: Burns with a bright white flame.
• Equation:

$$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$$

• pH of aqueous oxide: Basic (around 10), due to the formation of magnesium hydroxide ($Mg(OH)₂$) in water.

Aluminium ($Al$):

• Observation: Glows brightly when heated.
• Equation:

$$4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3(s)$$

• pH of aqueous oxide: Amphoteric, reacts with both acids and bases, showing neutral pH (~7) when tested with water but capable of reacting as both a base and an acid.

Silicon ($Si$):

• Observation: Requires high temperature, and forms solid oxide.
• Equation:

$$Si(s) + O_2(g) \rightarrow SiO_2(s)$$

• pH of aqueous oxide: Acidic (around 7), due to the weak reactivity of silicon dioxide in water (does not dissolve well).

Phosphorus ($P$):

• Observation: White phosphorus ignites spontaneously, producing dense white fumes.
• Equation:

$$P_4(s) + 5O_2(g) \rightarrow P_4O_{10}(s)$$

• pH of aqueous oxide: Strongly acidic (around 1-2), due to the formation of phosphoric acid ($H₃PO₄$) when dissolved in water.

Sulfur ($S$):

• Observation: Burns with a blue flame.
• Equation:

$$S(s) + O_2(g) \rightarrow SO_2(g)$$ $$2SO_2(g) + O_2(g) \xrightarrow{\text{V}_2O_5} 2SO_3(g)\ \text{(with vanadium oxide catalyst)}$$

• pH of aqueous oxide: Acidic (around 2-3), due to the formation of sulfurous acid ($H₂SO₃$) and, if oxidized further, sulfuric acid ($H₂SO₄$).

Conclusion:

• The metallic Period 3 elements ($Na$, $Mg$, $Al$) react vigorously with oxygen to form ionic oxides. Sodium and magnesium oxides are basic, while aluminium oxide is amphoteric.
• The non-metallic elements ($Si$, $P$, $S$) form covalent oxides. Phosphorus and sulfur form acidic oxides, while silicon dioxide has weak reactivity and does not significantly affect the pH of water.
• The pH of the oxides reflects their acidic, basic, or amphoteric nature and helps classify the elements into metals and non-metals based on their reactions with oxygen and water.