3.2.1. What happens when Metals are burnt in Air?

You have seen in Activity 3.8 that magnesium burns in air with a dazzling white flame. Do all metals react in the same manner? Let us check by performing the following Activity.

This introductory statement sets the stage for exploring how different metals behave when they are burned in air. It prompts curiosity about whether all metals will react similarly to oxygen in the air, starting with magnesium as a clear example of a metal that reacts dramatically.

Think of firecrackers: just like they produce different colors and sounds when ignited, different metals can produce varying reactions when burned in the air.

CAUTION: The following activity needs the teacher’s assistance. It would be better if students wear eye protection. n Hold any of the samples taken above with a pair of tongs and try burning over a flame. Repeat with the other metal samples. n Collect the product if formed. n Let the products and the metal surface cool down. n Which metals burn easily? n What flame colour did you observe when the metal burnt? n How does the metal surface appear after burning? n Arrange the metals in the decreasing order of their reactivity towards oxygen. n Are the products soluble in water?

This section outlines an experimental activity to observe how various metals react with oxygen. It highlights the importance of safety precautions and gives a structured approach: burn the metals, collect the products, observe the flame color, assess reactivity, and check product solubility. This step-by-step guidance helps students learn through direct observation and experimentation.

Imagine conducting a science experiment akin to a cooking show. Each metal is like a different ingredient, and by 'cooking' them with flame, students get to see how they transform, similar to how ingredients react when heated.

Almost all metals combine with oxygen to form metal oxides. Metal + Oxygen → Metal oxide. For example, when copper is heated in air, it combines with oxygen to form copper(II) oxide, a black oxide. 2Cu + O → 2CuO (Copper) (Copper(II) oxide). Similarly, aluminium forms aluminium oxide. 4Al + 3O → 2Al2O3 (Aluminium) (Aluminium oxide).

This section describes the chemical reaction that occurs when metals burn in air. It explains that metals generally react with oxygen to create metal oxides, illustrating the concept with copper and aluminum as examples. The provided chemical equations reinforce how elements combine during combustion.

Think of a metal like a friend who loves to team up with others. Just like how friends can form a strong group to achieve something great together, metals 'team up' with oxygen to form new substances (metal oxides) that have different properties than the metals alone.

Recall from Chapter 2, how copper oxide reacts with hydrochloric acid. We have learnt that metal oxides are basic in nature. But some metal oxides, such as aluminium oxide, zinc oxide show both acidic as well as basic behaviour. Such metal oxides which react with both acids as well as bases to produce salts and water are known as amphoteric oxides.

This chunk explains that most metal oxides are basic, which means they can react with acids. However, some metal oxides exhibit amphoteric behavior, meaning they can react with both acids and bases. This duality allows them to produce salts and water, illustrating the versatility of certain metal oxides.

Think of amphoteric oxides as bilingual friends who can speak two languages, allowing them to interact comfortably in various situations—whether it involves acids or bases, they can communicate effectively.

We have observed in Activity 3.9 that all metals do not react with oxygen at the same rate. Different metals show different reactivities towards oxygen. Metals such as potassium and sodium react so vigorously that they catch fire if kept in the open. Hence, to protect them and to prevent accidental fires, they are kept immersed in kerosene oil.

This section discusses the variability in the reactivity of metals when exposed to oxygen. It highlights that while some metals, like potassium and sodium, are highly reactive and can catch fire, others are less reactive. This information is crucial for understanding how metals behave in different environments and the precautions needed for storage.

Picture the metals as animals in a wild. Just as some animals are more aggressive or active than others, certain metals are 'wild' and react aggressively with their environment, necessitating careful handling to avoid accidents.

At ordinary temperature, the surfaces of metals such as magnesium, aluminium, zinc, lead, etc., are covered with a thin layer of oxide. The protective oxide layer prevents the metal from further oxidation. Iron does not burn on heating but iron filings burn vigorously when sprinkled in the flame of the burner. Copper does not burn, but the hot metal is coated with a black coloured layer of copper(II) oxide. Silver and gold do not react with oxygen even at high temperatures.

This chunk explains the protective nature of oxide layers that form on certain metals, which can shield them against rapid oxidation. It also mentions that while some metals burn easily, others like iron and copper may not burn in bulk but can still oxidize when in fine forms or when heated. Notably, precious metals like silver and gold remain uninhibited by oxidation.

Think of the oxide layer like a superhero shield protecting a city. Just as a shield can prevent destruction, the oxide layer guards the metal beneath, preserving it from further damage by air or moisture.

Anodising is a process of forming a thick oxide layer of aluminium. Aluminium develops a thin oxide layer when exposed to air. This aluminium oxide coat makes it resistant to further corrosion. The resistance can be improved further by making the oxide layer thicker. During anodising, a clean aluminium article is made the anode and is electrolysed with dilute sulphuric acid. The oxygen gas evolved at the anode reacts with aluminium to make a thicker protective oxide layer.

Anodising involves electrochemically enhancing the protective oxide layer on aluminum. This thicker layer not only offers better resistance to corrosion but also allows for the possibility of dyeing it for aesthetic purposes. This process is important for extending the lifespan of aluminum products.

Imagine coating an ice cream cone with chocolate to protect it from getting soggy. Just like that chocolate layer strengthens and preserves the ice cream, anodising strengthens aluminium against corrosion.

After performing Activity 3.9, you must have observed that sodium is the most reactive of the samples of metals taken here. The reaction of magnesium is less vigorous implying that it is not as reactive as sodium. But burning in oxygen does not help us to decide about the reactivity of zinc, iron, copper or lead. Let us see some more reactions to arrive at a conclusion about the order of reactivity of these metals.

This conclusion emphasizes that direct observation of burning metals does not provide enough information to determine the complete reactivity series of metals. By indicating the need for further experiments, it encourages students to engage in active learning and exploration of chemical properties.

Just like how you may need more than a single test to understand someone’s personality, determining a metal's reactivity requires a comprehensive study of various reactions to truly understand its character.

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