Chemistry Study Notes: Reactivity of Metals
Hello! Welcome to your study notes on the fascinating world of metals. Ever wondered why a gold ring stays shiny forever, but an old iron gate gets rusty? Or how we get useful metals like iron and aluminium from rocks? It all comes down to one key idea: reactivity.
In this chapter, we're going to explore:
• What the metal reactivity series is and how to use it.
• How metals can 'swap places' in displacement reactions.
• The different methods used to extract metals from the Earth.
Don't worry if this sounds tricky at first. We'll break it down with simple explanations, real-world examples, and even a memory trick to help you ace this topic!
1. The Metal Reactivity Series: A League Table for Metals
What is Reactivity?
Think of reactivity as a competition. Some metals are really eager to react with other substances like air, water, or acids, while others are quite lazy and prefer to stay as they are.
In chemical terms, reactivity is a measure of how easily a metal atom loses its outer electrons to form a positive ion (cation). The more easily a metal loses electrons, the more reactive it is.
Most Reactive Metals: Lose electrons very easily.
Least Reactive Metals: Do not lose electrons easily.
Building the "League Table"
Scientists have ranked metals from most reactive to least reactive based on experiments. This ranking is called the Reactivity Series. They observe how metals react with:
• Dilute Acids: Reactive metals (like magnesium) fizz vigorously, producing hydrogen gas. Less reactive metals (like copper) don't react at all.
• Water/Steam: Very reactive metals (like sodium) react with cold water. Others (like zinc and iron) only react with steam. The least reactive ones don't react at all.
The Metal Reactivity Series
Here is the series you need to know. Remember, the metal at the top is the most reactive, and the one at the bottom is the least reactive.
Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Zinc (Zn)
Iron (Fe)
Lead (Pb)
(Hydrogen)
Copper (Cu)
Mercury (Hg)
Silver (Ag)
Gold (Au)
Platinum (Pt)
(Note: Hydrogen is not a metal, but it's included as a reference point. Metals above hydrogen react with dilute acids to produce hydrogen gas; metals below it do not.)
A Handy Mnemonic to Remember the Series!
Memorising this list is super important. Here's a silly sentence to help you remember the order:
"Please Send Charlie's Monkeys And Zebras In Lead Cages Securely Guarded."
(Potassium, Sodium, Calcium, Magnesium, Aluminium, Zinc, Iron, Lead, Copper, Silver, Gold)
Key Takeaway: Section 1
The reactivity series is a list that ranks metals from most reactive to least reactive. A metal's position in this series tells us how likely it is to react and lose electrons to form a positive ion.
2. Displacement Reactions: The Great Metal Swap!
What is a Displacement Reaction?
Now that we have our league table, we can predict what happens when metals compete! A displacement reaction is where a more reactive metal displaces (kicks out) a less reactive metal from its compound in a solution.
Analogy: Imagine a dance. If a very popular person (a more reactive metal) wants to dance with someone who is already dancing with a less popular person (a less reactive metal), they can just cut in and take their place!
How to Predict a Displacement Reaction (Step-by-Step)
It's easy! Just follow these steps:
1. Identify the two metals involved: one is a pure metal element, and the other is an ion in a compound (a salt solution).
2. Find them in the reactivity series.
3. Compare their positions:
• If the pure metal element is HIGHER in the series, it is more reactive and a displacement reaction WILL happen.
• If the pure metal element is LOWER in the series, it is less reactive and NO REACTION will occur.
Example 1: A Reaction Happens!
Question: What happens when a piece of zinc metal is placed in blue copper(II) sulphate solution?
1. Metals: Zinc (Zn) and Copper (Cu, in copper(II) sulphate).
2. Reactivity Series: Zinc is higher than Copper.
3. Prediction: Reaction happens! Zinc will displace copper.
Observations:
• The grey zinc metal slowly dissolves.
• A reddish-brown layer of copper metal forms on the zinc.
• The blue colour of the copper(II) sulphate solution fades and becomes colourless (as zinc sulphate is formed, which is colourless).
Equations for this reaction:
Word Equation:
Zinc + Copper(II) sulphate → Zinc sulphate + Copper
Balanced Chemical Equation (with state symbols):
$$Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$$
Balanced Ionic Equation (showing the real action!):
The sulphate ion ($$SO_4^{2-}$$) is a 'spectator' - it doesn't change. We can remove it to see what's really happening. The zinc atom loses two electrons to become a zinc ion, and the copper ion gains those two electrons to become a copper atom.
$$Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$$
Example 2: No Reaction!
Question: What happens if we put a piece of copper in a zinc sulphate solution?
1. Metals: Copper (Cu) and Zinc (Zn, in zinc sulphate).
2. Reactivity Series: Copper is lower than Zinc.
3. Prediction: No reaction. Copper is not reactive enough to displace zinc.
Equation:
$$Cu(s) + ZnSO_4(aq) \rightarrow No \ reaction$$
Quick Review: Displacement Rules
• More reactive metal + Less reactive metal's salt = REACTION
• Less reactive metal + More reactive metal's salt = NO REACTION
Key Takeaway: Section 2
A more reactive metal can displace a less reactive metal from its salt solution. We can use the reactivity series to predict whether these reactions will happen.
3. Getting Metals from the Ground: Extraction Methods
Where Do We Find Metals?
Very unreactive metals like gold (Au) and platinum (Pt) are so stable that they can be found in the Earth's crust as pure elements. We call these native metals.
However, most metals are reactive, so they have already reacted with other elements (like oxygen or sulphur) over millions of years. They exist as compounds locked inside rocks, which we call metal ores. For example, iron is found as iron oxide in an ore called haematite.
To get the pure metal, we must separate it from the other elements in its ore. This process is called extraction, and it is a reduction reaction (usually involving the removal of oxygen).
Reactivity Decides the Extraction Method
The method we use to extract a metal depends entirely on its reactivity. The more reactive the metal, the more stable its ore is, and the harder it is to extract the metal.
Method 1: Electrolysis (For Very Reactive Metals)
• Metals: Potassium, Sodium, Calcium, Magnesium, Aluminium.
• Method: These metals hold onto oxygen so strongly that we need a very powerful method to split them apart. We pass a strong electric current through the molten metal ore to break it down. This is called electrolysis. It uses a huge amount of energy and is very expensive.
Method 2: Heating with Carbon (For Moderately Reactive Metals)
• Metals: Zinc, Iron, Lead.
• Method: These metals are less reactive. We can extract them by heating their ore with carbon (in the form of coke, which is cheap). Carbon is more reactive than these metals, so it displaces the metal by "stealing" the oxygen from the metal oxide. This is a reduction process.
Example: Extracting Iron in a Blast Furnace
Word Equation:
Iron(III) oxide + Carbon → Iron + Carbon monoxide
Balanced Chemical Equation:
$$Fe_2O_3(s) + 3C(s) \rightarrow 2Fe(l) + 3CO(g)$$
Method 3: Heating Alone (For Some Unreactive Metals)
• Metals: Mercury, Silver.
• Method: These metals are so unreactive that their compounds are quite unstable. Just heating their ores is often enough to decompose the compound and release the pure metal.
Example: Extracting Mercury
$$2HgO(s) \rightarrow 2Hg(l) + O_2(g)$$
Did you know? The History of Metals!
The order in which humans discovered and used metals is directly related to their ease of extraction! Gold and Silver were used first because they exist as native metals. Copper was next, as it's relatively easy to extract. The "Iron Age" came much later because extracting iron requires high temperatures (blast furnace). Aluminium, despite being very common, was only extracted commercially in the 1880s because it needs expensive electrolysis!
Conservation of Metals - Why Recycling is Key
The metal ores in the Earth's crust are a finite resource – they will eventually run out. Mining is also expensive, uses lots of energy, and damages the environment.
Recycling metals like aluminium and steel is crucial. It:
• Saves energy: Recycling an aluminium can uses only 5% of the energy needed to extract it from its ore.
• Conserves resources: It makes our limited supply of ores last longer.
• Protects the environment: It reduces waste in landfills and reduces the damage caused by mining.
Key Takeaway: Section 3
A metal's position in the reactivity series determines its extraction method. The more reactive the metal, the more difficult and energy-intensive the extraction. Unreactive metals are found as elements, while reactive metals are extracted from ores using methods like heating with carbon or electrolysis.