reactivity of metals Flashcards
describe how metals react with oxygen and state the compound they form.
Reaction of Metals with Oxygen:
When a metal reacts with oxygen, it forms a metal oxide.
The metal loses electrons (oxidation), and the oxygen gains electrons (reduction).
General Formula:
The general formula for the metal oxide is usually written as metal + oxygen → metal oxide.
Example: 2Mg + O₂ → 2MgO (Magnesium reacts with oxygen to form magnesium oxide).
Reactivity:
Reactive metals (like sodium, potassium, magnesium) react vigorously with oxygen and form white oxides.
Less reactive metals (like copper, gold, and platinum) may react slowly or not at all with oxygen, forming oxides under certain conditions.
Properties of Metal Oxides:
Many metal oxides are basic and can react with acids to form salts and water.
Some, like aluminium oxide, form protective layers that prevent further oxidation.
define oxidation
The loss of electrons by a substance.
The increase in oxidation state of an element.
define reduction
The gain of electrons by a substance.
The decrease in oxidation state of an element.
describe the arrangements of metals in the reactivity series, including carbon and hydrogen, and use the reactivity series to predict the outcome of displacement reactions.
Reactivity Series (Order from Most to Least Reactive):
Potassium (K), Sodium (Na), Lithium (Li), Calcium (Ca), Magnesium (Mg), Zinc (Zn), Iron (Fe), Tin (Sn), Lead (Pb), Hydrogen (H), Copper (Cu), Silver (Ag), Gold (Au), Platinum (Pt).
Metals like potassium, sodium, and calcium are highly reactive, and metals like gold and platinum are very unreactive.
Carbon is placed between zinc and iron in the reactivity series because carbon can be used to extract metals from their ores (e.g., extracting iron from iron ore).
Hydrogen is included because some metals react with acids to release hydrogen gas (e.g., zinc reacting with hydrochloric acid).
Displacement Reactions:
A displacement reaction occurs when a more reactive metal displaces a less reactive metal from its compound. For example, if a metal is placed in a solution of a salt containing a metal lower in the reactivity series, it will displace the less reactive metal and form its own salt.
Key points for predicting outcomes of displacement reactions:
More reactive metal displaces a less reactive metal from its compound.
Example: Magnesium (Mg) can displace copper (Cu) from copper sulfate:
Mg
+
CuSO
4
→
MgSO
4
+
Cu
Mg+CuSO
4
→MgSO
4
+Cu
Magnesium is more reactive than copper, so it displaces copper.
Metals react with acids to produce hydrogen gas. The more reactive a metal, the easier it is to release hydrogen.
Example: Zinc (Zn) reacts with hydrochloric acid (HCl) to produce hydrogen gas:
Zn
+
2
HCl
→
ZnCl
2
+
H
2
Zn+2HCl→ZnCl
2
+H
2
Metals above hydrogen in the series can displace hydrogen from acids.
Example: Magnesium reacts with hydrochloric acid to release hydrogen gas, while copper, being less reactive than hydrogen, does not react with acids to produce hydrogen.
Using the Reactivity Series:
When predicting if a displacement reaction will happen, check if the metal is higher in the reactivity series than the metal in the compound. If it is, a displacement reaction will occur.
For example: Iron can displace copper from copper sulfate, but copper cannot displace iron from iron sulfate.
recall and describe the reactions, if any of potassium lithium sodium calcium magnesium zinc and carbon with water or dilute acids.
Reactions with Water:
Potassium (K):
Very reactive with water.
Reacts explosively, forming potassium hydroxide (KOH) and hydrogen gas (H₂).
Reaction:
2
𝐾
+
2
𝐻
2
𝑂
→
2
𝐾
𝑂
𝐻
+
𝐻
2
2K+2H
2
O→2KOH+H
2
Potassium can ignite hydrogen, causing a small explosion or flame.
Sodium (Na):
Reacts vigorously with water.
Forms sodium hydroxide (NaOH) and hydrogen gas (H₂).
Reaction:
2
𝑁
𝑎
+
2
𝐻
2
𝑂
→
2
𝑁
𝑎
𝑂
𝐻
+
𝐻
2
2Na+2H
2
O→2NaOH+H
2
Sodium moves around on the surface of the water, often with a fizzing motion, and can ignite the hydrogen gas.
Lithium (Li):
Reacts more gently with water compared to sodium or potassium.
Forms lithium hydroxide (LiOH) and hydrogen gas (H₂).
Reaction:
2
𝐿
𝑖
+
2
𝐻
2
𝑂
→
2
𝐿
𝑖
𝑂
𝐻
+
𝐻
2
2Li+2H
2
O→2LiOH+H
2
Lithium will float and fizz on the surface.
Calcium (Ca):
Reacts fairly vigorously with water.
Forms calcium hydroxide (Ca(OH)₂ and hydrogen gas (H₂).
Reaction:
𝐶
𝑎
+
2
𝐻
2
𝑂
→
𝐶
𝑎
(
𝑂
𝐻
)
2
+
𝐻
2
Ca+2H
2
O→Ca(OH)
2
+H
2
The reaction is less violent than sodium or potassium, but calcium still releases gas and heats up the water.
Magnesium (Mg):
Reacts slowly with water at room temperature.
Forms magnesium hydroxide (Mg(OH)₂ and hydrogen gas (H₂).
Reaction:
𝑀
𝑔
+
2
𝐻
2
𝑂
→
𝑀
𝑔
(
𝑂
𝐻
)
2
+
𝐻
2
Mg+2H
2
O→Mg(OH)
2
+H
2
Magnesium needs to be heated to react more readily with water.
Zinc (Zn):
Reacts very slowly with cold water, and often does not react significantly at room temperature.
When heated, it reacts with water to form zinc hydroxide (Zn(OH)₂) and hydrogen gas (H₂).
Reaction:
𝑍
𝑛
+
2
𝐻
2
𝑂
→
𝑍
𝑛
(
𝑂
𝐻
)
2
+
𝐻
2
Zn+2H
2
O→Zn(OH)
2
+H
2
Zinc reacts more readily with acids than with water.
Carbon (C):
Does not react with water under normal conditions.
Reactions with Dilute Acids:
Potassium (K):
Reacts very vigorously with dilute acids.
Forms potassium salts (e.g., potassium chloride) and hydrogen gas (H₂).
Reaction with hydrochloric acid:
2
𝐾
+
2
𝐻
𝐶
𝑙
→
2
𝐾
𝐶
𝑙
+
𝐻
2
2K+2HCl→2KCl+H
2
The reaction is very exothermic, producing heat and hydrogen gas quickly.
Sodium (Na):
Reacts vigorously with dilute acids.
Forms sodium salts (e.g., sodium chloride) and hydrogen gas (H₂).
Reaction with hydrochloric acid:
2
𝑁
𝑎
+
2
𝐻
𝐶
𝑙
→
2
𝑁
𝑎
𝐶
𝑙
+
𝐻
2
2Na+2HCl→2NaCl+H
2
Sodium reacts quickly, releasing hydrogen gas and heat.
Lithium (Li):
Reacts gently with dilute acids.
Forms lithium salts (e.g., lithium chloride) and hydrogen gas (H₂).
Reaction with hydrochloric acid:
2
𝐿
𝑖
+
2
𝐻
𝐶
𝑙
→
2
𝐿
𝑖
𝐶
𝑙
+
𝐻
2
2Li+2HCl→2LiCl+H
2
The reaction is less vigorous than sodium or potassium.
Calcium (Ca):
Reacts quite vigorously with dilute acids.
Forms calcium salts (e.g., calcium chloride) and hydrogen gas (H₂).
Reaction with hydrochloric acid:
𝐶
𝑎
+
2
𝐻
𝐶
𝑙
→
𝐶
𝑎
𝐶
𝑙
2
+
𝐻
2
Ca+2HCl→CaCl
2
+H
2
Calcium reacts rapidly, producing bubbles of hydrogen gas.
Magnesium (Mg):
Reacts fairly vigorously with dilute acids.
Forms magnesium salts (e.g., magnesium chloride) and hydrogen gas (H₂).
Reaction with hydrochloric acid:
𝑀
𝑔
+
2
𝐻
𝐶
𝑙
→
𝑀
𝑔
𝐶
𝑙
2
+
𝐻
2
Mg+2HCl→MgCl
2
+H
2
Magnesium reacts readily, producing bubbles of hydrogen gas.
Zinc (Zn):
Reacts with dilute acids, but less vigorously than magnesium or calcium.
Forms zinc salts (e.g., zinc chloride) and hydrogen gas (H₂).
Reaction with hydrochloric acid:
𝑍
𝑛
+
2
𝐻
𝐶
𝑙
→
𝑍
𝑛
𝐶
𝑙
2
+
𝐻
2
Zn+2HCl→ZnCl
2
+H
2
Zinc reacts steadily with dilute acids to produce hydrogen gas.
Carbon (C):
Does not react with dilute acids under normal conditions.
relate the reactivity of metals to its tendency to form positive ions and be able to deduce an order of reactivity of metals based on experimental results.
Key Concepts:
Formation of Positive Ions:
Metals are oxidized during reactions, meaning they lose electrons to form positive ions (cations).
The easier it is for a metal to lose electrons, the more reactive the metal is.
Metals in the top of the reactivity series (like potassium and sodium) lose their outer electrons more easily than metals lower down the series (like copper or gold).
Reactivity and Electron Loss:
Highly reactive metals (like potassium, sodium, and magnesium) have few electrons in their outer shell, so they lose them easily to form positive ions.
Less reactive metals (like copper and gold) have more electrons in their outer shell or a more stable electron configuration, so they do not lose electrons as easily, and hence are less reactive.
Relating Reactivity to the Tendency to Form Positive Ions:
The more reactive a metal is, the more likely it is to lose electrons and form positive ions in reactions, such as with acids or water.
The reactivity series shows metals ordered from most reactive to least reactive based on their tendency to form positive ions.
Experimental Evidence for Reactivity:
In experiments, the reactivity of metals can be determined based on:
Reactions with water: More reactive metals (like potassium, sodium, and calcium) react vigorously with water, forming hydroxides and hydrogen gas. Less reactive metals (like magnesium and zinc) react more slowly or not at all with cold water.
Reactions with dilute acids: More reactive metals (like magnesium, zinc, and calcium) react vigorously with dilute acids, displacing hydrogen gas and forming salts. Less reactive metals (like copper and gold) do not react with dilute acids to produce hydrogen gas.
Displacement reactions: A more reactive metal can displace a less reactive metal from a compound in solution. For example, magnesium can displace copper from copper sulfate solution, but copper cannot displace magnesium from magnesium sulfate.
Order of Reactivity Based on Experimental Results:
Based on experimental observations, we can order metals from most reactive to least reactive in the reactivity series:
Potassium (K) – reacts explosively with water, very vigorous with acids.
Sodium (Na) – reacts vigorously with water and acids.
Lithium (Li) – reacts gently with water and acids.
Calcium (Ca) – reacts fairly vigorously with water and acids.
Magnesium (Mg) – reacts slowly with water (more vigorously with acids).
Zinc (Zn) – reacts slowly with acids but not with water.
Iron (Fe) – reacts slowly with acids and water.
Tin (Sn) – reacts very slowly with acids.
Lead (Pb) – reacts very slowly with acids.
Hydrogen (H) – included for comparison, does not react with water or dilute acids.
Copper (Cu) – does not react with water or dilute acids.
Silver (Ag) – very unreactive, does not react with water or dilute acids.
Gold (Au) – inert, does not react with water or dilute acids.
Summary of Key Points:
Reactivity increases as the tendency to lose electrons (and form positive ions) increases.
Metals at the top of the reactivity series lose electrons more easily, so they are more reactive.
Experimental results such as reactions with water, acids, and displacement reactions help determine the order of reactivity in the series.
recall what native metals are and explain how metals can be extracted from the compounds in which they are found in nature by reduction with carbon.
Native Metals:
Native metals are metals that are found in their pure, uncombined form in nature, usually as metallic elements, not in compounds.
These metals do not need to undergo chemical reactions to be extracted because they are already in their pure form.
Common native metals include gold (Au), platinum (Pt), and copper (Cu), which are found naturally in the Earth’s crust in a pure state, often in veins or nuggets.
Gold is highly unreactive, so it does not easily form compounds with other elements, which is why it can be found in its native form.
Extraction of Metals from Compounds:
Most metals, however, are found as compounds in ores (minerals that contain metal compounds). To extract the metal from these compounds, a process called reduction is used. Reduction is the process of removing oxygen from a metal ore to obtain the metal itself.
Reduction with Carbon:
Carbon is commonly used to extract metals from their oxides (metal ores that contain oxygen) because carbon is a reducing agent.
In reduction with carbon, carbon monoxide (CO) is produced when carbon reacts with oxygen from the metal ore. This carbon monoxide removes the oxygen from the metal oxide, leaving the metal behind.
Process of Reduction with Carbon:
Metal oxide + Carbon → Metal + Carbon dioxide
The carbon reacts with the oxygen in the metal oxide, forming carbon dioxide (CO₂) and leaving the pure metal.
Example: Extraction of Iron (Fe):
Iron is extracted from its ore, iron oxide (Fe₂O₃), using carbon (usually in the form of coke) in a blast furnace.
The reaction is:
𝐹
𝑒
2
𝑂
3
+
3
𝐶
→
2
𝐹
𝑒
+
3
𝐶
𝑂
2
Fe
2
O
3
+3C→2Fe+3CO
2
Carbon removes the oxygen from iron oxide, producing iron and carbon dioxide.
Example: Extraction of Copper (Cu):
Copper can be extracted from copper(II) oxide (CuO) by reduction with carbon:
𝐶
𝑢
𝑂
+
𝐶
→
𝐶
𝑢
+
𝐶
𝑂
2
CuO+C→Cu+CO
2
The carbon removes the oxygen from copper oxide to produce copper metal.
When is Carbon Used for Reduction?
Carbon is used for extracting metals that are less reactive than carbon itself. This is why it is used for metals like iron, copper, zinc, and tin.
Highly reactive metals (like potassium, sodium, and calcium) cannot be extracted using carbon because they are more reactive than carbon and would not easily lose their oxygen to carbon. For these metals, more powerful methods, such as electrolysis, are used.
Summary:
Native metals are metals found in their pure, metallic form in nature (e.g., gold, copper, platinum).
Most metals are found as compounds in ores, and to extract them, reduction is used to remove oxygen from the metal oxide.
Reduction with carbon is a common method for extracting metals like iron and copper, where carbon removes the oxygen from the metal oxide to produce the metal and carbon dioxide.
evaluate specific metal extraction processes when given appropriate information and identify which species are oxidised or reduced.
- Extraction of Iron (Fe) in a Blast Furnace:
Reaction Overview:
Iron is extracted from its ore, iron oxide (Fe₂O₃), using carbon in a blast furnace.
The blast furnace uses coke (a form of carbon) to reduce the iron ore into iron.
Key Reactions:
Coke (carbon) reacts with oxygen to form carbon dioxide (CO₂):
𝐶
+
𝑂
2
→
𝐶
𝑂
2
(oxidationofcarbon)
C+O
2
→CO
2
(oxidationofcarbon)
Carbon monoxide (CO), produced in the furnace, reduces the iron(III) oxide (Fe₂O₃) to iron (Fe):
𝐹
𝑒
2
𝑂
3
+
3
𝐶
𝑂
→
2
𝐹
𝑒
+
3
𝐶
𝑂
2
(reductionofironoxide)
Fe
2
O
3
+3CO→2Fe+3CO
2
(reductionofironoxide)
Redox Evaluation:
Oxidation: The carbon in coke loses electrons to form carbon dioxide (
𝐶
→
𝐶
𝑂
2
+
2
𝑒
−
C→CO
2
+2e
−
).
Reduction: The iron(III) ions (Fe³⁺) in iron oxide gain electrons to form iron metal (
𝐹
𝑒
3
+
+
3
𝑒
−
→
𝐹
𝑒
Fe
3+
+3e
−
→Fe).
Species Oxidised and Reduced:
Oxidised: Carbon (C) is oxidised to carbon dioxide (CO₂).
Reduced: Iron (III) oxide (Fe₂O₃) is reduced to iron (Fe).
Evaluation:
The blast furnace is an effective and cost-efficient method for extracting iron from its ore, particularly since carbon is relatively cheap and abundant. However, it produces carbon dioxide, a greenhouse gas contributing to climate change.
2. Extraction of Copper (Cu) from Copper(II) Oxide:
Reaction Overview:
Copper can be extracted from its copper(II) oxide (CuO) by reduction with carbon.
Key Reaction:
Copper(II) oxide reacts with carbon (usually in the form of coke) to produce copper metal and carbon dioxide:
𝐶
𝑢
𝑂
+
𝐶
→
𝐶
𝑢
+
𝐶
𝑂
2
CuO+C→Cu+CO
2
Redox Evaluation:
Oxidation: The carbon (C) loses electrons to form carbon dioxide (CO₂) (
𝐶
→
𝐶
𝑂
2
+
2
𝑒
−
C→CO
2
+2e
−
).
Reduction: The copper(II) ions (Cu²⁺) in copper oxide gain electrons to form copper metal (Cu) (
𝐶
𝑢
2
+
+
2
𝑒
−
→
𝐶
𝑢
Cu
2+
+2e
−
→Cu).
Species Oxidised and Reduced:
Oxidised: Carbon (C) is oxidised to carbon dioxide (CO₂).
Reduced: Copper(II) ions (Cu²⁺) are reduced to copper metal (Cu).
Evaluation:
This method is used for extracting copper from copper oxide and works well because carbon is a relatively cheap reducing agent.
However, copper oxide must first be obtained, and large-scale copper extraction can still have environmental impacts, such as mining waste and energy use.
