Metals and Non-metals-Notes
Chemistry - Notes
Metals
Metals are defined as opaque, lustrous, fusible, ductile substances that conduct heat and electricity well. They are also malleable, meaning they can be pressed into thin sheets without breaking.
Properties of metals
- Conductivity:
Metals are good conductors of heat and electricity. - Malleability:
Metals can be hammered or pressed into thin sheets without breaking. - Ductility:
Metals can be drawn into wires. - Reflectivity:
Metals are highly reflective, making them shiny. - Melting and boiling points:
Most metals have high melting and boiling points. - Density:
Metals are highly dense. - Sonorous:
Metals are sonorous
Examples of metals: Gold, Silver, Aluminium, Copper, and Iron.
Non-Metals
Non-metals are chemical elements that lack metallic properties and are poor conductors of heat and electricity. They are usually gases, but can also be liquids or solids.
Physical properties
- Non-metals are usually lighter than metals
- They are often brittle, soft, and dull in appearance
- They have low melting and boiling points
- They are poor conductors of heat and electricity
- They are non-sonorous
Chemical properties
- Non-metals form negative ions by gaining or accepting electrons
- They have relatively high electronegativity, meaning they attract electrons in chemical bonds with other elements
- Their oxides tend to be acidic
Difference between Metals and Non-Metals
| Metals | Non-Metals |
|---|---|
| These are solids at room temperature except mercury | These exist in all three states (Solid, Liquid and Gas) |
| These are very hard except sodium | These are soft except diamond |
| These are malleable and ductile | These are brittle and can break down into pieces |
| These are shiny | These are non-lustrous except iodine |
| Electropositive in nature | Electronegative in nature |
| Have high densities | Have low densitie |
| Sonorous | Non Sonorous |
Exceptions
- All metals except mercury exist as solids at room temperature. Metals have high melting. points but gallium and caesium have very low melting points. These two metals will melt if you keep them on your palm.
- Iodine is non metal but lustrous
- Diamonds are hardest natural substance. It is a non-metal with very high melting and boiling point
- Graphite is non metal but good conductor of electricity.
- Alkali metals (lithium, sodium, potassium) are so soft that they can be cut with a knife. They have low densities and low melting points.
CHEMICAL PROPERTIES OF METALS
Reaction with Oxygen
Almost all metals combine with oxygen to form metal oxides.
\[\small\color{orange}\boxed{\text{Metal}+ \text{Oxygen}\rightarrow \text{Metal Oxide}}\]
\[
\begin{aligned}\mathrm{2Cu+O_2}&\rightarrow \mathrm{2CuO}\\
\mathrm{4Al} + \mathrm{3O_2} &\rightarrow \mathrm{2Al_2O_3}
\end{aligned}
\]
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.
Aluminium oxide reacts in the following manner with acids and bases –
\[\scriptsize
\begin{aligned}
\mathrm{Al_2O_3 + 6HCl} &\rightarrow \mathrm{2AlCl_3 + 3H_2O}\\\\
\mathrm{Al_2O_3 + 2NaOH} &\rightarrow \mathrm{\underset{\text{(Sodium aluminate)}}{2NaAlO_2} + H_2O}
\end{aligned}
\]
Most metal oxides are insoluble in water but some of these dissolve in water to form alkalis. Sodium
oxide and potassium oxide dissolve in water to produce alkalis as follows
\[\small
\begin{aligned}
\mathrm{Na_2O(s) + H_2O(l)} &\rightarrow \mathrm{2NaOH(aq)}\\\\
\mathrm{K_2(s) + H_2O(l)} &\rightarrow \mathrm{2KOH(aq)}
\end{aligned}
\]
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.
Reaction with Acids
Metals tend to react with dilute acids like hydrochloric acid to produce a salt and hydrogen gas. For
example:
\[\scriptsize\color{orange}\boxed{\mathrm{Metal + Dilute Acid} \rightarrow \mathrm{Salt + Hydrogen}}\]
Sodium reacts vigorously with dilute HCl to produce sodium chloride and hydrogen gas.
Metals below hydrogen in the reactivity series, such as zinc and iron, also react but less vigorously
Reaction with Water
Metals react with water and produce a metal oxide and hydrogen gas. Metal oxides that are soluble in water dissolves in it to further form metal hydroxide \[\scriptsize \color{orange}\boxed{ \begin{aligned} \mathrm{Metal + Water} &\rightarrow \mathrm{Metal oxide + Hydrogen}\\\\ \mathrm{Metal oxide + Water} &\rightarrow \mathrm{Metal hydroxide} \end{aligned}} \] Metals like aluminium, iron and zinc do not react with either cold or hot water. But they react with steam to form the metal oxide and hydrogen. \[\scriptsize \begin{aligned} \mathrm{2Al(s) + 3H_2O(g)} &\rightarrow \mathrm{Al_2O_3 + 3H_2O(g)}\\\\ \mathrm{3Fe(s) + 4H_2O(g)} &\rightarrow \mathrm{Fe_3O_4 (s) + 4H_2(g)} \end{aligned} \] Metals such as lead, copper, silver and gold do not react with water at all.
Reactions with Solutions of other Metal Salts
\[\scriptsize\color{orange}\boxed{ \mathrm{\text{Metal A} + \text{Salt solution of B} \rightarrow \text{Salt solution of A} + \text{Metal B}}} \] If metal A displaces metal B from its solution, it is more reactive than B.
The Reactivity Series
Reactivity series is a systematic arrangement of metals based on their tendency to undergo
chemical reactions, particularly with water, acids, and other metal compounds.
The series provides a clear comparison of how vigorously different metals react, helping learners
understand why some metals,
like potassium and sodium, are highly reactive, while others, such as gold and platinum, are far less so.
Key points
- The order of common metals in the series, starting from the most reactive (like potassium) down to the least reactive (like gold).
- How the position of a metal in the reactivity series predicts its ability to displace another metal from its salt solution, explaining important processes like displacement reactions.
- The significance of the series in practical activities, such as metal extraction, corrosion, and protection methods.
- Real-world applications and experiments that highlight why metals are used differently in industry and daily life based on their reactivity.
Reactivity Series of Metals (NCERT Class 10)
Ordered List (Most to Least Reactive):
- Potassium (K)
- Sodium (Na)
- Calcium (Ca)
- Magnesium (Mg)
- Aluminium (Al)
- Zinc (Zn)
- Iron (Fe)
- Lead (Pb)
- Copper (Cu)
- Silver (Ag)
- Gold (Au)
Easy-to-Remember Rhyme:
Each word’s first letter stands for a metal:
- Please → Potassium
- Send → Sodium
- Cats → Calcium
- Monkeys → Magnesium
- And → Aluminium
- Zebras → Zinc
- In → Iron
- Lead → Lead
- Cages → Copper
- Securely → Silver
- Guarded → Gold
Reaction of Metals and Non-Metals
Compounds are formed in Metals and Non-Metals by the transfer of electrons from a metal to a non-metal are known as ionic compounds or electrovalent compounds.
Properties of Ionic Compound
- Ionic compounds are crystalline solids
- They have high melting and boiling points
- They are hard and brittle
- They are good conductors of electricity in the molten state and in aqueous solutions
- They are generally soluble in water
OCCURRENCE OF METALS
Extraction of Metals
Some metals are found in the Earth’s crust in the free state. Some are found in the form of their
compounds. The metals at the bottom of the activity series are the least reactive. They are often found
in a free state. For example, gold, silver, platinum and copper are found in the free state. Copper and
silver are also found in the combined state as their sulphide or oxide ores.
The metals at the top of the activity series (K, Na, Ca, Mg and Al) are so reactive that they are never
found in nature as free elements. The metals in the middle of the activity series (Zn, Fe, Pb, etc.) are
moderately reactive. They are found in the earth’s crust mainly as oxides, sulphides or carbonates. You
will find that the ores of many metals are oxides. This is because oxygen is a very reactive element and
is very abundant on the earth.,
Several steps are involved in the extraction of pure metal from ores.
Gangue
Ores mined from the earth are usually contaminated with large amounts of impurities such as soil, sand, etc., called gangue.
Extracting Metals Low in the Activity Series
Metals low in the activity series are very unreactive. The oxides of these metals can be reduced to metals by heating alone. For example, cinnabar (HgS) is an ore of mercury. When it is heated in air, it is first converted into mercuric oxide (HgO). Mercuric oxide is then reduced to mercury on further heating. \[\scriptsize \begin{aligned} \mathrm{2HgS(s) + 3O_2(g)}&\mathrm{\xrightarrow{Heat} 2HgO(s) + 2SO_2(g)}\\\\ \mathrm{2HgO(s)}&\mathrm{\xrightarrow{Heat}2Hg(l) + O_2(g)} \end{aligned} \] Similarly, copper, which is found as \(\mathrm{Cu_2S}\) in nature, can be obtained from its ore by just heating it in air. \[\scriptsize \begin{aligned} \mathrm{2Cu_2S+3O_2(g)}&\mathrm{\xrightarrow{Heat}2Cu_2O(s) + 2SO_2(g)}\\\\ \mathrm{2Cu_2O+Cu_2S}&\mathrm{\xrightarrow{Heat}6Cu(s)+SO_2(g)} \end{aligned} \]
Roasting
Involves heating ores to a temperature below their melting point in the presence of excess air.
Calcination
Involves heating ores below their melting point in the absence of air or with limited air.
Extracting Metals in the Middle of the Activity Series
The metals in the middle of the activity series, such as iron, zinc, lead, and copper, are moderately
reactive. These are usually present as sulphides or carbonates in nature. It is easier to obtain a metal
from its oxide, as compared to its sulphides and carbonates.
Therefore, prior to reduction, the metal sulphides and carbonates must be converted into metal oxides.
The sulphide ores are converted into oxides by heating strongly in the presence of excess air. This
process is known as roasting. The carbonate ores are changed into oxides by heating strongly in limited
air. This process is known as calcination. The chemical reaction that takes place during roasting and
calcination of zinc ores can be shown as follows
Roasting
\[\scriptsize
\mathrm{2ZnS(s) + 3O_2(g)\xrightarrow{\text{Heat}} 2ZnO(s)+2SO_2}
\]
Calcination
\[\small
\mathrm{ZnCO_3(s)\xrightarrow{\text{Heat}}ZnO(s)+CO_2(g)}
\]
The metal oxides are then reduced to the corresponding metals by using suitable reducing agents such as
carbon. For example, when zinc oxide is heated with carbon, it is reduced to metallic zinc.
\[
\mathrm{ZnO(s)+C(s)\rightarrow Zn(s)+CO_2(g)}
\]
Thermite Reaction
The thermite reaction is a chemical reaction between aluminium and iron oxide that produces molten iron,
aluminium oxide, and a large amount of heat. It's an exothermic reaction, which means it releases energy
in the form of heat.
Equation
\[\scriptsize\mathrm{Fe_2O_3 (s) + Al (s) → Al_2O_3 + 2Fe(l) + Heat}\]
Uses
- Welding:
Thermite welding is used to join railway tracks, cracked machine parts, and other metals. - Preparing metals:
Thermite can be used to extract metals from their oxides. - Incendiary devices:
Thermite can be used to produce incendiary weapons.
Extracting Metals towards the Top of the Activity Series
The metals high up in the reactivity series are very reactive. They cannot be obtained from their compounds by heating with carbon. These metals are obtained by electrolytic reduction. For example, sodium, magnesium and calcium are obtained by the electrolysis of their molten chlorides. The metals are deposited at the cathode (the negatively charged electrode), whereas chlorine is liberated at the anode (the positively charged electrode). The reactions are: \[\small \begin{aligned} \text{At Anode: } \mathrm{Na^+ + e^-}&\rightarrow \mathrm{Na}\\\\ \text{At Cathode: } \mathrm{2Cl^-}&\rightarrow \mathrm{Cl_2 +2e^-} \end{aligned} \]
Refining of Metals
Once a metal is extracted from its ore, it is often not pure and may contain various impurities that
need to be removed in order to make the metal suitable for practical use. This step is called
refining.
The process of refining aims to increase the quality of the metal, making it more valuable and
functional for specific applications in industries and daily life.
There are several methods of refining metals, each chosen based on the nature of the metal and the
type
of impurities present. One of the most common techniques is electrolytic refining, where the impure metal is made the anode and a pure strip of the same metal is used as the cathode. Both are immersed
in
a solution containing a suitable salt of the metal. When an electric current is passed through the
solution, pure metal from the anode dissolves and deposits onto the cathode, while impurities either
remain in solution or settle at the bottom as “anode mud.”
Other methods include distillation, liquation, and zone refining, each tailored for specific metals
and
impurities. The result of refining is a metal with high purity, which is essential for making tools,
wires, machinery, and other products that require specific material qualities.
Electrolytic Refining
Electrolytic refining is a specialised technique used to purify metals that have been obtained from ores but still contain certain unwanted impurities. This process takes advantage of electricity to separate pure metal from its contaminants, resulting in a product that meets the strict standards required for usage in industries and technology.
In electrolytic refining, an apparatus is set up where a block of impure metal acts as the anode (positive electrode) and a thin strip of the pure metal serves as the cathode (negative electrode). Both electrodes are immersed in a solution (known as the electrolyte) that contains a suitable salt of the metal being purified, such as copper sulphate solution for copper refining.
When electrical current passes through this setup, atoms of the impure metal at the anode lose electrons and dissolve into the electrolyte. These metal ions then gain electrons at the cathode and deposit there as pure metal. The impurities present in the impure metal either remain dissolved in the solution or fall off and collect at the bottom as a residue known as “anode mud.”
Electrolytic refining is widely used for metals like copper, silver, gold, and nickel, ensuring the removal of even minute traces of unwanted substances. Through this process, industries obtain metals of remarkable purity, crucial for manufacturing fine wires, electrical parts, and advanced machinery where both performance and reliability are essential.
Corrosion and Prevention of Corrosion
Corrosion is a natural but unwanted process where metals gradually deteriorate as a result of chemical reactions with substances present in their environment. One of the most familiar examples is the rusting of iron, which occurs when iron reacts with oxygen and moisture in the air. This rust not only weakens the metal but also spoils its appearance and can ultimately render it useless.
Corrosion does not affect all metals equally—highly reactive metals like iron, zinc, and aluminium are more prone to corrosion, while noble metals like gold and platinum remain unaffected even after years of exposure. The process is slow, yet persistent, and leads to huge economic losses as metallic structures, machinery, pipelines, and everyday items lose their strength and integrity over time.
Prevention Methods:
- Painting or Coating: Covering metal surfaces with paint, oil, or grease creates a shield that blocks exposure to air and moisture.
- Galvanization: Iron or steel objects are coated with a thin layer of zinc. Since zinc is more reactive, it corrodes instead of the underlying metal, protecting it effectively.
- Alloying: Making alloys like stainless steel (a mixture of iron with chromium and nickel) significantly increases resistance to rust.
- Cathodic Protection: Connecting the metallic structure to a more reactive metal, which sacrifices itself by corroding in place of the protected item.
By understanding corrosion and actively employing methods to prevent it, engineers, builders, and even households ensure the longevity and safety of metallic objects, saving resources and minimising waste. This knowledge highlights the practical importance of chemistry in preserving valuable assets and infrastructure in everyday life.
Important Points
- Elements can be classified as metals and non-metals.
- Metals are lustrous, malleable, ductile and are good conductors of heat and electricity. They are solids at room temperature, except mercury, which is a liquid.
- Metals can form positive ions by losing electrons to non-metals.
- Metals combine with oxygen to form basic oxides. Aluminium oxide and zinc oxide show the properties of both basic as well as acidic oxides. These oxides are known as amphoteric oxides.
- Different metals have different reactivities with water and dilute acids.
- A list of common metals arranged in order of their decreasing reactivity is known as an activity series.
- Metals above hydrogen in the Activity series can displace hydrogen from dilute acids
- A more reactive metal displaces a less reactive metal from its salt solution.
- Metals occur in nature as free elements or in the form of their compounds.
- The extraction of metals from their ores and then refining them for use is known as metallurgy.
- An alloy is a homogeneous mixture of two or more metals, or a metal and a non-metal.
- The surface of some metals, such as iron, is corroded when they are exposed to moist air for a long period of time. This phenomenon is known as corrosion.
- Non-metals have properties opposite to those of metals. They are neither malleable nor ductile. They are bad conductors of heat and electricity, except for graphite, which conducts electricity.
- Non-metals form negatively charged ions by gaining electrons when reacting with metals.
- Non-metals form oxides which are either acidic or neutral.
- Non-metals do not displace hydrogen from dilute acids. They react with hydrogen to form hydrides.
