Physical Properties and Major Exceptions of Metals and Non-metals
| Property | Metals (Typical) | Exceptions | Nonmetals (Typical) | Exceptions |
|---|---|---|---|---|
| State at Room Temperature | Solid | Mercury (liquid), Gallium (liquid above 30°C) | Gases, some solids | Bromine (liquid) |
| Appearance (Lustre) | Lustrous (shiny) | Sodium dull, some metals not shiny | Dull | Iodine and graphite (lustrous) |
| Hardness | Hard | Sodium, Potassium, Lithium (soft) | Brittle (solids) | Diamond (very hard), Carbon fibers |
| Density | Usually high | Lithium, Sodium, Potassium (low density) | Low density | Iodine, Diamond (high density) |
| Melting/Boiling Point | Very high | Gallium, Caesium, Sodium, Potassium (low) | Low | Diamond, Graphite, Carbon, Silicon, Boron (high) |
| Malleability/Ductility | Malleable, ductile | Zinc, Mercury (not malleable/ductile) | Non-malleable, non-ductile | Carbon fibers ductile |
| Conductivity (Electric) | Good conductors | Lead, Mercury (poor), Tungsten (less conductive) | Poor conductors | Graphite (good conductor) |
| Conductivity (Heat) | Good conductors | Lead, Mercury (poor) | Poor conductors | Diamond, graphite (good for heat) |
| Sonority | Sonorous | Mercury (not sonorous) | Non-sonorous | None |
Chemical Properties of Metals
What happens when Metals are burnt in Air?
When metals are burnt in air, they react with oxygen present in the air to form metal oxides. This is a chemical reaction known as combustion.
For example, when magnesium burns, it reacts with oxygen to form magnesium oxide, which is a white powder:
2Mg + O2 → 2MgO
During this reaction, heat and light are produced, often visible as a bright flame or spark. Metal oxides are generally solid and have different properties than the original metal.
The burning of metals is an oxidation process where metals lose electrons to oxygen atoms, forming ionic compounds called metal oxides.
Note: Different metals burn with different flame colors and produce different types of oxides.
What happens when Metals react with Water?
When metals react with water, they generally form metal hydroxides and release hydrogen gas. The nature and speed of the reaction depend on the metal's reactivity.
General Reaction Formula:
Metal + Water → Metal Hydroxide + Hydrogen gas
Example:
2Na + 2H2O → 2NaOH + H2↑
- Highly reactive metals (like sodium and potassium) react vigorously and instantly with cold water to produce metal hydroxides and hydrogen gas. This reaction releases a lot of heat and hydrogen gas may ignite.
- Moderately reactive metals (like magnesium and calcium) react slowly with cold water or quickly with hot water/steam, producing metal hydroxides or oxides and hydrogen gas.
- Less reactive metals (like iron, zinc) react only with steam, not cold water.
- Unreactive metals (like copper, silver, gold) do not react with water.
Reactivity of metals with water and products
| Metal | Reaction with Water | Products Formed | Reaction Type |
|---|---|---|---|
| Potassium (K) | Reacts violently with cold water | Potassium hydroxide (KOH) + Hydrogen gas (H₂) | Vigorous |
| Sodium (Na) | Reacts quickly with cold water | Sodium hydroxide (NaOH) + Hydrogen gas (H₂) | Rapid |
| Calcium (Ca) | Reacts less strongly with cold water | Calcium hydroxide (Ca(OH)₂) + Hydrogen gas (H₂) | Moderate |
| Magnesium (Mg) | Reacts very slowly with cold water; reacts faster with steam | Magnesium hydroxide (Mg(OH)₂) + Hydrogen gas (H₂) (with water) Magnesium oxide (MgO) + Hydrogen gas (H₂) (with steam) |
Slow (cold water), Fast (steam) |
| Zinc (Zn) | Reacts only slowly with steam | Zinc oxide (ZnO) + Hydrogen gas (H₂) | Slow |
| Iron (Fe) | Reacts only slowly with steam | Iron oxide (Fe₃O₄) + Hydrogen gas (H₂) | Slow |
| Copper (Cu) | No reaction with water or steam | None | None |
Reaction of metals with cold water vs steam vs acids
| Type of Metal | Reaction with Cold Water | Reaction with Steam | Reaction with Dilute Acids | Example Metals |
|---|---|---|---|---|
| Highly Reactive Metals | React vigorously | React vigorously | React vigorously | Potassium (K), Sodium (Na), Calcium (Ca) |
| Moderately Reactive Metals | React slowly or no reaction | React readily | React readily | Magnesium (Mg), Aluminium (Al), Zinc (Zn) |
| Less Reactive Metals | No reaction | React slowly | React slowly | Iron (Fe), Lead (Pb), Tin (Sn) |
| Least Reactive Metals | No reaction | No reaction | No reaction | Copper (Cu), Silver (Ag), Gold (Au) |
What Happens When Metals React with Acids?
When metals react with acids, they usually produce a salt and hydrogen gas (H2). This reaction is common for metals that are more reactive than hydrogen, as these metals displace hydrogen from the acid.
General Equation:
Metal + Acid → Salt + Hydrogen gas
Example Chemical Equation:
Zn + 2HCl → ZnCl2 + H2↑
In this example, zinc reacts with hydrochloric acid to form zinc chloride (a salt) and hydrogen gas, which can be seen as bubbles or fizzing.
The salt formed depends on the acid used. For example:
- Hydrochloric acid forms metal chlorides
- Sulfuric acid forms metal sulfates
- Nitric acid forms metal nitrates
More reactive metals react more vigorously with acids, while less reactive metals react slowly or not at all.
Reaction of Metals with Solutions of Other Metal Salts
When a more reactive metal is put into a solution of the salt of a less reactive metal, it displaces the less reactive metal from the solution. This is called a displacement reaction.
General reaction:
Metal A + Salt solution of Metal B → Salt solution of Metal A + Metal B
Example:
Zn (s) + CuSO4 (aq) → ZnSO4 (aq) + Cu (s)
In this reaction, zinc displaces copper from copper sulfate solution forming zinc sulfate and solid copper.
This reaction shows that zinc is more reactive than copper.
Reactivity Series of Metals
The reactivity series is a list of metals arranged from most reactive to least reactive. It helps predict how metals react with water, acids, and other substances.
| Position | Metal | Reactivity |
|---|---|---|
| 1 | Potassium (K) | Most Reactive |
| 2 | Sodium (Na) | Highly Reactive |
| 3 | Lithium (Li) | Highly Reactive |
| 4 | Calcium (Ca) | Highly Reactive |
| 5 | Magnesium (Mg) | Moderately Reactive |
| 6 | Aluminium (Al) | Moderately Reactive |
| 7 | Zinc (Zn) | Moderately Reactive |
| 8 | Iron (Fe) | Less Reactive |
| 9 | Lead (Pb) | Less Reactive |
| 10 | Hydrogen (H) | Reference Point |
| 11 | Copper (Cu) | Low Reactivity |
| 12 | Silver (Ag) | Low Reactivity |
| 13 | Gold (Au) | Least Reactive |
| 14 | Platinum (Pt) | Least Reactive |
Key Points about the Reactivity Series:
- Metals at the top are the most reactive (easily lose electrons and react strongly with substances like water, acids, and oxygen).
- Metals at the bottom are the least reactive (do not react easily, even with acids or water).
- The series helps predict how metals react with water, acids, and other metal salt solutions.
- It also determines whether a metal can displace another from its compound in solution.
How do Metals and Non Metals React
Metals and non-metals react differently due to their distinct chemical nature. Metals and non-metals react when electrons transfer from the metal to the non-metal, forming positively charged metal ions (cations) and negatively charged non-metal ions (anions). This electron transfer results in a strong electrostatic attraction between the opposite charges, creating an ionic bond and forming a stable ionic compound.
Formation of Ionic Compounds:
When a metal loses electrons and a non-metal gains them, oppositely charged ions are formed. These ions then attract each other electrostatically, similar to magnets, to form a strong chemical bond known as an ionic bond. The resulting compound, held together by ionic bonds, is called an ionic compound.
Example:
Sodium chloride (NaCl): Sodium (a metal) loses an electron to become a positive sodium ion (Na⁺). Chlorine (a non-metal) gains that electron to become a negative chloride ion (Cl⁻). The Na⁺ and Cl⁻ ions are attracted to each other, forming the ionic compound sodium chloride.
Other Reactions of Metals and Non-Metals
Beyond reacting with each other, metals and non-metals exhibit vastly different behaviors when exposed to common reactants like oxygen, water, and acids.
| Reaction Type | Metals | Non-Metals |
|---|---|---|
| With Oxygen | Form basic or amphoteric oxides (e.g., MgO) | Form acidic or neutral oxides (e.g., CO2, SO2) |
| With Water | Form metal hydroxides + H2 gas (e.g., 2Na + 2H2O → 2NaOH + H2) | Generally no reaction (except chlorine forms acidic solution with water) |
| With Acids | Form salt + H2 gas (e.g., Zn + 2HCl → ZnCl2 + H2) | Do not react to release H2 gas |
| With Salt Solutions | More reactive metal displaces less reactive metal | Gain electrons to form ionic compounds with metals |
| With Other Non-Metals | -- | Form covalent compounds by sharing electrons |
Properties of Ionic Compounds
Ionic compounds have distinct properties due to the strong electrostatic forces between oppositely charged ions. Some key properties of ionic compounds are given below:
- High Melting and Boiling Points: Ionic compounds have high melting and boiling points because a large amount of energy is required to break the strong ionic bonds between the ions.
- Hardness and Brittleness: Ionic compounds are hard but brittle. When force is applied, layers of ions shift, and ions of the same charge align, causing repulsion and the crystal to shatter.
- Electrical Conductivity:
- In solid state, ionic compounds do not conduct electricity because ions are fixed in place.
- When melted or dissolved in water, they conduct electricity well as ions are free to move and carry charge.
- Solubility: Ionic compounds are generally soluble in polar solvents like water because water molecules can surround and separate the ions.
- Crystal Lattice Structure: Ionic compounds form a highly ordered three-dimensional crystal lattice structure that maximizes the attraction between opposite charges and minimizes repulsion.
Occurrence of Metals
Most metals are found in the Earth's crust in the form of combined compounds called minerals. These minerals may be oxides, sulfides, carbonates, or other compounds. Only a few metals such as gold, silver, copper, and platinum occur in their native (free) state because they are less reactive and resist natural oxidation. Highly reactive metals like sodium, potassium, magnesium, and calcium occur only in combined states because they readily react with moisture, oxygen, and carbon dioxide in the environment.
An ore is a mineral from which a metal can be extracted economically. For example, bauxite is the ore of aluminum, while iron is commonly extracted from hematite and magnetite ores.
Extraction of Metals
Extraction involves separating the metal from its ore. Since metals mostly occur in combined states, extraction requires chemical processes such as roasting, reduction, or electrolysis depending on the metal's reactivity and ore type. This industrial process is called metallurgy.
Enrichment of Ores
After mining, ores contain impurities called gangue (rock or earthy materials). Enrichment involves removing this gangue to increase the concentration of the metal in the ore. Common methods include:
- Handpicking for large visible impurities
- Washing or gravity separation for lighter gangue
- Froth flotation for sulfide ores
- Magnetic separation for magnetic ores
Enriched ore is then suitable for further extraction processes.
Extracting Metals Low in the Activity Series
Metals like gold, silver, and platinum are found in their native state and do not easily form compounds. They are extracted by physical means such as:
- Heating (casting and melting)
- Chemical reactions if needed for impurities
These metals do not require reducing agents due to their low reactivity.
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.
2HgS(s) + 3O2(g) → 2HgO(s) + 2SO2(g) → Heat
2HgO(s) → 2Hg(l) + O2(g)
Extracting Metals in the Middle of the Activity Series
Metals such as zinc, iron, lead, and copper are moderately reactive. They are commonly extracted from their oxides or sulfide ores by the following steps:
Roasting
Heating sulfide ores in the presence of excess air to convert them into oxides and release sulfur dioxide gas.
2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2(g)
Reduction
Reducing the metal oxide to free metal usually using carbon (coke) or carbon monoxide as the reducing agent.
ZnO(s) + C(s) → Zn(s) + CO(g)
For example, iron is extracted in a blast furnace by reducing iron oxides with carbon.
Extracting Metals towards the Top of the Activity Series
Highly reactive metals like potassium, sodium, calcium, magnesium, and aluminium cannot be reduced by carbon due to their strong affinity for oxygen. They are extracted by:
- Electrolysis of their molten salts (e.g., electrolytic reduction of molten sodium chloride)
At cathode: Na+ + e– → Na
At anode: 2Cl– → Cl2 + 2e–
This method requires large amounts of electricity but is essential for these active metals.
Refining of Metals
Refining is the process of removing impurities from crude metals to obtain pure metals suitable for use.
- Distillation: Used for metals with low boiling points like mercury and zinc. The impure metal is heated to vaporize the pure metal, which is then condensed, leaving impurities behind.
- Liquation: Suitable for metals with low melting points such as tin and lead. The crude metal is heated just above its melting point. Pure metal melts and flows away from solid impurities.
- Electrolytic Refining: The most common method where the impure metal acts as the anode and a pure metal strip as the cathode in an electrolyte solution containing metal salts. The pure metal deposits on the cathode while impurities either remain in solution or form anode mud. Example: copper refining.
- Zone Refining: Used for very high-purity metals like silicon, germanium. A molten zone passes through the metal rod, carrying impurities along and leaving behind pure metal.
- Vapour Phase Refining: Metal is converted into a volatile compound which decomposes to give pure metal. Used for metals like nickel.
- Fire Refining: Impurities are oxidized and removed as slag or gases by blowing oxygen into the molten metal.
Corrosion
Corrosion is the natural and irreversible process where metals deteriorate due to chemical or electrochemical reaction with the surrounding environment, forming oxides or other compounds that weaken the metal.
Examples
- Iron when exposed to moist air for a long time acquires a coating of a brown flaky substance called rust.
- Silver articles become black after some time when exposed to air. This is because it reacts with sulphur in the air to form a coating of silver sulphide.
- Copper reacts with moist carbon dioxide in the air and slowly loses its shiny brown surface and gains a green coat. This green substance is basic copper carbonate.
The Electrochemical Mechanism of Rusting (Anode/Cathode)
Experimental Setup: Conditions Necessary for Rusting (Air & Water)
Causes of Corrosion
- Moisture and oxygen exposure
- Corrosive gases (chlorine, sulfur oxides)
- Contact with acidic or salty environments
- Electrical currents and mechanical stress
- Dirt and bacterial contamination
Effects
- Weakens metal structures
- Loss of strength and conductivity
- Rust formation (e.g., iron oxide in iron)
- Economic and safety hazards
Common Types of Corrosion
- General (uniform) corrosion
- Galvanic corrosion (between different metals)
- Pitting corrosion (localized holes)
- Stress corrosion cracking
- Erosion corrosion
Prevention of Corrosion
Corrosion of metals can be prevented by various methods that protect the metal surface or slow down the chemical reactions causing corrosion. Here are some common and effective methods for preventing corrosion:
- Protective Coatings: Paint, grease, oil, and plastic coatings create barriers against moisture and air.
- Galvanization: Coating iron/steel with zinc to protect it from corrosion.
- Alloying: Forming alloys like stainless steel that resist corrosion.
- Cathodic Protection: Using a sacrificial more reactive metal to protect the main metal.
- Electroplating: Applying a thin metal coating to prevent corrosion.
- Corrosion Inhibitors: Chemicals added to slow down corrosion reactions.
- Design and Maintenance: Avoiding water traps and regular cleaning and maintenance.