Hello, and welcome to today's chemistry lesson! Today, we are going to explore the fascinating world of hydrogen — the lightest, simplest, and one of the most abundant elements in the entire universe. By the end of this lesson, you will understand where hydrogen comes from, how we prepare it in the laboratory and on an industrial scale, its remarkable properties, and why this gas is considered the fuel of the future. We will also unravel the important concepts of oxidation and reduction, and see how hydrogen plays a central role in these chemical processes.
Let us begin with some essential facts about hydrogen. The symbol for hydrogen is H, and since hydrogen exists as diatomic molecules, its formula is H₂. Hydrogen has a valency of one and an atomic number of 1. Its atomic mass is approximately 1.00794 atomic mass units.
Hydrogen holds a special place in scientific history. It was first prepared by Robert Boyle in 1672, but it was Henry Cavendish who studied its properties in 1776 and called it "inflammable gas" because of its tendency to burn. In 1783, Antoine Lavoisier gave it the name "hydrogen," which literally means "water producer" — derived from the Greek words "hydra" for water and "gen" for producing.
Now, where do we find hydrogen? In the free state, hydrogen is incredibly abundant in the universe — it makes up the outer atmospheres of the sun and stars. However, on Earth, free hydrogen is rare and found only in traces, mainly in volcanic gases. Because hydrogen is highly reactive, it occurs mainly in the combined state.
Water, with the formula H₂O, is the most important compound of hydrogen, covering more than seventy percent of Earth's surface. Hydrogen is also present in acids, alkalis, petroleum, natural gas, and organic compounds. In fact, hydrogen is a key component of all organic compounds, which form the basis of life itself.
Let us now explore how we prepare hydrogen gas. The principal sources on Earth are water, acids, and alkalis.
One major method is the electrolysis of water. When an electric current is passed through acidulated water — that is, water containing a small amount of H₂SO₄ — it decomposes into hydrogen and oxygen. The chemical equation is: H₂O decomposes to give H₂ and O₂. Hydrogen gas collects at the cathode, the negative electrode, while oxygen collects at the anode, the positive electrode. The ratio by volume is two to one — twice as much hydrogen as oxygen. This method is used for large-scale production.
Before we continue, let us understand some key terms related to electrolysis. Electrolysis is defined as a process in which an electric current is passed through an aqueous solution or molten state of a compound, called an electrolyte, to bring about a chemical change.
Electrolytes are compounds that conduct electricity because they contain free mobile ions. Strong electrolytes, like NaCl, NaOH, and strong acids, dissociate almost completely into ions. Weak electrolytes, like ammonium hydroxide and acetic acid, only partially dissociate. Non-electrolytes, such as sugar and pure water, do not conduct electricity at all.
Electrodes are the solid conductors — usually metals or graphite — through which current enters or leaves the electrolyte. The anode is connected to the positive terminal and attracts negative ions called anions. The cathode is connected to the negative terminal and attracts positive ions called cations.
Another common method to prepare hydrogen is by the action of dilute acids on active metals. Metals like magnesium, zinc, and iron react with dilute hydrochloric acid or dilute sulphuric acid to produce hydrogen gas along with a salt.
In the laboratory, we specifically use granulated zinc with dilute hydrochloric acid or dilute sulphuric acid. The reaction with hydrochloric acid produces zinc chloride and hydrogen gas. With sulphuric acid, we get zinc sulphate and hydrogen. Granulated zinc is preferred because it contains copper as an impurity, which acts as a positive catalyst and speeds up the reaction.
The apparatus consists of a round-bottom flask containing zinc granules, fitted with a thistle funnel and delivery tube. When acid is added, effervescence occurs as hydrogen gas is liberated. The gas is collected by downward displacement of water — hydrogen rises and pushes water out of an inverted jar. We discard the first few bubbles as they may contain air, and we never bring a flame near the apparatus because hydrogen is highly inflammable.
Why do we not use dilute nitric acid? Nitric acid is a strong oxidising agent, even when dilute. It oxidises the hydrogen gas into water, preventing its collection.
That is why hydrochloric and sulphuric acids are preferred.
Zinc is particularly suitable for laboratory preparation because other metals have drawbacks. Sodium and potassium react too violently. Calcium and magnesium are expensive. Aluminium forms a protective oxide layer. Iron needs heating and gives a reversible reaction. Lead forms insoluble salts that stop the reaction. Copper, mercury, and silver are below hydrogen in the reactivity series and cannot displace hydrogen from acids.
Hydrogen can also be prepared by the action of water or steam on metals. Very active metals like sodium, potassium, and calcium liberate hydrogen from cold water, forming metal hydroxides. Sodium and potassium react violently, so they are stored in kerosene oil. Calcium reacts more slowly.
Magnesium liberates hydrogen from boiling water, though the reaction is slow. When steam is passed over heated magnesium, zinc, aluminium, or iron, hydrogen is produced along with the metal oxide. The reaction between iron and steam is particularly important but reversible.
For large-scale manufacture, we use the Bosch process. This involves three main steps.
First, steam is passed over hot coke at one thousand degrees Celsius to produce water gas — a mixture of carbon monoxide and hydrogen. This is an endothermic reaction.
Second, water gas is mixed with more steam and passed over ferric oxide with chromium trioxide as a promoter at four hundred fifty degrees Celsius. This converts carbon monoxide to carbon dioxide and produces more hydrogen. This step is exothermic.
Third, the mixture is passed through water under pressure to dissolve carbon dioxide, or through caustic potash solution. Finally, any remaining carbon monoxide is removed by passing the gas through ammoniacal cuprous chloride solution. Pure hydrogen gas is thus obtained.
Now let us examine the properties of hydrogen.
Physically, hydrogen is a colourless, odourless, and tasteless gas that is non-poisonous. It is sparingly soluble in water — only about twenty millilitres dissolve in one litre of water at ordinary conditions. Hydrogen is extremely difficult to liquefy, requiring temperatures of minus two hundred forty degrees Celsius and twenty atmospheres of pressure. Most remarkably, hydrogen is the lightest of all gases — air is fourteen point four times heavier than hydrogen. This is why hydrogen-filled balloons rise rapidly in air.
Chemically, hydrogen is neutral to litmus — it does not change the colour of either red or blue litmus paper. It is combustible but does not support combustion. When pure hydrogen burns in air, it produces a pale blue flame silently. Ordinary hydrogen, containing impurities, burns with a characteristic pop sound — this is actually a test for hydrogen.
When hydrogen burns in oxygen, it forms water with the release of tremendous energy. This highly exothermic reaction is used to propel space rockets, where liquid hydrogen and liquid oxygen are stored as fuel. However, hydrogen forms an explosive mixture with air, so it must be handled with extreme care.
In diffused sunlight, hydrogen combines with chlorine to form hydrogen chloride gas. Direct sunlight causes an explosive reaction, so we always use diffused light.
With nitrogen, hydrogen produces ammonia in the presence of iron catalyst at four hundred fifty degrees Celsius and two hundred atmospheres — this is the Haber process. With boiling sulphur, hydrogen forms hydrogen sulphide gas, which has a rotten egg smell.
When heated with reactive metals like sodium and calcium, hydrogen forms metal hydrides. Most metal hydrides are unstable.
One of the most important chemical properties of hydrogen is its ability to act as a reducing agent. When passed over hot metal oxides like copper oxide, lead oxide, or iron oxide, hydrogen removes oxygen and reduces them to the free metals. For example, copper oxide plus hydrogen gives copper plus water. Note that oxides of highly active metals like sodium, potassium, and aluminium cannot be reduced by hydrogen.
Let us now understand oxidation and reduction — concepts closely linked to hydrogen.
Oxidation is defined as a chemical process involving addition of oxygen or removal of hydrogen from a substance. For example, when carbon burns in oxygen to form carbon dioxide, carbon is oxidised. When hydrogen sulphide reacts with chlorine, hydrogen is removed and sulphur is oxidised.
An oxidising agent is the substance that supplies oxygen or removes hydrogen. During the reaction, the oxidising agent itself gets reduced.
Reduction is the reverse process — addition of hydrogen or removal of oxygen from a substance. When nitrogen reacts with hydrogen to form ammonia, nitrogen is reduced. When mercuric oxide is heated to give mercury and oxygen, the mercury is reduced.
A reducing agent supplies hydrogen or removes oxygen. Hydrogen itself is an excellent reducing agent. During reduction, the reducing agent gets oxidised.
Redox reactions are those in which oxidation and reduction occur simultaneously. One substance is reduced while another is oxidised. For example, when hydrogen reduces copper oxide to copper, the copper oxide is reduced while hydrogen is oxidised to water.
Finally, let us look at the important uses of hydrogen.
First, hydrogen and oxygen together produce an oxy-hydrogen flame with temperatures between two thousand eight hundred and three thousand degrees Celsius — ideal for cutting and welding metals.
Second, hydrogen serves as a fuel. Its high heat of combustion makes it valuable as coal gas, water gas, and especially as liquid hydrogen for rocket fuel. Importantly, hydrogen is pollution-free — its only combustion product is water. This is why it is seen as the mass fuel of the future.
Third, hydrogen is used in the hydrogenation of vegetable oils. When hydrogen is passed through oils like groundnut or coconut oil in the presence of nickel, platinum, or palladium catalyst, the liquid oils solidify into vanaspati ghee. This process is called catalytic hydrogenation.
Fourth, hydrogen is essential for manufacturing chemicals — ammonia for fertilisers, hydrochloric acid, methanol, and urea.
Fifth, as a reducing agent, hydrogen helps extract less reactive metals like copper, lead, and tin from their oxides.
Sixth, hydrogen was once used to fill weather balloons, though it has now been replaced by helium due to safety concerns.
Let us quickly recap the key points from today's lesson.
First, hydrogen is the lightest element with symbol H and formula H₂, and it is the most abundant element in the universe.
Second, hydrogen is prepared by electrolysis of water, by the action of dilute acids on metals like zinc, and on a large scale by the Bosch process.
Third, hydrogen is a colourless, odourless, combustible gas that does not support combustion, and it forms an explosive mixture with air.
Fourth, hydrogen acts as a reducing agent, removing oxygen from metal oxides.
Fifth, oxidation involves addition of oxygen or removal of hydrogen, while reduction involves addition of hydrogen or removal of oxygen — these processes occur simultaneously in redox reactions.
Sixth, hydrogen has diverse uses including fuel, metal cutting, hydrogenation of oils, and manufacture of chemicals.
That brings us to the end of our lesson on hydrogen. You have learned about an element that powers stars, forms water, and may one day power our vehicles and homes. Keep exploring, stay curious, and remember — the simplest elements often hold the greatest potential. Thank you for listening, and I look forward to our next chemistry adventure together!