Study of Hydrogen

Hydrogen sits first in the periodic table: atomic number 1, with just one proton and one electron. It is the lightest of all substances and the most abundant element in the universe, yet on Earth it is found mostly combined - in water, acids and almost every living thing. As a free gas it is colourless, odourless, tasteless and only very slightly soluble in water. Because a single electron can be lost (like the alkali metals) or gained (like the halogens), hydrogen shows a dual nature and is given a special place at the top of the table.

The usual laboratory method reacts granulated zinc with dilute sulphuric acid: Zn + H2SO4 -> ZnSO4 + H2. The reaction is steady at room temperature, which is why zinc is chosen rather than a violent metal like sodium. Which metal reacts with what depends on the activity series. Very reactive metals (potassium, sodium, calcium) liberate hydrogen even from cold water; moderately reactive metals (magnesium, aluminium, zinc, iron) need a dilute acid; and metals below hydrogen, such as copper and silver, do not displace hydrogen from acids at all. A few amphoteric metals - zinc, aluminium and lead - are special: they give hydrogen with dilute acids and also with hot concentrated alkali.

Hydrogen is collected by the downward displacement of water because it is almost insoluble and lighter than air; the gas bubbles up and pushes water out of an inverted jar. If a dry sample is needed it is passed through a drying agent such as anhydrous calcium chloride. The first few jars are discarded because they contain air mixed with hydrogen, and a hydrogen-air mixture is explosive - so the gas is never lit until it has been tested and is pure.

Hydrogen burns in air with a pale-blue, nearly invisible flame, combining with oxygen to form water - a strongly exothermic reaction (the oxy-hydrogen flame reaches very high temperatures used in welding). Its most important property is being a reducing agent: when dry hydrogen is passed over a heated oxide of a less active metal, it pulls the oxygen away. Black copper(II) oxide turns into red-brown copper, and lead(II) oxide into grey lead, while water forms: CuO + H2 -> Cu + H2O. Hydrogen cannot, however, reduce the oxides of very active metals such as sodium, calcium or aluminium, which hold their oxygen too firmly.