If acids are the sour side of chemistry, bases are the bitter, slippery side. That characteristic soapy feel when you rub baking soda between your fingers? That is a base dissolving the thin layer of oils on your skin. Bases are defined as proton acceptors (Bronsted-Lowry definition) or hydroxide donors (Arrhenius definition). When sodium hydroxide (NaOH) dissolves in water, it releases OH⁻ ions that gobble up any free H⁺ ions, raising the pH above 7. This is the chemical opposite of what acids do, which is why mixing an acid with a base neutralizes both — producing water and a salt in one of chemistry's cleanest reactions.
Bases are everywhere in daily life. Baking soda (sodium bicarbonate, NaHCO₃) neutralizes stomach acid and makes cakes rise by releasing CO₂ gas when it reacts with acidic ingredients. Soap has been made for thousands of years by reacting fats with a strong base — traditionally lye (NaOH). Ammonia (NH₃) in household cleaners is a weak base that dissolves grease. Antacid tablets are bases (calcium carbonate, magnesium hydroxide) designed to neutralize excess stomach acid. Even your blood is slightly basic, maintained at a remarkably stable pH of 7.35-7.45 by a sophisticated buffering system — if blood pH drops below 7.0 or rises above 7.8, it becomes life-threatening.
Strong bases like sodium hydroxide (caustic soda) and potassium hydroxide (caustic potash) are as corrosive as strong acids and just as important industrially. Sodium hydroxide is used to manufacture paper, textiles, soap, and biodiesel. Calcium hydroxide (slaked lime) treats water and reduces soil acidity for farming. Concrete, the most widely used material on Earth after water, relies on calcium hydroxide chemistry to set and harden. The word "alkali" — an older name for soluble bases — comes from the Arabic "al-qali," meaning "the calcined ashes," because early chemists extracted bases from wood ash.