Base (chemistry) - Fundamental Concepts of Bases

Definitions of Bases Understanding What a Base Is Bases are one of the most important chemical concepts you'll encounter. However, the definition of a base has evolved over time as chemistry advanced. Understanding each definition will help you tackle different types of chemistry problems. Arrhenius Definition The Arrhenius definition was historically the first and is still useful for introductory chemistry. An Arrhenius base is a substance that dissociates in aqueous solution to produce hydroxide ions ($\text{OH}^-$). The classic example is sodium hydroxide: $$\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$$ When a base releases hydroxide ions into water, it lowers the concentration of hydrogen ions, which raises the pH above 7. This is why Arrhenius bases feel slippery and taste bitter—properties that early chemists used to identify them. However, the Arrhenius definition has limitations. It only works in aqueous solutions and requires the formation of hydroxide ions. This is why scientists developed broader definitions. Brønsted–Lowry Definition The Brønsted–Lowry definition is more general and works in any solvent, not just water. A Brønsted–Lowry base is a substance that can accept a proton (hydrogen cation, $\text{H}^+$) from another substance. This definition explains why hydroxide ions work as bases: they accept a proton to form water: $$\text{OH}^- + \text{H}^+ \rightarrow \text{H}2\text{O}$$ A crucial example is ammonia ($\text{NH}3$). Ammonia doesn't contain hydroxide ions, yet it acts as a base. It accepts a proton from water: $$\text{NH}3 + \text{H}2\text{O} \rightleftharpoons \text{NH}4^+ + \text{OH}^-$$ Notice that this reaction produces hydroxide ions indirectly, but ammonia itself is the proton acceptor. This is why the Brønsted–Lowry definition is more useful for understanding a broader range of bases. Lewis Definition The Lewis definition is the most general and powerful. A Lewis base is an electron-pair donor—a substance that can share a pair of electrons with an electron acceptor (called a Lewis acid). This definition includes many substances that wouldn't be considered bases under the other definitions. For example, the carbonate ion ($\text{CO}3^{2-}$) can donate electron pairs to boron trifluoride ($\text{BF}3$): $$\text{CO}3^{2-} + \text{BF}3 \rightarrow \text{(CO}3\text{)(BF}3)$$ The advantage of the Lewis definition is that it doesn't require protons or hydroxide ions at all. It applies to bases in non-aqueous solutions and explains basic behavior in terms of fundamental electron chemistry. Key insight: Each definition is nested within the others. Every Brønsted–Lowry base is a Lewis base, and most Arrhenius bases are also Brønsted–Lowry bases. The definitions become progressively more general. General Properties of Bases Now that you understand what bases are, let's examine their observable properties. These characteristics will help you identify bases in practical situations. Electrical Conductivity Aqueous solutions of bases conduct electricity. This is because bases dissociate into ions in water—for example, sodium hydroxide breaks apart into Na$^+$ and OH$^-$ ions. These mobile ions carry electrical current through the solution. Indicator Reactions Scientists use indicators—dyes that change color in the presence of bases—to identify bases and measure their strength. You should memorize these color changes, as they appear frequently on exams: Litmus paper: Red litmus turns blue in basic solutions (this is why "litmus test" means a definitive test) Phenolphthalein: Colorless in neutral solutions but turns bright pink in basic solutions Bromothymol blue: Remains blue in basic solutions Methyl orange: Turns yellow in basic solutions These indicators work because their molecular structure changes in response to pH, causing different wavelengths of light to be absorbed. pH Scale The pH scale measures the concentration of hydrogen ions in solution. A pH greater than 7 indicates a basic (alkaline) solution. Recall that pH is defined as: $$\text{pH} = -\log[\text{H}^+]$$ In basic solutions, the concentration of hydroxide ions is higher than hydrogen ions, resulting in pH > 7. For reference, a neutral solution (pH = 7) has equal concentrations of H$^+$ and OH$^-$. Classification of Bases Bases aren't all equally strong. Understanding the difference between strong and weak bases is critical for predicting reactions and calculating equilibrium concentrations. Strong Bases A strong base is one that completely ionizes (or protonates) in water. When a strong base dissolves, it fully dissociates into its ions—essentially 100% of the molecules break apart. The common strong bases you must know are: Group 1 hydroxides (alkali metals): LiOH, NaOH, KOH, RbOH, CsOH Group 2 hydroxides (alkaline-earth metals): Ca(OH)$2$, Sr(OH)$2$, Ba(OH)$2$ For example, when sodium hydroxide dissolves: $$\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$$ This reaction goes essentially to completion. If you dissolve 0.1 moles of NaOH in water, you'll get 0.1 moles of OH$^-$ ions. Why are these bases so strong? The leveling effect explains this: water is such a strong proton donor that it fully converts any strong base to its conjugate acid. Since the conjugate acid of OH$^-$ is water itself (a weak acid), the reaction is essentially irreversible. Weak Bases A weak base does not fully ionize in water, so ionization is incomplete. Only a small fraction of the base molecules actually accept protons. The classic example is ammonia (NH$3$): $$\text{NH}3 + \text{H}2\text{O} \rightleftharpoons \text{NH}4^+ + \text{OH}^-$$ Notice the equilibrium arrow (⇌), not a single arrow. This indicates that the reaction doesn't go to completion. At 25°C, the equilibrium constant for this reaction is: $$Kb = 1.8 \times 10^{-5}$$ This small value means the equilibrium strongly favors the reactants. If you dissolve 0.1 moles of ammonia, only a tiny fraction converts to NH$4^+$. The key difference: Strong bases produce a high concentration of OH$^-$ ions; weak bases produce a low concentration. This is why strong bases have higher pH values than weak bases of the same molarity. <extrainfo> Superbases Superbases are bases stronger than the hydroxide ion. Examples include alkoxide ions (RO$^-$) and amide ions (NH$2^-$). These bases cannot exist in aqueous solution because water is basic enough to protonate them, forming hydroxide ions instead. This is another manifestation of the leveling effect: water "levels" all strong bases to the strength of hydroxide. </extrainfo> Monoprotic and Polyprotic Bases Some bases can accept multiple protons. The terminology describes how many hydroxide ions a base releases: A monoprotic base releases one hydroxide ion per formula unit when it completely ionizes. Ammonia (NH$3$ → NH$4^+$ + OH$^-$) is monoprotic. A polyprotic base (specifically, a diprotic base) releases two hydroxide ions. Calcium hydroxide (Ca(OH)$2$) and barium hydroxide (Ba(OH)$2$) are diprotic strong bases because they contain two hydroxide ions in their formula. When these dissociate in water: $$\text{Ca(OH)}2 \rightarrow \text{Ca}^{2+} + 2\text{OH}^-$$ For every mole of Ca(OH)$2$ that dissolves, you get two moles of OH$^-$. This is why even though Ca(OH)$2$ is slightly less soluble than NaOH, it can still produce high concentrations of hydroxide in saturated solutions.