Acids and Bases

Overview and Role of Acids and Bases

Acid–base reactions are a central type of proton transfer process. In such reactions, particles exchange protons $H^+$, leading to characteristic changes in composition, structure, and properties of substances. This chapter introduces the general idea of acids and bases, builds the bridge to the detailed theories discussed in the subsequent sections, and highlights why acid–base chemistry is so important in science, technology, and everyday life.

Phenomenological Description: How Acids and Bases Behave

Before using precise theoretical definitions, acids and bases can be recognized by simple, observable properties.

Typical Properties of Acids

Common aqueous acids (acidic solutions) often show:

• Taste: sour (e.g., citric acid in lemons, acetic acid in vinegar)
• Effect on indicators:
• Blue litmus paper turns red
• Universal indicator turns red, orange, or yellow depending on strength
• Conductivity: aqueous acid solutions conduct electricity (they contain ions)
• Reactivity with metals: many acids react with active metals (e.g., zinc) to form salts and hydrogen gas:
  $$\text{Zn} + 2\,\text{H}^+ \rightarrow \text{Zn}^{2+} + \text{H}_2 \uparrow$$
• Reactivity with carbonates: acids react with carbonates (e.g., $\text{CaCO}_3$) to produce $CO_2$ gas:
  $$\text{CaCO}_3 + 2\,\text{H}^+ \rightarrow \text{Ca}^{2+} + \text{CO}_2 \uparrow + \text{H}_2\text{O}$$

Typical Properties of Bases

Common aqueous bases (basic or alkaline solutions) often show:

• Taste: bitter (e.g., sodium bicarbonate has a mild basic taste; strong bases are not safe to taste)
• Feel: slippery or soapy on the skin (due to reaction with skin lipids and proteins)
• Effect on indicators:
• Red litmus paper turns blue
• Universal indicator turns blue or violet
• Conductivity: aqueous base solutions conduct electricity
• Reactivity with fats: strong bases (such as NaOH) react with fats to form soap-like substances (saponification)

Acids and Bases as Complementary Opposites

Acids and bases show a characteristic neutralizing behavior when combined:

• Mixing an acidic and a basic solution in suitable amounts often leads to:
• Loss of sour or slippery character
• Indicator colors shifting toward “neutral”
• Formation of salt and water

Neutralization is treated in detail later; here it is sufficient to see that acids and bases can cancel each other’s key properties.

From Behavior to Definitions: Conceptual Framework

Several theoretical models formalize what “acid” and “base” mean. In this chapter, they are only sketched to provide orientation; later sections treat each theory precisely and with examples.

Arrhenius View (Restricted to Aqueous Solutions)

The Arrhenius concept is a simple, classical description:

• Arrhenius acid: substance that increases the concentration of $H^+$ (often present as $\text{H}_3\text{O}^+$) in water
• Arrhenius base: substance that increases the concentration of $\text{OH}^-$ in water

This is useful for many everyday acids and bases (e.g., HCl, NaOH) but does not describe all acid–base reactions, especially those outside water.

Brønsted–Lowry View: Proton Transfer

The Brønsted acid–base theory generalizes the Arrhenius idea and is centered on proton transfer:

• Acid: proton donor
• Base: proton acceptor

Every acid–base reaction is seen as a transfer of $H^+$ from an acid to a base. This is the core model for most of this chapter’s topic and applies to many solvents, not just water.

Lewis View: Electron Pair Interactions

The Lewis acid–base theory focuses on electron pairs:

• Lewis acid: electron pair acceptor
• Lewis base: electron pair donor

This view is more general and connects acid–base behavior with bonding in inorganic and organic chemistry.

Acid–Base Strength and Direction of Proton Transfer

In proton transfer reactions, not all acids and bases behave equally. Some acids donate protons easily (strong acids), others only weakly (weak acids). Corresponding statements apply to bases.

Qualitative Notion of Strong and Weak

Even without numerical measures:

• Strong acids transfer almost all available protons to a given base in a particular medium (e.g., HCl in water)
• Weak acids transfer only a fraction of their protons (e.g., acetic acid in water)
• Strong bases readily accept protons (e.g., $OH^-$ in water)
• Weak bases accept protons only partially (e.g., ammonia in water)

Strength is always relative to:

• The solvent (water, nonaqueous solvents, etc.)
• The other acid or base present

A “strong” acid in water might behave differently in another solvent.

Conjugate Acid–Base Pairs (Conceptual Link)

Any proton transfer connects two pairs of related species:

• An acid becomes a base after giving up $H^+$
• A base becomes an acid after accepting $H^+$

These linked species are called conjugate acid–base pairs and are central to describing acid–base equilibria and reaction direction. The detailed treatment of conjugate pairs and equilibrium is given in the Brønsted and acid–base equilibria sections.

Acids, Bases, and the Solvent

The surrounding medium strongly affects acid–base behavior.

Role of Water as Solvent

In aqueous solutions:

• Water can act as both acid and base (it is amphoteric)
• It participates directly in many proton transfer reactions:
• Accepting a proton to form $\text{H}_3\text{O}^+$
• Donating a proton to form $OH^-$

The balance between these processes leads to characteristic properties of pure water and defines what is considered “neutral” in aqueous systems.

Non-Aqueous and Gas-Phase Acid–Base Behavior

Outside water:

• Different solvents can stabilize or destabilize ions differently
• Substances that are weak acids or bases in water might be strong in other media
• Proton transfer can also occur in the gas phase, though then solvation effects are absent

These aspects highlight that acid–base strength is not an intrinsic property of a substance alone, but of the system (substance + solvent + conditions).

Observing and Measuring Acid–Base Behavior

To work with acids and bases experimentally, one needs ways to detect and quantify acidity or basicity.

Acid–Base Indicators

Indicators are substances that change color depending on the environment’s acidity:

• Organic dyes that exist in differently colored forms depending on protonation state
• Useful for:
• Detecting whether a solution is roughly acidic, neutral, or basic
• Locating the endpoint of a neutralization reaction

Different indicators operate in different acidity ranges, which will be linked later to quantitative measures.

pH as a Measure of Acidity (Conceptual Introduction)

The pH concept compresses the degree of acidity of aqueous solutions into a single number:

• Low pH: acidic
• pH around a characteristic “neutral” point: neither strongly acidic nor strongly basic
• High pH: basic

Exact definitions and calculations, as well as the connection to equilibria and logarithmic measures, are covered in later sections on acid–base equilibria.

Structural Features Behind Acidic and Basic Behavior

The tendency of a substance to donate or accept protons is related to its structure and bonding.

Factors Influencing Acidity

Qualitatively, common influences include:

• Bond strength to the acidic hydrogen (weaker $X–H$ bonds often mean stronger acids)
• Charge distribution and electronegativity (more polar $X–H$ bonds can facilitate proton loss)
• Stabilization of the conjugate base (by resonance, inductive effects, or solvation)

These ideas link acid–base behavior with topics in bonding and structure that are developed elsewhere and revisited in detail in the Brønsted and Lewis theory chapters.

Factors Influencing Basicity

For bases, factors include:

• Availability of an electron pair to bind a proton
• Stabilization or destabilization of the resulting conjugate acid
• Environment (solvent, presence of other ions, etc.)

Later chapters in organic and inorganic chemistry use these concepts to explain reactivity patterns.

Importance of Acids and Bases Across Chemistry

Acid–base processes permeate nearly all areas of chemistry and many technological and biological applications.

In Inorganic and Materials Chemistry

• Formation and decomposition of many salts proceed via acid–base reactions
• Corrosion and protection of metals often involve acidic or basic environments
• Control of solution pH is crucial in precipitation, complex formation, and purification steps

In Organic and Biological Chemistry

• Many organic reaction mechanisms involve proton transfers:
• Activation of functional groups (e.g., carbonyls)
• Formation and stabilization of reactive intermediates
• Biochemical processes (e.g., enzyme catalysis, metabolism) rely on precise acid–base behavior:
• Enzyme active sites contain acidic and basic groups
• Many biomolecules change structure and function with pH

In Everyday Life and Technology

• Household substances: cleaning agents, baking powder, antacids, food acids
• Environmental chemistry: acid rain, buffering of natural waters, soil acidity
• Industrial processes: pickling of metals, production of fertilizers, pH control in chemical manufacturing

Orientation Within the Chapter Group

The present chapter provides the general, phenomenological, and conceptual basis of acids and bases as proton- and electron-pair–related species. Subsequent sections:

• Develop the Brønsted acid–base theory rigorously and introduce conjugate pairs and equilibria
• Examine acid–base equilibria quantitatively, including pH, equilibrium constants, and related calculations
• Discuss amphoterism, neutralization reactions, and Lewis acid–base theory
• Highlight acids and bases in everyday life, connecting theory with practical examples

Together, these sections build a coherent picture of acid–base chemistry as a key type of proton and, more broadly, electron-transfer process.