An acid is a substance that increases the concentration of hydrogen ions (protons) in a solution or, more broadly, a substance that can donate a proton to another substance. What makes an acid "acidic" depends on the definition you use, but the core idea is that acids lower the pH of a solution and are able to furnish H+ or overall increase acidity in the system. Key ways chemists describe acidity

• Arrhenius definition (classic, aqueous solutions)
  • Acids are substances that increase the concentration of hydrogen ions (H+) in water.
  • Examples: HCl, H2SO4, and CH3COOH release H+ in solution, producing a solution with pH below 7. This is the basis for the common classroom description of acids as sour-tasting, corrosive, and capable of turning blue litmus red. In this framework, acidity correlates with hydronium ion (H3O+) concentration in the solution. [Acid definition and properties apply broadly to many acids in water.]
• Bronsted–Lowry definition (proton donor)
  • Acids are substances that donate a proton to a base.
  • The conjugate base that remains after proton donation is formed, and the strength of an acid is related to how readily it donates that proton.
  • Strong acids donate protons more completely in solution; weaker acids donate less readily, leaving more undissociated acid species. This concept helps explain acidity across a wide range of solvents, not just water. [Bronsted–Lowry concept and acid strength]
• Lewis definition (electron-pair acceptor)
  • Acids are electron-pair acceptors.
  • This broad view encompasses many species that accept electron pairs, including many metal cations and some nonmetals, and is useful for contexts beyond proton transfer. [Lewis acid concept]

What determines acidity in different contexts

• In water (common aqueous chemistry):
  • Acid strength is quantified by the acid dissociation constant, Ka, or its negative logarithm, pKa. A smaller pKa means a stronger acid (more willing to donate H+). [Acid strength in water]
• In organic chemistry:
  • Acidity is influenced by factors such as:
    • Stabilization of the conjugate base (through resonance, inductive effects, and hyperconjugation)
    • The atom bearing the acidic proton (size and electronegativity) and its ability to delocalize negative charge
    • Availability of charge distribution and resonance across the molecule after deprotonation
  • These factors explain why carboxylic acids are typically stronger acids in organic solvents than alcohols, and why electron-withdrawing groups nearby an acidic proton increase acidity. [Factors influencing organic acidity]
• In non-aqueous solvents:
  • Acid strength can differ from water-based expectations because solvent stabilization of ions changes. Some species classified as weak acids in water may behave differently in other solvents, and vice versa. [Solvent effects on acidity]

Common indicators of acidity

• pH below 7 in aqueous solutions (high H3O+ concentration)
• Sour taste (historical and practical observation for many acids, though tastings are not recommended for safety)
• Ability to turn blue litmus paper red
• Reactivity with metals to produce hydrogen gas and with carbonates to form carbon dioxide, depending on the acid. [Acidic properties and indicators]

If you’d like, specify the context (general chemistry, organic chemistry, aqueous vs. non-aqueous solvents), and the level of detail (high-level overview or rigorous quantitative treatment), and the explanation can be tailored with examples and key equations.