Acids, bases, pH and salts
In introductory aqueous chemistry, an acid produces hydrogen ions in water and an alkali produces hydroxide ions in water. A base neutralises an acid; an alkali is a base that dissolves in water. All alkalis are bases, but not every base is soluble. These definitions describe the school-level aqueous context and should not be applied carelessly to every possible solvent or reaction.
The pH scale expresses hydrogen- concentration logarithmically. In the usual dilute-solution treatment. At 25°C, a neutral aqueous solution has pH 7, acidic solutions have lower pH and alkaline solutions have higher pH. Temperature matters for the exact neutral value. A change of one pH unit represents a tenfold concentration change, rather than a one-unit change in ordinary concentration.
Strength and concentration are different ideas. A strong acid ionises essentially completely in dilute aqueous solution, while a weak acid only partly ionises. Concentration states the amount of acid per volume of solution. A dilute strong acid can contain fewer hydrogen ions per unit volume than a sufficiently concentrated weak acid, so neither the word ‘strong’ nor ‘weak’ alone tells you the whole numerical pH.
Neutralisation often has the form acid + base → salt + water. Hydrochloric acid and sodium hydroxide react as. Acids also react with carbonates to form a salt, water and carbon dioxide. The characteristic effervescence is evidence of gas production, but identifying the gas requires an appropriate test rather than assuming every bubbling reaction releases carbon dioxide.
An acid–base uses measured solution volumes and a known concentration to find an unknown amount. Read the meniscus at eye level, choose a suitable indicator and approach the endpoint slowly. The indicator endpoint should be close to the reaction’s equivalence point. Calculate moles from concentration and volume, apply the balanced ratio, and then divide by the required solution volume.
Laboratory safety supports accurate chemistry. Do not taste chemicals or use household experiments as substitutes for supervised laboratory procedures. When instructed to dilute concentrated acid, the usual safe procedure adds acid slowly to water with suitable protection and cooling, never water into concentrated acid. In written work, distinguish observations such as colour change from inferences such as complete neutralisation, and state assumptions in numerical pH problems.
Acid–base ideas beyond a single aqueous formula
A Brønsted acid donates a proton and a Brønsted base accepts a proton. In HCl + H₂O → H₃O⁺ + Cl⁻, water acts as a base by accepting a proton. In NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, ammonia is a base even though its molecular formula contains no hydroxide group. Conjugate acid–base partners differ by one proton. These ideas explain more reactions than a rule that all bases must contain OH. In introductory calculations, H⁺ is convenient shorthand for the hydrated proton in water; distinguish the shorthand from a claim that isolated bare protons simply float unassociated in solution.
Indicators and what their colours show
An acid–base indicator changes its dominant form over a particular pH range. Litmus, methyl orange and phenolphthalein therefore do not all change at exactly pH 7. Litmus can show a solution is acidic or alkaline but does not determine a precise concentration. Universal indicator gives an approximate range by colour comparison, while a calibrated pH meter can provide a numerical reading. For , choose an indicator whose transition range lies within the steep pH change near equivalence. A strong acid–strong base has several suitable choices, while a weak acid–strong base generally favours an alkaline-range indicator such as phenolphthalein under ordinary teaching conditions.
Strength, basicity and concentration
Acid strength describes extent of ionisation, whereas basicity of an acid counts the replaceable ionisable protons per in the relevant acid–base reactions. Hydrochloric acid is monoprotic and sulfuric acid is commonly treated as diprotic in elementary neutralisation stoichiometry. Ethanoic acid has four hydrogen atoms but only one ordinary acidic proton, so counting every H symbol gives the wrong basicity. Concentration states amount per solution volume, independent of whether ionisation is complete. Conductivity depends on concentration and mobility, so comparing acid strengths through conductance needs comparable concentrations, temperature and appropriate interpretation rather than a casual brightness observation alone.
pH, pOH and logarithmic comparisons
In the ordinary dilute-solution treatment at and [H⁺][OH⁻] ≈ 1.0 × 10⁻¹⁴. A solution with pH 3 has about ten times the hydrogen- concentration of pH 4 and about one hundred times that of pH 5. For a completely ionised dilute monoprotic acid, hydrogen- concentration approximately equals acid concentration. Do not use that shortcut for a weak acid without information. A dilute strong base with has pOH 2 and pH 12 at 25°C. The temperature qualification matters because water's ionisation changes with temperature.
Choosing a salt-preparation method
To prepare a soluble salt from an acid and insoluble base or carbonate, add the solid in excess, filter off the unreacted solid and concentrate the filtrate before controlled crystallisation. To prepare a soluble salt from an acid and an alkali, use to find the required volumes, then repeat with measured reactants without indicator if pure crystals are needed. Insoluble salts can be prepared by precipitation from two suitable soluble salt solutions, followed by filtration, washing and drying. A method must match the salt's solubility and reactants. Simply evaporating every mixture to dryness can trap impurities or decompose heat-sensitive substances.
Salt types and hydrolysis
An acid salt, such as sodium hydrogensulfate, retains replaceable hydrogen from incomplete neutralisation of a polyprotic acid. This name is different from saying that a salt solution is acidic. Ammonium chloride is a normal salt whose aqueous solution is acidic because NH₄⁺ can transfer a proton to water. Sodium ethanoate solution is alkaline because ethanoate can accept a proton from water. A strong acid–strong base salt such as sodium chloride gives an approximately neutral solution under ordinary dilute conditions. Double salts are crystalline combinations that dissociate into their constituent ions in solution; their classification does not that two unrelated solutions are automatically a new pure double salt.
reasoning and buffer action
At equivalence the reactants have combined in their balanced stoichiometric ratio, not necessarily at equal volumes or pH 7. An endpoint is the observed signal used to estimate that point. Read the burette consistently, remove air bubbles from its tip and add titrant dropwise near the endpoint under supervision. A buffer contains suitable conjugate partners that resist a large pH change on addition of small amounts of acid or base. It does not keep pH perfectly fixed under unlimited addition. A weak acid and its soluble conjugate-base salt provide one common buffer type. Buffer capacity depends on amounts and composition, not the word buffer alone.
Self-check and writing a sound explanation
An excess of copper(II) oxide can react with warm dilute sulfuric acid to prepare copper(II) sulfate solution. Filter away remaining oxide, then crystallise the salt from the clear solution; do not describe filtering as removing dissolved sulfate ions. In a separate 2:1 NaOH:H₂SO₄ , 0.010 mol sulfuric acid requires 0.020 mol sodium hydroxide. A pH 2 solution has one thousand times the approximate hydrogen- concentration of pH 5, since the difference is three logarithmic units. When asked to distinguish a strong acid from a concentrated acid, give the two definitions and a consistent example rather than repeating that both are dangerous.