Summary
Highlights
Acids react with bases to form salts and water, a process called neutralization. The H+ ions from acids react with OH- ions from alkalis to form water. Ionic equations focus on the reacting species, excluding spectator ions. Ammonia reactions with acids produce ammonium salts but no water.
Acids react with metals to produce salt and hydrogen gas. With metal oxides (treated as bases), they form salt and water. Metal hydroxides also produce salt and water. Metal carbonates react with acids to form salt, water, and carbon dioxide, which can be observed as fizzing. Ionic equations for these reactions highlight the essential changes.
A conjugate pair is linked by the transfer of a proton. A species that gains a proton becomes a conjugate acid, and one that loses a proton becomes a conjugate base. Water can act as both an acid and a base (amphoteric), accepting a proton to form H3O+ or donating one to form OH- depending on the reaction.
Weak acids and bases react reversibly with water, resulting in incomplete dissociation. Strong acids and bases dissociate almost completely. Examples include ethanoic acid (weak acid), HCl (strong acid), NaOH (strong base), and ammonia (weak base). The extent of dissociation determines the concentration of H+ or OH- ions, influencing pH.
Acid-base reactions involve proton exchange and are always in equilibrium. Le Châtelier's principle applies: adding more reactants shifts the equilibrium to the right, producing more products, while increasing product concentration shifts it to the left. Water acts as a base when an acid is added to it, accepting a proton to form H3O+.
Water exists in equilibrium with its ions (H+ and OH-). The equilibrium law applies, leading to the ionic product of water, Kw = [H+][OH-]. The concentration of water is considered constant due to its weak dissociation. At 298 K, Kw is 1 x 10^-14 mol^2 dm^-6, a value that changes with temperature. In pure water, [H+] = [OH-], so Kw = [H+]^2.
pH is a logarithmic scale measuring H+ concentration (pH = -log[H+]). It ranges from 0 (very acidic) to 14 (very basic), with 7 being neutral. For strong monoprotic acids, the acid's concentration directly equals the H+ concentration because they dissociate fully. For diprotic acids like sulfuric acid, the H+ concentration is twice the acid's concentration.
To calculate the pH of a strong base, we first assume full dissociation to find the OH- concentration. Then, use Kw = [H+][OH-] to find the H+ concentration, and finally, pH = -log[H+]. For example, a 0.15 mol dm^-3 NaOH solution at 298 K yields a pH of approximately 13.18, which is consistent with a strong base.
Weak acids dissociate slightly, so their H+ concentration cannot be assumed equal to the initial acid concentration. The acid dissociation constant (Ka) is used instead: Ka = [H+][A-]/[HA]. For weak acids, we make two key assumptions: the concentration of undissociated acid remains approximately constant at equilibrium, and all H+ ions come from the acid. This simplifies Ka to Ka = [H+]^2/[HA].
To calculate the pH of a weak acid, use the simplified Ka expression. Rearrange to find [H+] (which will be the square root of Ka * [HA]), and then use pH = -log[H+]. It's also possible to work backward to find the acid concentration or Ka if the pH is known. Remembering the Pka = -log(Ka) relationship makes calculations easier.
pH meters accurately measure pH but require calibration using buffer solutions of known pH (e.g., pH 4, 7, and 10). The probe must be cleaned with distilled water between measurements to avoid contamination. Careful calibration ensures accurate readings.
Titrations are vital for determining the unknown concentration of an acid or base. A solution of known concentration (in the burette) is added to a solution of unknown concentration (in the conical flask) until an indicator changes color at the endpoint. Accurate technique, including drop-by-drop addition near the endpoint, is crucial. Results should be concordant (within 0.1 cm^3).
Titration curves plot pH against the volume of titrant added, showing distinct 'S'-shapes for different acid-base combinations. The equivalence point (or endpoint) is the sharp vertical rise where neutralization occurs. Indicators must change color entirely within this vertical region. Methyl orange is suitable for strong acid-strong base or strong acid-weak base titrations, while phenolphthalein is for weak acid-strong base titrations. Weak acid-weak base titrations lack a sharp pH change, making a pH meter necessary.
Titration calculations involve balancing the reaction equation, calculating moles of the known substance (concentration x volume), using stoichiometry to find moles of the unknown substance, and then calculating the unknown's volume or concentration. Remember to convert volumes to dm^3 (divide by 1000) and ensure units are consistent.
A buffer is a chemical system that resists changes in pH when small amounts of acid or base are added. There are acidic buffers (maintaining pH<7) and basic buffers (maintaining pH>7). An acidic buffer typically consists of a weak acid and its salt (e.g., ethanoic acid and sodium ethanoate). Two equilibrium equations coexist in a buffer solution.
In an acidic buffer (weak acid + its salt): The weak acid is in equilibrium, producing H+ and its conjugate base (A-). The salt fully dissociates, providing a high concentration of A-. When H+ is added, it reacts with A- (from the salt) to form more of the weak acid. When OH- is added, it reacts with H+, and Le Châtelier's principle causes the weak acid to dissociate further, replenishing H+. This system effectively maintains a stable pH.
To calculate a buffer's pH, the KA expression is used. Unlike weak acid calculations, [H+] does not equal [A-] because A- also comes from the fully dissociated salt. Therefore, [A-] is approximated by the initial salt concentration, and [HA] by the initial weak acid concentration. The formula [H+] = Ka * ([HA]/[A-]) is then used, followed by pH = -log[H+]. This is one of the more complex calculations in A-level chemistry.
Calculating the pH change after adding a strong acid or base to a buffer is even more complex. It involves calculating the initial moles of the weak acid and its salt, determining how the added substance reacts with them, calculating the new moles of the weak acid and salt, and then adjusting for the new total volume. Finally, recalculate [H+] and pH using the buffer pH calculation method. It is crucial to account for the change in moles of both the acid and its conjugate base.
Buffers are crucial in many household products and biological systems. Blood is a classic example of a buffer system, maintaining a pH of approximately 7.4. This stability is essential for enzyme function and cell integrity. The carbonic acid (H2CO3) and hydrogen carbonate (HCO3-) system in blood plasma, regulated by respiration (CO2 levels) and kidneys (HCO3- levels), is vital for life.
Acids, bases, and buffers can be a challenging topic. The video encourages viewers to subscribe for more free revision content and offers the slides for purchase. The importance of practice and understanding key concepts is reiterated.
This video is a revision tool for OCR A Chemistry, specifically covering acids, bases, and buffers. It's designed to refresh knowledge and emphasize the importance of practicing exam questions. The content is tailored to the OCR A specification, providing all necessary information and nothing more.
Scientific theories evolve over time. Antoine Lavoisier in 1778, known for his work on oxygen, linked oxygen to acids, noting formulas like H2SO4 and HNO3. However, his theory was limited as he wasn't aware of acids like HCl, which don't contain oxygen.
Svante Arrhenius in 1884 proposed that acids donate H+ ions and bases release OH- ions. He also suggested that acid-base reactions form salt and water. His theory, however, couldn't explain why substances like ammonia (NH3) acted as bases without directly releasing OH- ions.
In 1923, Brønsted and Lowry improved Arrhenius's theory, defining acids as proton donors and bases as proton acceptors. They explained that H+ ions don't exist alone in water but form hydronium ions (H3O+). This theory successfully explained ammonia's basicity by stating it accepts a proton from water to form NH4+ and OH-.
Some acids can donate more than one proton. Monoprotic acids (e.g., HNO3) donate one proton per molecule. Diprotic acids (e.g., H2SO4) donate two protons, and triprotic acids (e.g., H3PO4) donate three. Understanding these variations is crucial for acid-base chemistry.