OCR A-LEVEL CHEMISTRY 5.1.3: ACIDS, BASES & BUFFERS

• How do you Analyse a Carbon-13 NMR Spectrum of an Organic Molecule?

  • No. of Peaks = No. of carbon environments.• Chemical Shift = Type of environment (use data sheet).• C atoms bonded to different atoms/groups differ in environments & absorb at different chemical shifts. • 2 C atoms positioned symmetrically are equivalent. • This means the same: chemical environment, absorbtion at chemical shift & contribution to the peak.
• How do you Analyse a Proton NMR Spectrum of an Organic Molecule?

  • No. of Peaks = No. of Proton Environments.• Chemical Shift = Type of environment (use data sheet).• Relative Peak Areas = Relative no. of each type of H⁺.• Splitting Pattern = No. of non-equivalent H⁺ on adjacent carbon atom + 1 (n + 1 rule), to give: singlets, doublets, triplets, quartets or multiplets (4+).
• What are Uses of TMS, Deuterated Solvents & D₂O in NMR Spectroscopy?

  • Tetramethylsilane, TMS, is the standard for chemical shift measurements & is given a chemical shift of 0 ppm.• All results of NMRs are calibrated by & relative to TMS.• Deuterated solvents (e.g. CDCl₃) contain ²H (Deuterium) rather than ¹H. • Deuterium gives no signal in frequency ranges used for ¹H and ¹³C NMR.• Adding D₂O exchanges & replaces OH or NH protons with ²H atoms: eliminating the broad OH & NH peaks.
• What are Brønsted–Lowry Acids & Alkalis?

  • Brønsted-Lowry acids are proton donors.• Brønsted-Lowry bases are proton acceptors.
• What are Conjugate Acid-Base Pairs?

  • A conjugate acid-base pair contains two species that can be interconverted by transfer of a proton.• HCI (aq) & Cl⁻ (aq) are a conjugate acid-base pair.• Forward direction: HCI releases a H⁺ to form its conjugate base, CI⁻.• Reverse direction: CI⁻ accepts a H⁺ to forms its conjugate acid, HCI.

  
• What are Mono/di/tribasic Acids?


  • Monobasic acids dissociate One H⁺ in solution.• Dibasic acids dissociate Two H⁺'s in solution.• Tribasic acids dissociate Three H⁺'s in solution.

  
• What is the Role of H⁺ in Reactions of Acids with Metals & Bases?

  • Acids react with metals, carbonates, metal oxides & alkalis to form salts. • The active species from the acid is H⁺ (aq). • Ionic equations can be written for each type of reaction to emphasise the role of H⁺ ions. • When writing ionic equations with H⁺ (aq), the acid does not matter: it's the same reaction.
• What is the Relationship between pH & Proton Concentration?

  • pH = -log[H⁺].• [H⁺] = 10⁻ᵖᴴ.• pH of 1 has 10 times more [H+] than solution of pH 2. • To dilute a solution from a pH 1 to 4 (3 pH units), you would need dilution by 10 x 10 x 10 = 1000 times. • A solution with a pH of 1 contains 10¹³ times more H⁺ (aq) ions as a solution with a pH of 14. • Comparing 1 with 14 is more manageable than comparing 1 x 10⁻¹ with 1 x 10⁻¹⁴.

  
• What is Relationship between Ka & pKa?

  • Acid dissociation constant, Ka, shows extent of acid dissociation & it changes with temperature.• Recorded Ka values are usually standardised at 25 °C. • Increased Ka means equilibrium has shifted right. • Larger Ka value = Greater dissociation & acid strength.• Ka = 10⁻ᵖᵏᵃ.• pKa = -logKa.

  
• What is Kw?


  • Ionic product of water: Kw = [H⁺(aq)] [OH⁻(aq)].• Kw varies with temperature. • At 298K (25°C) Kw is 1.00 x 10⁻¹⁴mol²dm⁻⁶.• Kw sets up the neutral point in the pH scale.• A solution is acidic when [H⁺(aq)] > [OH⁻(aq)].• A solution is neutral when [H⁺(aq)] = [OH⁻(aq)].• A solution is alkaline when [H⁺(aq)] < [OH⁻(aq)]. • Acidic solutions still have OH⁻ ions, but more H⁺. • For pH values that are whole numbers, the indicies for [H⁺(aq)] & [OH⁻(aq)] add to -14.

  
• How can you Calculate pH of Strong Monobasic Acids?

  
• How can you Calculate pH of Weak Monobasic Acids?


  Approximations & Their Limitations:1) Dissociation of water is negligible: • Breaks down with very weak acids or dilute solutions. 2) [HA(aq)] >> [H⁺(aq)] concentration at equilibrium:• Breaks down for stronger weak acids or very dilute solutions & only holds true for very weak acids.

  
• How can you Calculate pH of Strong Bases Using Kw?

  
• What are Buffer Solutions?

  • Systems that minimise pH changes on addition of small amounts of an acid or a base.• They contain weak acids & their conjugate bases.• The weak acid, HA, removes added alkali. • Small concentration of H⁺ reacts with the OH-(aq) ions: H⁺(aq) + OH⁻(aq) → H₂O(l). • HA dissociates, shifting equilibrium position to the right to restore most of H⁺(aq) ions.• The conjugate base, A⁻, removes added acid.• Conjugate base reacts with A⁻, shifting equilibrium to the left; to remove most of the H⁺ ions.• The pH still changes with buffers, but changes less.• If either HA or A⁻ run out, the buffer stops working.• pH of buffer solution is same as the pka value of HA.

  
• How can Buffer Solutions be Prepared?


  From a Weak Acid & its Salt:• If CH₃COOH is added to water, the acid partially dissociates & amount of CH₃COO⁻ is very small. • Ethanoic acid is the source of the weak acid component of the buffer solution. • CH₃COOH (aq) ⇌ H⁺ (aq) + CH₃COO⁻ (aq).• Salts of weak acids are ionic compounds and provide a convenient source of the conjugate base. • When added to water, the salt completely dissolves. • Dissociation into ions is complete, and so the salt is the source of the conjugate base of the buffer solution. • CH₃COONa (s) + aq → CH₃COO⁻ (aq) + Na (aq).From Excess Weak Acid & Strong Alkali:• Weak acid (CH₃COOH) is partially neutralised by a base (NaOH) to form the conjugate base.• Some of the weak acid is left over unreacted.• The resulting solution is a mixture of the salt of the weak acid & unreacted weak acid.
• How can pH be of a Buffer Solution be Calculated?

  
• How is Blood pH Controlled?

  On Addition of Acid, H⁺(aq):• [H⁺(aq)] increases.• H⁺ ions react with the conjugate base, HCO₃⁻(aq). • Equilibrium shifts left, removing most H⁺(aq) ions. • H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq).On Addition of Alkali, OH⁻(aq):• [OH⁻(aq)] increases.• H⁺(aq) + OH⁻(aq) → H₂O(l). • Equilibrium shifts right, restoring most H⁺(aq) ions. • H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq).
• What is the pH Titration Curve for Addition of Base to Acid?

  • A pH meter has an electrode, dipped into a solution, connected to a meter that displays the pH reading. • It records pH values to two decimal places. • Indicator paper is usually matched from colour charts to the nearest whole number, pH meters give more accurate measurements of pH during a titration.• At first, acid is in excess & pH increases very slightly. • As vertical section approaches, pH starts to increase more quickly as the acid is used up more quickly. • Eventually, pH increases rapidly during addition of a small volume of base, producing the vertical section. • Only drops of solution are needed for vertical section. • After the vertical section, the pH will rise very slightly as the base is now in great excess.• Equivalence point of titration is when volume of one solution exactly reacts with volume of another solution. • Solutions have then exactly reacted with one another & amounts used match stoichiometry of the reaction. • Equivalence point is centre of the vertical section. • The curve may be a different shape for combinations of add and base with different strengths.

  
• What is the pH Titration Curve for Addition of Acid to Base?

  
• What are Acid-Base Indicators?

  • An acid-base indicator is a weak acid.• Has very different colour from conjugate base, A⁻. • For methyl orange HA is red & A⁻ is yellow. • At end point of titration, indicator has equal concentrations of HA & A⁻; colour will be in between.• Most indicators range over two pH units.

  
• What are the Different pH Titration Curves?

  • No indicator is suitable for a weak acid-weak base titration as there is no vertical section and, even at its steepest, the pH requires several cm³ to pass through a typical pH indicator range of 2 pH units.