Acid-Base Chemistry Quiz

Biochemistry Laboratory Session Reminder

• Biochemistry lab session is on Thursday (tomorrow).
• Complete the Health and Safety videos and short quiz.
• Review the Learning Science virtual simulations.
• Bring a printed copy of the lab instructions (proforma).
• Arrive on time for the health and safety briefing. Late arrivals will not be permitted to participate in the lab session.
• Check your personal timetable for the specific time and location of the lab session.

Practical One: Bradford Assay

• Find the necessary information in the "PRACTICAL ONE Bradford Assay" folder on the Biochemistry Blackboard site.
• This folder contains lecture slides, tutorial materials, practical instructions and other relevant information.
• The practical is formative, but the skills will be assessed in a summative practical test in January 2025.

Fundamentals of Chemistry: Acids and Bases and pH

• Learning Outcomes:

  • Define pH.
  • Understand how pH relates to the mechanisms of action of acids and bases in water.
  • Understand the relationship between Ka and pKa and pH.
  • Explain how buffers function.

• Introduction:

  • pH is a measure of acidity in a solution (short for "power of hydrogen").
  • It measures the activity of dissolved hydrogen ions.
  • Hydrogen ions exist as protons (H+) and hydronium ions (H₃O⁺).
  • Water dissociates as follows: H₂O + H₂O ⇌ H₃O⁺ + OH⁻ or H₂O ⇌ H⁺ + OH⁻.

• Introduction (continued):

  • In pure water at 25°C, the concentration of hydrogen ions (H⁺) equals the concentration of hydroxide ions (OH⁻).
  • Neutral pH is 7.0.
  • Acidic solutions have a higher concentration of H⁺ ions than OH⁻ ions and a pH less than 7.0.
  • Basic solutions have a higher concentration of OH⁻ ions than H⁺ ions and a pH greater than 7.0.

pH of Everyday Solutions

• Various everyday solutions and their corresponding pH values are presented in a table format (refer to the provided image).

Definitions

• pH is a measurement of the concentration of hydrogen ions in a solution.
• Low pH values indicate a high concentration of hydrogen ions.
• High pH values indicate a low concentration of hydrogen ions.
• Pure water has a pH of 7.0.

Definitions Continued

• Acids are solutions with a pH below 7 (more hydrogen ions than water).
• Bases are solutions with a pH greater than 7 (fewer hydrogen ions than water).
• The definition of strong or weak acids/bases doesn't refer to pH, but to how readily they ionize in solution.

Alternative Definition of an Acid or Base (Brønsted-Lowry Theory)

• Acid: a proton donor.
• Base: a proton acceptor.
• This definition is independent of water.

Any Questions?

• Questions can be asked via synchronous chat during the live session or via the asynchronous Question Board on the Biochemistry Blackboard module.

Definition of pH

• pH = -log[H₃O⁺] orpH = -log[H⁺]
• pH scale: indicates acidity/alkalinity of a solution.
• Measured by the number of H⁺ ions (or equivalents) in the solution.
• Ranges from 0 (very acidic) to 14 (very basic).
• pH is a measure of the effective concentration of hydrogen ions, not the actual concentration.
• In practice, hydrogen ions can be shielded or hidden, so they are not always available to participate in chemical reactions (e.g., inside a protein molecule).

pH in Biological Systems

• Table of bodily fluids and their corresponding pH values.

The Case of Water

• Pure water is mostly molecular (covalently bonded).
• A very small amount is ionized (dissociated).
• In pure water at room temperature [H₃O⁺] = 1 x 10⁻⁷ M and [OH⁻] = 1 x 10⁻⁷ M

pH Calculations

• pH calculation using the formulas pH = -log[H⁺] or pH = -log[H₃O⁺].
• Example calculations are shown, including how to find the pH value given a concentration.
  • Calculate pH when [H₃O⁺] = 1 x 10⁻⁷ M

pH of Pure Water

• Normally, pure water has a pH of 7.
• When exposed to air and absorbing CO₂, the pH drops to around 5.7.

pH Scale

• The difference between each pH value is 10-fold
• Show the relationship between pH and the concentration of H+.

Strong and Weak Acids

• Tables of examples of strong and weak acids.
• Definitions and explanations of strong and weak acids

Strong and Weak Acids (continued)

• Strong acids completely dissociate in water, while weak acids only partially dissociate.
• HCl example reaction with water: HCl + H₂O → Cl⁻ + H₃O⁺.
• Ethanoic acid (CH₃COOH) example reaction with water: CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺.

Calculating pH for Weak and Strong Acids

• Strong acids: Dissociation reaction goes to completion; no unreacted acid remains in solution.
• Reaction is simplified as HX → H⁺ + X⁻.
• To calculate pH for strong acids, only need to know the concentration of the acid (HX) present.
• Weak acids: Dissociation reaction does not go to completion; an equilibrium is reached between the hydrogen ions and the conjugate base.
• Reaction is: HX ⇌ H⁺ + X⁻ .
• The acidity constant (Ka), which is equal to the equilibrium constant, must be used.

Calculation of pH for Weak and Strong Acids (Continued)

• Ka is used in calculating pH values for weak acids
• Comparing Ka values will help to determine which acids are stronger based on their ionization/dissociation in solution.
• Differences in Ka values reveal which acids ionize/dissociate more greatly and which are weaker.

Strong vs Weak Acids (Examples)

• Example reactions of strong and weak acids given and their different characteristics explained.

Calculating pH of Weak Acids - (Continued)

• When calculating pH of weak acids, the equilibrium constant denoted as Ka is used.
• Equilibrium constant (Ka) is a crucial factor to determine the behaviour of weak acids and their dissociation.

Bases

• Strong bases are completely ionized; examples include NaOH, KOH, and Ca(OH)₂.
• Weak bases are partially ionized; ammonia (NH₃) is an example.
• Alkalis are bases that are soluble in water.

Bases and pH

• Strong bases ionize to form hydroxide ions (OH⁻).
• Pure water dissociates to form hydronium (H₃O⁺) and hydroxide (OH⁻) ions.
• H₃O⁺ and OH⁻ ions combine to create water molecules.
• Hydroxide ions remove hydronium and hydrogen ions, thus reducing concentration of H₃O⁺ and H⁺ and subsequently causing the pH to rise.

Relationships between pOH, pH, and Kw

• Equilibrium constant of water ionization (Kw) = [H₃O⁺][OH⁻]. (equal to 1.0 ×10⁻¹⁴)
• pKw = pH + pOH = 14, where pOH= -log[OH⁻].
• Calculate pH of basic solutions; pH = 14 - pOH.
• Example: Calculate pH for 0.5 M NaOH in water, [OH⁻] = 0.5 M, pOH = - log (0.5) = +0.3, pH = 14 - 0.3 = 13.7

pH is Concentration of H⁺ Ions

• Values of pH and corresponding concentration of H⁺ are displayed.

Neutralization

• Neutralisation reaction between an acid and a base produces a salt and water. Example: HCl (aq) + NaOH (aq) → NaCl (aq) + H₂O (l).
• Neutralization reaction is not always equal to a pH of 7.
• It only produces a pH of 7 if the acid and base have similar strength.

Buffer Solution

• Buffer solution resists significant pH changes upon addition of acidic or basic components.
• A buffer system consists of a weak acid and its salt, or weak base and its salt.
• Buffers are essential in biochemical and biological reactions.

Buffers

• Buffer solution: an aqueous solution mixing weak and conjugate acid/base
• Buffer helps to maintain the pH in a biochemical reaction.
• Example: Blood Plasma, maintain pH between 7.35 and 7.45 via bicarbonate-carbonic acid.

What Happens to pH if...?

• This slide shows how the pH of an unbuffered solution changes with the addition of either H⁺ or OH⁻.

Buffers (Continued)

• Buffer solutions: consist of mixtures either one weak acid plus its salt, or weak base plus its salt.
• Example reaction: CH₃COOH (aq) + H₂O (l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq).

Additives in Buffers

• Example of a buffer composition, and effect of addition of different substances (H⁺, OH⁻) is explained.

Buffering - The Henderson-Hasselbalch Equation

• Henderson-Hasselbalch equation is used to calculate pH of buffer solutions.
• pH = pKa + log ([base]/[acid]), where pKa is -log(Ka).

Buffering Capacity

• Graph of pH change versus volumes of added base, showcasing the impact of the buffer to maintain relatively constant pH.

Buffering Range

• The pH range in which a buffer solution function optimally will be approximately ±1 pH unit from its pKa value
• The highest buffering capacity is when the pH value equals the pKa value.
• Amino acids and Proteins have pKa values

Buffers in Biology

• Buffers are crucial in biological systems for maintaining stable pH.
• An ideal buffer has a pKa matching the desired pH.
• Buffer solutions are necessary to preserve the precise pH range needed by enzymes in various biochemical reactions.
• Example in blood: carbonic acid (H₂CO₃) and bicarbonate (HCO₃⁻).

MCQ Quiz for Lecture 5

• Multiple Choice Questions (MCQs) for Lecture 5, Fundamentals of Chemistry Part 4.
• Answers to quizzes provided during seminar sessions, which need to be attended before the seminar session.
• These quizzes are part of the Biochemistry module and help with on-going studies.

Quiz Questions (Q1-Q5):

• Five multiple choice questions regarding the topics of pH calculations, weak bases, and acidic constants.
• Correct answers are directly related to the formula and explanation provided in the previous notes.

Test your knowledge on acid-base chemistry with this quiz, which covers concepts such as pH levels, strong and weak acids, and the effects of carbon dioxide on pH. Explore key definitions and relationships that are crucial for understanding chemical interactions. Perfect for students studying general chemistry or environmental science!