Lecture 5 Chemistry p4 Acids Bases MCQ PDF

REMINDER: Thursday (tomorrow) you will be doing your FIRST laboratory session for Biochemistry module. You MUST go through and complete the Health and Safety videos and short quiz. You should go through the short Learning Science virtual simulations. They will help you get the most from th...

REMINDER: Thursday (tomorrow) you will be doing your FIRST laboratory session for Biochemistry module. You MUST go through and complete the Health and Safety videos and short quiz. You should go through the short Learning Science virtual simulations. They will help you get the most from the lab session. You MUST bring a paper copy of the proforma (instructions) with you. You MUST arrive on time for the health and safety briefing. Late arrivals will not be allowed to do the laboratory session. Look at your PERSONAL timetable for time and room  You will find the information you need in PRACTICAL ONE Bradford Assay folder on Biochemistry Blackboard site. This is a formative practical BUT the skills you learn here will be assessed in the in summative in class Practical test in January 2025 Fundamentals of Chemistry: Acids and Bases and pH 4BICH001W Biochemistry Dr Sarah K Coleman  Learning outcomes Define what pH is Understand how pH relates to mechanism of action of acids and bases in water Understand how Ka and pKa are related to pH Explain what buffers do  Introduction pH = the measure of the acidity of a solution (pH stands for “power of hydrogen”) a measure of the activity of dissolved hydrogen ions Hydrogen ions occur as various of cations including protons (H+) and hydronium ions (H3O+) in solution Dissociation of water: H2O + H2O  H3O+ + OH- or H2O  H+ + OH-  Introduction cont… In pure water at 25 °C the concentration of hydrogen ions (H+) equals the concentration of hydroxide ions (OH-) "neutral" corresponds to a pH level of 7.0 ACID Solutions are where concentration of [H+] exceeds that of [OH- ]. They have a pH value lower than 7.0 BASIC Solutions are where [OH- ]exceeds [H+]. They have a pH value greater than 7.0  pH of Everyday Solutions  Definitions pH = a measurement of concentration of hydrogen ions in a solution low pH values  solutions with high concentration of hydrogen ions high pH values  solutions with low concentration of hydrogen ions Pure water has a pH of 7.0; other solutions are often described with reference to this value  Definitions continued Acids  solutions that have a pH less than 7 (i.e. more hydrogen ions than water) Bases  a pH greater than 7 (i.e. less hydrogen ions than water) The definition of weak and strong acids OR weak and strong bases does not refer to pH value It describes how well an acid or base ionizes in solution Alternative definition of an acid or base Brønsted-Lowry Theory of Acids and Bases Acid is a proton donor Base is a proton acceptor NOTE: this definition is INDEPENDENT of water. What is a proton? Any questions: You can type in the chat function box during this live session (synchronous)? Or onto the Question Board in the Biochemistry Blackboard module and I will look at them later (asynchronous). The pH of a solution is defined as: pH = -log[H3O+] pH scale - a scale that indicates the acidity or basic nature of a solution Measure of number of H+ ions (or equivalent) in solution Ranges from 0 (very acidic) to 14 (very basic) The pH of a solution is defined as: pH = -log[H+] A measure of the effective concentration of hydrogen ions (rather than the actual concentration) In practice hydrogen ions could be shielded / hidden so not available to participate in chemical reactions e.g. inside a protein molecule Having solutions at the correct pH is critical to living organisms  pH in biological systems Bodily Fluid pH gastric acid 0.7 lysosome 5.5 granule of chromaffin cell 5.5 Pure, neutral H2O at 37 ºC 6.81 Cytosol 7.2 Cerebral Spinal Fluid 7.3 arterial blood plasma 7.4 mitochondrial matrix 7.5 exocrine secretions of pancreas 8.1  The Case of Water… Pure water almost 100% molecular (covalently bonded). Very small amount is ionised (dissociated) H2O + H2O  H3O+ + OH- In pure water at room temperature: [H3O+] = 1 x 10-7 M [OH-] = 1 x 10-7 M (Square brackets [ ] mean the concentration of a substance)  pH = -log[H+] or pH = -log[H3O+] [H3O ] = 1 x 10 M + -7 (1 x 10-7 = 0.0000001) So, pH = -log[1x10-7]  So the pH of pure water is: pH = - log[H3O+] However, when pure water is exposed to the atmosphere CO2 will be absorbed and pH = - (- 7) react with water to form carbonic acid (HCO3- and H+). pH = 7 So pH lowered to approx 5.7 pH is a Logarithmic  pH 7  0.0000001 M Scale  pH 6  0.000001 M  pH 5  0.00001 M  pH 4  0.0001 M The difference  pH 3  0.001 M between each value  pH 2  0.01 M is 10-fold  pH 1  0.1 M  pH 0  1M  Strong and Weak Acids Examples of Strong Acids Examples of Weak Acids HCl Hydrochloric acid HClO Hypochlorous acid HNO3 Nitric acid HNO2 Nitrous acid H2SO4 Sulphuric acid H2SO3 Sulphurous acid HI Hydroiodic acid HF Hydrofluoric acid CH3COOH Ethanoic acid H2CO3 Carbonic acid How is a strong acid defined? Any questions: You can type in the chat function box during this live session (synchronous)? Or onto the Question Board in the Biochemistry Blackboard module and I will look at them later (asynchronous).  Strong and Weak Acids cont… HCl H+ + Cl- acid + H2O HCl (l) + H2O (l) Cl- (aq) + H3O+ (aq) Hydronium ion CH3COOH CH3COO- + H+ acid + H2O CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) Hydronium ion Calculation of pH for weak and strong acids STRONG acids: dissociation reaction goes to completion – no unreacted acid remains in solution Reaction is: HX + H2O H 3 O+ + X − usually simplified to: HX H + + X− To calculate pH only need to know concentration of the acid (HX) present  Calculation of pH for weak and strong acids WEAK acids: dissociation reaction does not go to completion – unreacted acid remains in solution An equilibrium is reached between the hydrogen ions and the conjugate base Reaction is: HX ⇌ H+ + X− For methanoic acid: HCOOH(aq) ⇌ H+ + HCOO− To calculate pH need to know balance of equilibrium reaction (equilibrium constant) :- here called ACIDITY constant (Ka) Calculation of pH for weak and strong acids If reaction is: HX ⇌ H+ + X− The Acidity constant (Ka) ≡ Equilibrium constant (Keq) Ka Square brackets [ ] indicate concentration of a molecule STRONG ACID WEAK ACID HCl Cl - + H+ CH3COOH CH3COO - + H+ 50 50 HCl Cl - + H+ CH3COOH CH3COO - + H+ 0 50 50 49 1 1  Calculation of pH for weak and strong acids If reaction is: HX ⇌ H+ + X− Ka pH = -log[H+] Acidity constant, Ka, is NOT the same as pH Higher the Ka value the stronger the acid − because less of HX left Weak acids have small Ka values − lots of HX left at equilibrium of dissociation Most organic acids are weak acids Can look up Ka values for acids (or pKa values)  Calculation of pH for weak and strong acids When calculating the pH it is assumed that the water does not provide any hydrogen ions simplifies calculation concentration of H+ provided by water (1×10−7 Molar) is insignificant Any questions: You can type in the chat function box during this live session (synchronous)? Or onto the Question Board in the Biochemistry Blackboard module and I will look at them later (asynchronous).  Bases A strong base is a base which hydrolyzes completely, raising the pH of the solution towards 14 Arrhenius bases are water-soluble and donate hydroxide ions (OH-) Alkalis are bases. However, not all bases are alkalis Alkalis are Arrhenius bases and hydroxides of the alkali metals e.g. sodium, potassium Strong bases: NaOH, KOH, Ca(OH)2 Weak bases: NH3 (ammonia)  Bases and pH The strong base sodium hydroxide ionizes into hydroxide and sodium ions in solution: NaOH → Na+ + OH− Pure water dissociates : pH = -log[H3O+] 2H2O(l) → H3O+(aq) + OH−(aq) When these are mixed the H3O+ and OH− ions combine to form water molecules: H3O+ + OH− → 2 H2O The hydroxide ion ‘removes’ the available hydronium / hydrogen ions, lowering their concentration, so pH value goes up  Bases and pH Relationship of [H3O+] and [OH-] and pH is via water self-ionisation Equilibrium constant for water ionisation pH = -log[H3O+] Kw=[H3O+][OH−]=1.0×10−14 pOH = -log[OH-] pKw=pH+pOH=14 So to calculate pH of a basic solution: pH = 14 - pOH e.g. 0.5 M NaOH in water gives a pH of: (strong base so fully ionised) [OH-] = 0.5 M; pOH = -log 0.5 = -(-0.30) = +0.3 pH of the solution is 14 - (+0.3) = 13.7  pH is concentration of H ions +  pH 7  0.0000001 M  0.0000001 M  pH 7  pH 6  0.000001 M  0.00000001 M  pH 8  pH 5  0.00001 M  0.000000001 M  pH 9  pH 4  0.0001 M  0.0000000001 M  pH 10  pH 3  0.001 M  0.00000000001M  pH 11  pH 2  0.01 M  pH 1  0.1 M  0.000000000001 M  pH 12  pH 0  1M  0.0000000000001 M  pH 13  0.00000000000001 M  pH 14 pH = -log[H+]  Neutralization the reaction between an acid and a base will produce a salt and neutralized base hydrochloric acid and sodium hydroxide form sodium chloride and water: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) H+ (aq) + OH- (aq) H2O (l) Neutralization will not always give a solution with pH 7.0 – Only if similar strength acids and bases reacted – a strong acid and a weak base will give a weakly acidic salt and vice-versa  Buffer solution - solution which resists large changes in pH when small volumes of acids or bases are added. Buffers consist of either: a weak acid and its salt or a weak base and its salt Buffers are essential to life and biochemical reactions  Buffers Buffer solution: an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid Buffers mean that pH of solution changes very little when a small amount of acid or base is added to it Buffer solutions are used as a means of keeping pH at a nearly constant value in biochemical applications Buffers are in cells and blood to maintain physiological pH – E.g. blood plasma is at pH 7.4 via bicarbonate-carbonic acid (will be covered in your Human Physiology module)  What happens to the pH if…? CH3COOH CH3COO + H - + pH = -log[H+] 49 999 1 1 CH3COOH CH3COO - + H+ H+ 49 999 1 1 (a) added OH- CH3COOH is what? added (b) (In a non-buffered solution) Any questions: You can type in the chat function box during this live session (synchronous)? Or onto the Question Board in the Biochemistry Blackboard module and I will look at them later (asynchronous).  Buffer solution - solution which resists large changes in pH when small volumes of acids or bases are added. So Buffers consist of either: a weak acid and its salt or a weak base and its salt CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq) 50 000 50 000 Un-ionised Ethanoate Hydronium Ethanoic Acid ions ions The added salt here could be sodium ethanoate CH3COO- Na+  Buffers CH3COOH(aq) + H2O(l)  CH3COO- Na+(aq) + H3O+(aq) When H+ are added to solution, some react with CH3COO-, equilibrium moves to the left a little When OH- are added to solution, equilibrium moves to the right (as hydrogen ions are removed) Thus, some of the added acid or base is neutralized in shifting the equilibrium So the pH of solution changes by less than it would if it were not buffered REMEMBER pH is Concentration of FREE hydrogen ions Addition of Acid (H+) to a Buffer Solution. H+ CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq) 50 000 50 000 Un-ionised Ethanoate Hydronium Ethanoic Acid ions ions CH3COO- (aq) + H+ (aq) CH3COOH(aq) Addition of Base (OH-) to a Buffer Solution. OH- CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq) 50 000 50 000 Un-ionised Ethanoate Hydronium Ethanoic Acid ions ions CH3COOH (aq) + OH- (aq) CH3COO-(aq) + H2O(l) and H3O+ (aq) + OH- (aq) H2O(l)  Buffers CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq) When H+ are added to solution, equilibrium moves to the left (as hydrogen ions are on righthand side of equilibrium equation) When OH- are added to solution, equilibrium moves to the right (as hydrogen ions are removed) Thus, some of the added acid or base is neutralized in shifting the equilibrium So, the pH of solution changes by less than it would if it were not buffered  The Henderson–Hasselbalch equation Useful for estimating the pH of a buffer solution and finding the equilibrium pH in acid-base reactions Where pKa is −log(Ka) Ka negative log of the acid dissociation constant Optimal buffering capacity is when pH = pKa So when [base] = [acid]  14 Halfway to end-point 12 Buffering 10 8 End-point pH 6 At halfway point solution has equal amounts of acid 4 and salt molecules. 2 Thus, can best act as a buffer to prevent changes in 0 pH 1 2 3 4 5 6 7 8 Volume of base Hydroxide / ml added CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq) 50 000 50 000  14 12 Halfway to end-point Buffering 10 So at halfway point the pH 8 scale gives the pKa value pH 6 4 2 0 1 2 3 4 5 6 7 8 Volume of base Hydroxide / ml added CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq) 50 000 50 000  So what is the buffering range here? 14 Buffering capacity of 12 solution is largest when 10 pH=pKa pH 8 pKa 6 Amino acids have a pKa 4 values 2 Thus, proteins have a pKa 0 1 2 3 4 5 6 7 8 values Volume of base Hydroxide / ml added  Buffers in Biology Resistance to changes in pH make buffer solutions essential for many biochemical processes An ideal buffer for a particular pH has a pKa equal to that pH as such a solution has maximum buffer capacity Buffer solutions are necessary to keep the correct pH for enzymes to work - many enzymes work only under very precise conditions A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3−) is present in blood plasma - maintains a pH between 7.35 and 7.45 Any questions: You can type in the chat function box during this live session (synchronous)? Or onto the Question Board in the Biochemistry Blackboard module and I will look at them later (asynchronous). MCQ quiz for Lecture 5: Fundamentals of Chemistry part 4 Answers will be given in your Seminar sessions – with further discussion. You must attempt before your seminar session. These quizzes are part of your learning for the Biochemistry module They will aid your on-going studies at the University of Westminster Q1) pH is calculated via which of the following? a) +Log10 H+ ion concentration b) +Log10 OH- ion concentration c) -Log10 H+ ion concentration d) -Log10 OH- ion concentration e) +Log10 H2O concentration Q2) If a solution is described as a weak base, its properties in will be? a) pH > 7 and partially ionised b) pH < 7 and partially ionised c) pH = 7 and partially ionised d) pH > 7 and fully ionised e) pH < 7 and fully ionised Q3) You are using a hydrochloric acid solution of pH2.5. The closest estimation of the original hydrochloric acid concentration is? a) +2.5 Molar b) +300 milli-moles c) +0.003 moles d) -300 milli-Molar e) +300 milli-Molar Q4) Which of the following statements about the acidity constant, Ka, is true? a) Ka is the same as pH. b) pKa is the same as pH. c) If an acid has a very high Ka value, the acid is only very slightly ionised in solution. d) The lower the Ka value the weaker the acid. e) pKa is always 7 in buffer solutions. Q5) Which of the following statements is incorrect? a) Acids can be defined as proton donors, independent of water. b) Buffers are very important in biochemical reactions, allows enzymes to function. c) Not all bases are derived from the alkalis metals. d) All the equations given in the lecture need to be memorised. e) Water which has been left exposed to the air will not be not pH7.