In order to continue enjoying our site, we ask that you confirm your identity as a human. Thank you very much for your cooperation. Show
Explanation for the (left) curve involving strong acid being titrated with weak base: Suppose the reaction: HCl(aq) + NH4OH(aq) -> NH4Cl(aq) + H2O(l) Initially, the electrical conductivity is high due to the dissociation of the strong acid into H+ ions and its conjugate base ions. Anyways, back to the left graph, as you add the weak base, the OH– dissociated from the base molecules will react with the H+ dissociated from the acid molecules to form undissociated water molecules. This means that the overall electrical conductivity of the solution will decrease due to the decreased total concentration of free ions in solution. The minimum concentration of total ions is when equivalence point has been reached where the OH– ions neutralised all the H+ ions that is in solution. Upon adding excess weak base after equivalence point has been reached, the excess weak base will react with water to form OH– ions in solution (by accepting H+) from water. However, as seen from the graph, the total electrical conductivity of the solution will NOT increase upon adding excess weak base after the equivalence point. The reason for this is due to the common ion effect. When equivalence point has been reached, there is high concentration of NH4+ (and Cl–) ions in solution. Upon adding more weak base (NH4OH) into the solution, there will be an increase in [NH4+] ion in solution. As we have learnt in Module 5, according to Le Chatelier’s Principle, due to the addition of the common ion in solution, the system will shift the equilibrium position of the following reaction to the left to minimise the increase [NH4+] ion in the solution. NH4OH(aq) <-> NH4+(aq) + Cl–(aq). Therefore, more undissociated NH4OH molecules will remain in solution and therefore lowering the solubility of NH4+ ions being added into solution. Due to common ion effect, it effectively suppresses the dissociation of NH4OH molecules that are being added into solution. This therefore means that electrical conductivity won’t increase (at least not to significant extent)
Explanation for the (right) curve involving weak acid being titrated with weak base: Suppose the reaction: CH3COOH(aq) + NH4OH(aq) -> CH3COONH4(aq) + H2O(l) Initially there is a decrease in conductance of the solution as the dissociated H+ ions from the weak acid reacts with the OH– ions that is dissociated from the weak base. However, as the CH3COONH4 is formed and is dissociated into its constituent ions in solution (CH3COO– and NH4+), the electrical conductivity of the solution increases upon the equivalence point is reached where no more CH3COONH4 ions are formed and dissociated. Notice that if you add more of the weak acid after the equivalence point, the conductivity WILL NOT increase due to the common ion effect as explained in the strong acid – weak base titration scenario. However, if you add more of the strong acid after the equivalence point, such as NaOH, then the conductance of the solution will increase! Learning Objective #3 - Calculate and apply the acid dissociation constant (Ka) and pKa to determine the difference between strong and weak acidsWe have briefly introduced on Ka in Module 5 under different types of Keq or equilibrium constants. In this learning objective, we will dive into greater detail and explore Ka, i.e. the acid dissociation constant. So, Ka is one type of equilibrium constant that is used to determine the extent of dissociation of an acid in solution (e.g. water). Like other equilibrium constant, Ka is basically the concentration of the products divided by the concentration of the reactants. For a generic acid with the formula HA, its Ka can be written as: Ka = [H+] [A–] / [HA] ; where A- is the conjugate base of acid HA. Last week, we have already explored the difference between strong and weak acids. Strong acids ionise completely into its constituent ions in water whereas weak acids have less than 100% degree of ionisation. For strong acids, this would mean that the concentrations of [H+] and [A–] would much much higher the concentration of [HA], where the [HA] is very small. This means that a strong acid have a very high Ka value compared to weak acid. If you recall, from the previous week’s note, the ‘p’ in front of H in ‘pH’, effectively means -log10. So, pKa means -log10 x Ka. So,
Base dissociation constant = Kb = [BH+] [OH–] / [B] ; where B+ is the conjugate acid of base B Due to the nature of logarithmic functions, as the value of Ka increases, the value of pKa decreases. You can substitute different Ka values into your calculator to test this inverse relationship out. Since strong acids have large Ka values, they have a smaller pKa value that weak acids. We will have homework questions at the end of this week’s notes for you to calculate pKa when given Ka and vice versa. But here is the a formula that you need to know when you are given pKa and you need to determine Kb. We talked about the autoionisation of water, Kw, last week. Now that we have learnt, Ka and Kb, we can also express Kw as: Kw = Ka x Kb = 1.0 x 10-14 (ONLY at 25 degrees celsius as Kw changes with temperature like all other equilibrium constant which we have explanation in the derivation of Keq section in Module 5 notes)
Yes, there is also a pKw.
Learning Objective #4 - Describe the importance of buffers in natural systemsA buffer system is a solution comprised of approximately equal amounts of weak acid and its conjugate base OR approximately equal amounts of a weak base with its conjugate acid. The importance and use of buffer are derived from its ability to resist changes in pH when small amounts (volume) of strong acids or base are added into the buffered system. Only small amounts of strong acid or base can be added in order for the buffer to resist changes in pH. That is, the strong acid or base that is being added into the buffer is a limiting reagent. This is because if excess amounts strong or base are added, all of the weak acid/base and its conjugate base/acid will be consumed, effectively exceeding the buffer capacity to resist changes in pH. Buffer capacity refers to the total amount of acid or base added into the buffer and where any significant changes in pH is resisted using the weak acid/base and its conjugate base/acid components in the buffer.
Sometimes you see that buffer capacity also being defined as the number of moles of H+ or OH- required to be added into the buffer system to increase or decrease the buffer’s pH by 1 for a one litre of the buffer solution.
Using Le Chatelier’s Principle, it is possible to account for how buffers are able to to resist changes in pH when small amounts of acid or base are added into the buffer system (or buffer solution). Let’s explore two non-specific examples before looking into a specific buffer system in our blood. Weak acid - Conjugate Base BufferHA(aq) + H2O (l) <-> H3O+ (aq) + A- (aq) Situation 1: Adding strong acid Suppose if you add small amounts of strong acid into the weak acid-conjugate base buffer, the added hydronium ions will react with the conjugate base in the buffer (A-) to shift the equilibrium position to the left to minimise the increased concentration of hydronium ions, forming more weak acid (HA(aq)) as per Le Chatelier’s Principle. This effectively means that if you add a strong acid into the weak acid buffer solution, you will not see the pH rise as much as you would have expected. Situation 2: Adding strong base Suppose if you add small amounts of strong base into the weak acid-conjugate base buffer, the added hydroxide ions will react with the hydronium ions which results in the increase in pH of the buffer system as there is less hydronium ions in solution. This is what you will expect when you add a base into any solution containing an acid which results in a increase in pH by consuming the hydronium ions.
Weak base – conjugate acid buffersB(aq) + H2O (l) <-> BH+ (aq) + OH–(aq) Situation 1: Adding strong base Suppose if you add small amounts of strong base into the weak base-conjugate acid buffer, the added hydroxide ions will react with the conjugate acid (BH+) to shift the equilibrium position to the left to minimise the increased concentration of hydroxide ions, forming more weak base (B aq) as per Le Chatelier’s Principle. This effectively means that if you add a strong base into the weak base buffer solution, you will not see the pH rise as much as you would have expected. Situation 2: Adding strong acid Suppose if you add small amounts strong acid into the weak base-conjugate acid buffer, the added hydronium ions will react with the hydroxide ions which results in the decrease in pH of the buffer system as there is less hydroxide ions in solution. This is what you will expect when you add an acid into any solution containing a base which results in a decrease in pH by consuming the hydroxide ions.
Specific Buffer Example (Natural Buffer): Carbonic acid - hydrogen carbonate bufferThe carbonic acid – hydrogen carbonate buffer occurs in a natural system which is in our blood. This buffer allows the pH of our blood to remain in a constant range between 7.3 – 7.4. The chemical expression of this buffer is: H2CO3(aq) + H2O(l) <-> H3O+(aq) + HCO3-(aq) Note that carbonic acid is a weak acid and its conjugate base is the hydrogen carbonate ion. When cells perform cellular respiration, carbon dioxide is a waste product that is produced. This carbon dioxide diffuses through the cells’ membranes and into the blood plasma (or blood stream). Since blood plasma is mainly water, the carbon dioxide that dissolved in the blood plasma will react with water to carbonic acid where the chemical reaction can be expressed as: CO2(g) + H2O(l) <-> H2CO3(aq). Shifting the carbonic acid – hydrogen carbonate bufferThe kidney has the function of filtering the blood for nitrogenous waste which are mostly acidic. The kidney absorbs hydronium ions from the blood and secretes hydrogen carbonate (HCO3-) ions back into the blood to be absorb as the hydrogen carbonate ion is important in assisting the organism’s digestion processes. Comparatively, the hydrogen ion (or hydronium ion) is absorbed from the blood and secreted out the blood as a component of nitrogenous waste (urea). When the kidney absorbs hydronium ions from the blood, it will shift the buffer system to the right as per Le Chatelier’s Principle to minimise the decrease in [H3O+]. This would mean that more carbonic acid (weak acid) will ionise in water. The overall effect of this buffer allows the change in the blood pH to be minimised and be maintain within a narrow range of 7.3 – 7.4. When the kidney reabsorbs hydrogen carbonate ions by secreting it into the blood, it will shift the buffer system to the left as per Chatelier’s Principle to minimise the increase in [HCO3-] where the absorbed hydrogen carbonate ions will react with hydronium ions to form carbonic acid. The carbonic acid molecules will dissociate into carbon dioxide and water where the carbon dioxide will be secreted out of the body when exhaled as increasing the [H2CO3] will shift the equilibrium position of CO2(g) + H2O(l) <-> H2CO3(aq) to the left as per Le Chatelier’s Principle. The overall effect of this buffer allows the change in the blood pH to be minimised and be maintain within a narrow range of 7.3 – 7.4. Importance of maintaining the pH within a narrow rangeEnzymes are examples of biological catalyst which are in the human body. Enzymes are operate most optimally around a specific pH range and, if the pH of the surrounding environment which it occupies in deviates too much from its specific pH range, it will denature. This means that the enzymes’ shape and active site will be be altered such that it will no longer be able to catalyst metabolic processes in the body. Failure to catalyst metabolic processes can have many consequences such as accumulation of wastes, inability to transport nutrients in and out of cells, etc. If the pH of the blood deviates from 7.3-7.4 for prolonged performs of time, the inability for the required chemical reactions to be catalysed could result in death.
Thus the blood’s pH should be maintain around 7.4 where the enzymes can function optimally and catalyse the appropriate (and necessary) chemical reactions. Therefore, the carbonic acid-hydrogen carbonate buffer is essential in the blood to maintain the wellbeing of the organism (e.g. human). Learning Objective #5 - Conduct a practical investigation to prepare a buffer and demonstrate its propertiesTo prepare a buffer, the Henderson-Hasselbalch Equation is useful in assisting us in calculating the concentration of weak acid/base and its conjugate base/acid to make an effective buffer with high buffering capacity at a given pH level. Originally, Henderson-Hasselbalch equation was used to help chemist approximate the pH of a buffer which can be expressed as: pH = pKa + log10 [conjugate base] / [weak acid]
pOH = pKb + log10 [conjugate acid] / [weak base]
Creating a buffer with the pH of 7.4Since we have explored the carbonic acid – hydrogen carbonate buffer that exist in a pH of 7.4, let’s keep things in our way and attempt to create a buffer with a pH of 7.4. Every buffer systems functions most effectively when its buffering range is +/- 1 of the weak acid’s pKa. For example, we cannot create an effective acetic acid – acetate buffer at a pH of 7.4 as the Ka of acetic acid is 1.75 x 10^-5 at 25 degrees celsius. This means that its pKa of acetic acid is 4.77. This means an effective acetic acid-acetate buffer can only be prepared at a pH of 3.77 – 5.77 and not at a pH of 7.4. The most effective buffer is one where its pH = pKa. This means that, from the Henderson-Hasselbalch equation, the most effective buffer is when the concentration of the conjugate base is equal to the concentration of weak acid. This is because log[1] = 0. This means pH = pKa.
An effective buffer is when pH is +/- 1 of the pKa. Therefore, the pKa1 for carbonic acid chosen to make the buffer with the pH of 7.4. pH = pKa1 + log[HCO3-]/[H2CO3] 7.4 = 6.4 + log[HCO3-]/[H2CO3] log[HCO3-]/[H2CO3] = 1.0 [HCO3-]/[H2CO3] = 10^(1) = 10 NOTE: Notice that the above ratio DOES NOT tell you the volume of carbonic acid you need to use to make the buffer solution. Please have a look at the Solution to Question #10 of this week’s homework set (located after this week’s notes) to see an example of how you can prepare a buffer in practice by calculating the exact volumes of solution that you need to add to make a buffer solution at school. So, we can make a carbonic acid – carbonate buffer with a pH of 7.4 by making a solution that has 1M of hydrogen carbonate ions and 0.1M of carbonic acid. Since the volume cancels out, we can say that we can make the buffer using 1 mole of hydrogen carbonate ions and 0.1 moles of carbonic acid. This means that pH does not depend on volume so diluting the buffer solution will not affect pH as both the conjugate’s concentration and weak acid/base are affected (changed/diluted) equally* * This is true in theory when examining the Henderson-Hasselbatch equation but not in reality. This is because, in practice, as dilution increases (adding more and more distilled water), more of the pH will be determined by the autoionisation of water instead by the [conjugates] and [weak acid] or [weak base]. This is because you are adding in more water into the solution and the solution is mostly water. Water will autoionise into [H+] and [OH-] in equal amounts so, over time, as the solution gets more and more dilute, the buffer solution will approach a pH of 7. This means that the Henderson Hasselbatch equation on the pH of buffers do not hold when the solution is very dilute or very strong (strong acid/base). NOTE: Dilution will have less effect on buffer solutions than solutions that are not buffers. NOTE: if the solution is NOT a buffer, then dilution will decrease the concentration of the hydronium ions and increase the pH as pH is dependent on the concentration of hydronium ions. If you continue adding distilled water, the pH will approach 7 but never exceed 7 because distilled water used for the dilution is neutral (pH = 7). The diluted solution will also never be equal to exactly 7 as there are acids present in the solution. Learning Objective #6 - Explore acid/base analysis techniques that are applied to:- In Industries- By Aboriginal and Torres Strait Islander Peoples- Using digital probes and instrumentsIn Industries - Wine IndustryMost of the wine solutions that you encounter comprises of dilute, weak organic acids. These acids in wine can be divided into two areas, those being fixed and volatile acids. These two division of acids determines total acidity of a wine solution.
During the fermentation process of manufacturing wine, volatile acids (such as acetic acid) can be formed as a result of bacteria present in the wine. The amount of volatile acid in wine is called volatile acidity. It is important to analyse the volatile, organic acid content in wine as their concentration is proportional to the susceptibility of the wine being spoiled. This is because if bacteria enters the wine, the higher the concentration of volatile organic acids that is present in the wine, the more easily the bacteria can convert alcohol into acetic acid (example of volatile acid). This is called the spoiling of wine.
Altogether, the volatile and non-volatile acids give the tartiness or sharpness of the wine’s taste. The higher the acid concentration in the wine, the more sharp the wine will taste. Acetic acid and other volatile acids can be analysed for their concentration in wine by using steam distillation (technique). After steam distillation, the volatile acids will be the distillate. The volatile acids can be to titrated against a standardised NaOH solution to determine the concentration of the volatile acids. Typically, three drops of phenolphthalein indicator is used in the titration. The remaining solution can be then analysed to determine the components of fixed acids in wine using thin layer chromatography (technique) or paper chromatography (technique). The alcohol content is required to be specified on the wine as per legalisation. Alcohol content such as ethanol present in wine can be determined using liquid gas chromatography (technique) or by using hydrometry. Using a hydrometer, different wine with known ethanol densities (i.e. wine with different known amount of ethanol) are measured against their boiling point. This means that a density calibration curve can be constructed which consists of boiling (Y-axis) against of the densities for a range of wine with known ethanol densities, % by mass (x-axis). From this, the calibration curve can be used to determine the unknown ethanol concentration in a wine sample. Wines also have added sulfur dioxide used as a preservative to prevent the growth of bacteria and yeast. Sulfur dioxide also inhibit the oxidation of wine by acting as an antioxidant to keep the wine smelling fresh. If oxidation of wine occurs, it would lower the quality of wine and this is often known as a wine fault in the industry. When sulfur dioxide is added into the wine, most of them will react with organic aldehyde and ketone acid molecules present in the wine, which you will learn in Module 7. The rest of the sulfur dioxide will react with water in the wine to form bisulfite, sufite and hydrogen ions. SO2(g) + H2O(l) <-> H+ (aq) + HSO3-(aq) <-> 2H+(aq) + SO32-(aq) HSO3–= bisulfite ion SO32-= sulfite ion It is important to determine the concentration of sulfur dioxide in wine because there is a limit for amount allowed in wines. As sulfur dioxide produces a strong, undesirable aroma that covers the smell and flavour of the wine, the reduction of sulfur dioxide improves the smell and taste of the wine. Furthermore, some of the sulfur dioxide can dissolve in the wine to produce sulfite ions which cause skin rashes. Therefore, by lowering the content of sulfur dioxide in the wine production process, there is also health benefits for wine consumers. To determine the total [SO2] in the wine, titration is used. First, the sodium hydroxide is added into the wine sample to decompose bisulfite molecules (freeing SO2). Next, the solution is acidified with sulfuric acid. Finally, redox titration (technique) is performed with iodine as a titrant and starch as an indicator. Description of the mentioned acid-base analysis techniquesSteam distillation – involves adding wine sample to boiling water in a flask that is being heated (i.e. using a heating mantle, water bath, etc) to evaporate volatile acids which are then collected as distillate using a condenser. Thin layer chromatography – involves the separation of substances according to their solubility as they distribute themselves throughout the stationary phase by travelling in a mobile phase. A silica gel is typically used as the stationary phase. The non-volatile acids that we want to test are placed on the silica gel whereby the gel is then partially submerged in an appropriate solvent (e.g. water). The more soluble the acid is, the higher up the silica gel in which the acid will travel. From this, the different components in the sample can be separated based off solubility. Paper chromatography – The procedure is the same as thin layer chromatography but the stationary phase is not silica gel but a chromatography paper is used instead. The solvent for paper chromatography is typically water. Liquid gas chromatography –The procedure is the same as the two chromatography techniques mentioned above. However, the stationary phase is typically a liquid inside a chamber and the mobile phase is a gas. The sample is inserted into a separated heated chamber whereby the sample’s component can be separated according to their mass. This is because the components that are heavier will travel at a slower speed thus will be detected by the detector at a later point in time compared to the component in the sample that are of lower molecular mass. This is a quantitative measure of the components so compared to the two chromatography techniques mentioned above which are qualitative. Hydrometry – procedure already discussed in the above learning objective involving the use of hydrometer and calibration curve. Titration – already discussed in Learning Objective #1 in this week’s notes. Redox Titration – Normal titration, however, the titrant and analyte has an oxidant and reductant relationship such that a redox reaction can occur. The titrant does not necessarily need to be the oxidant or reductant. In Industry - Juice IndustryAnother example of acid-base analysis techniques used in juice industry in the manufacture of orange juice. Titration can also used to determine the acid content in juices such as ascorbic acid in orange juice by titrating it against a strong base such as hydrochloric acid. Once the concentration of the ingredients are identified the correct nutritional information can be printed on the product’s nutrition label, allowing customers to select according to their nutritional demand. This is important because consumers expect that the stated ingredients that are either contained or removed from the product and their concentration to accurate. If not, the manufacturer may face legislative risk from their customers. By Aboriginal and Torres Strait Islander PeoplesAboriginal and Torres Strait Islander Peoples demonstrated their technique of using acids and bases through their choice of selecting specific fruits and plants to meet their specific needs. By performing volumetric analysis such as titration, chemists are able to determine the nature and concentration of various substances (acids and bases) that are contained within the fruits and plants used by Aboriginal and Torres Strait Islander Peoples. The Davidson plum is a natural Australian fruit that approximately 100 times more ascorbic acid (vitamin C) than contained in an orange. For such reason, it has a very sour taste due to the high concentration of acid. Aboriginal and Torres Strait Islander Peoples consumed the Davidson plum as way to boost their body’s nutrient level which reduced their chance of having scurvy disease. When exposed to water, the soap tree’s leaves is able to lather or produce a soap solution that have antibacterial properties and thus act as an antiseptic. The reason for this is that the leaves contain saponin acid which has the ability to suppress bacteria growth. Aboriginal and Torres Strait Islander Peoples used the soap tree leaves as a way to heal cuts on their skin. Aboriginal and Torres Strait Islander Peoples also used yellow ochre (hydrated iron hydroxide) to treat stomach upsets. The chemistry behind is that the yellow ochre is basic and thus can react and neutralise with any excess hydrochloric acid in the stomach. This served as a way for Aboriginal and Torres Strait Islander People to remove any heartburns or stomach upsets. The yellow ochre has base similar to antacids which is an example of medicine used in everyday life as mentioned earlier in Module 6’s notes. The Grey Mangroves are used to treat stingray injury by preventing infection and neutralise the mildly acidic stingray venom. This is done by smashing the Grey Mangroves’s leaves and adding water to create a mixture that is a base which can be applied to the wound caused by the stingray. Digital probes and instrumentsPlease refer to previous week’s notes on using pH probe and meter to determine the acidity of a range of solutions Learning Objective #7 - Conduct a chemical analysis for a range of common household substances for its acidity or basicity, for example:- Soft drink- Wine- Juice- MedicineTitration can be produced on soft drinks, wine, juice and medicine (e.g. aspirin tablets) to determine its composition and pH. In this learning objective, we will briefly examine the acid or base content in each of the four substances. A youtube video will be released in Term 1 showing the titration of each of the four substances to determine their composition and pH. Soft DrinksThe carbon dioxide inside soft drinks react with water to form carbonic acid. It is the carbonic acid that gives soda drinks their tartiness or sharpness. Flavours are then added into this carbonated water, giving soft drinks their range of tangy flavours. The higher the carbonic acid concentration, the sharper the soft drink will taste. Since carbonic acid is acidic, soft drinks are acidic. In order to dissolve carbon dioxide in water to occur, or carbonation, it is accompanied with low temperature and high pressure. The high temperature allows carbon dioxide to be dissolved in water and low temperature enhances the solubility of gases such as carbon dioxide. You can relate this back to Le Chatelier’s Principle, a principle that we explored in Module 5. CO2(g) <-> CO2(aq) [Equation 1] CO2(aq) + H2O(l) <-> H2CO3(aq) [Equation 2] H2CO3(aq) <-> H+(aq) + HCO3-(aq) [Equation 3] The forward reaction for both equation 2 and equation 1 are both exothermic. Therefore, increasing temperature will shift both Equation 1 and 2 to the left. This will decrease the concentration of carbonic acid in solution and therefore resulting in the soda going flat (losing its sharp taste). Increasing pressure will cause the equilibrium reaction in equation 1 to shift to the right. This would mean that the carbon dioxide gas will dissolve in water. The increase in [CO2] will shift equation 1 to dissolve the carbon dioxide in water which shifts equation 2 to the right to form carbonic acid. Vice versa, if you open the can of soda, it will cause equation 1 to shift the left to order to increase gas concentration and thus shift both equation 2 and 3 to the left. This lowers the concentration of carbonic acid and, overtime, the soda can will taste flat.
|
Why does the addition of extra distilled water to the flask not affect the amount of sodium?
Postagens relacionadas
direito autoral © 2024 estudarpara Inc.
