Biochemistry Lab - 01 pH and Buffer Preparation

BIOCHEMISTRY LABORATORY MODULE 1 - pH OF COMMON SOLUTIONS AND BUFFER PREPARATION pH • • • pH is measured using pH paper or a pH meter. Most fruits, vegetables, and body fluids are acidic. While pure water is neutral, natural water may be either acidic or basic. Cleaners tend to be basic. Not All Liquids Have a pH Value, pH only has meaning in an aqueous solution (in water). Many chemicals, including liquids, do not have pH values. If there's no water, there's no pH. For example, there is no pH value for vegetable oil, gasoline, or pure alcohol. • pH- the acidity and basicity of a substance can be measured in terms of the molar concentration of the hydrogen (H+) or hydroxide (OH–) ions, respectively pH OF COMMON ACIDS • • • • pH paper o the pH of a solution may be measured using pH paper o pH paper measures the pH of the solution by changing its color depending on the concentration of H+ and OH– ions o this paper is treated with an indicator, which is flavin, an anthocyanin commonly found in red cabbages o flavin usually turns: ▪ red if the solution is acidic ▪ greenish if the solution is neutral ▪ purple if the solution is basic pH meter pH meter o a more accurate method than pH paper o contains a probe that passes the electrical signals to the meter itself which interprets the electrical signals and display the results on its screen pH meter probe o two electrodes: ▪ sensor electrode (contains a pH 7 buffer solution) ▪ reference electrode (usually contains saturated KCl solution) o made up of porous glass; have both a silver wire covered with silver chloride or calomel o when the probe is immersed in a solution, the hydrogen ions accumulate around the bulb of the electrode, generating an electrical flow captured by the silver wire o the voltage of this electrical flow is measured by the meter and then converts it to pH value LESSON 1 - pH OF COMMON SOLUTIONS pH • • A measure of how acidic or basic a chemical is when it's in an aqueous (water) solution. o A neutral pH value (neither an acid nor a base) is 7. o Substances with a pH greater than 7 up to 14 are considered bases. o Chemicals with a pH lower than 7 down to 0 are considered acids. o The closer the pH is to 0 or 14, the greater its acidity or basicity, respectively. Here's a list of the approximate pH of some common chemicals. pH is a measure of how acidic or basic an aqueous solution is. pH usually ranges from 0 (acidic) to 14 (basic). A pH value around 7 is considered neutral. Fruits and vegetables tend to be acidic. Citrus fruit, in particular, is acidic to the point where it can erode tooth enamel. Milk is often considered to be neutral since it's only slightly acidic. Milk becomes more acidic over time. The pH of urine and saliva is slightly acidic, around a pH of 6. Human skin, hair, and nails tend to have a pH of around 5. pH 0 2.0 2.2 Common Acids Hydrochloric Acid (HCl) Battery Acid (H2SO4 sulfuric acid); stomach acid Lemon Juice Vinegar 3.0 3.0 to 3.5 3.5 to 3.9 4.0 Apples, Soda Sauerkraut Pickles Wine and Beer 4.5 4.5 to 5.2 around 5.0 5.0 Tomatoes Bananas Acid Rain Black Coffee 1.0 NEUTRAL pH CHEMICALS • • Distilled water tends to be slightly acidic because of dissolved carbon dioxide and other gases. Pure water is nearly neutral, but rainwater tends to be slightly acidic. Natural water rich in minerals tends to be alkaline or basic. 7.0 - Pure Water pH OF COMMON BASES • Many common cleaners are basic. Usually, these chemicals have a very high pH. Blood is close to neutral but is slightly basic. pH Common Bases 7.0 to 10 Shampoo 7.4 Human Blood 7.4 Human Tears 7.8 Egg around 8 Seawater 8.3 Baking Soda (Sodium Bicarbonate) around 9 Toothpaste 10.5 Milk of Magnesia 11.0 Ammonia 11.5 to 14 Hair Straightening Chemicals 12.4 Lime (Calcium Hydroxide) 13.0 14.0 Lye Sodium Hydroxide (NaOH) OTHER pH VALUES • • • • • • • • There are multiple ways to test the pH of substances. The simplest method is to use pH paper test strips. You can make these yourself using coffee filters and cabbage juice, use Litmus paper, or other test strips. The color of the test strips corresponds to a pH range. Because the color change depends on the type of indicator dye used to coat the paper, the result needs to be compared against a chart of standard. Another method is to draw a small sample of a substance and apply drops of pH indicator and observe the test change. Many home chemicals are natural pH indicators. pH test kits are available to test liquids. Usually, these are designed for a particular application, like aquaria or swimming pools. pH test kits are fairly accurate but may be affected by other chemicals in a sample. The most accurate method of measuring pH is using a pH meter. pH meters are more expensive than test papers or kits and require calibration, so they are generally used in schools and labs. pH Equation o The equation for calculating pH was proposed in 1909 by Danish biochemist Søren Peter Lauritz Sørensen: ▪ pH = -log[H+] ▪ where the log is the base-10 logarithm and [H+] stands for the hydrogen ion concentration in units of moles per liter solution. The term "pH" comes from the German word "potenz," which means "power," combined with H, the element symbol for hydrogen, so pH is an abbreviation for "power of hydrogen." ACIDS AND BASE THEORIES • • Soil pH ranges from 3 to 10. Most plants prefer a pH between 5.5 and 7.5. Stomach acid contains hydrochloric acid and other substances and has a pH value of 1.2. Pure water free of undissolved gases is neutral, not much else is. However, buffer solutions may be prepared to maintain a pH near 7. Dissolving table salt (sodium chloride) in water does not change its pH. HOW TO MEASURE pH • • • Arrhenius Theory o acid is any species that can increase the concentration of H+ ions (protons) in an aqueous solution o base is any species that can increase the concentration of OH– in an aqueous solution o however, it is limited in describing the acidity or basicity of a substance as it dissociates in aqueous solution o there are many substances that behave as an acid or base in gaseous phase where no solvent is present and there are those that remain unionized in certain solvents (HCl in benzene, liquid ammonia), and amides (NaNH2) Brønsted-Lowry Theory o an acid is any species that can donate a proton (H+) o a base is any species that can accept a o proton Lewis Theory o an acid is a species that can accept a pair of nonbonding electrons ▪ a base is a species that can donate a pair of non-bonding electrons o suggest that acids react with bases, sharing a pair of electrons, with no change in the oxidation numbers of any atoms LESSON 2 - BUFFER PREPARATION BUFFERS • • most organisms, including humans, need to maintain pH within a fairly narrow range in order to survive (e.g. human blood needs to keep its pH right around 7.4) buffers, solutions that can resist changes in pH, are key to maintaining stable [H+] in biological systems o ↑ [H+], the buffer will absorb some of them, bringing pH back up o ↓ [H+], a buffer will donate some of its own [H+] to reduce the pH typically consist of an acid-base pair, with the acid and base differing by the presence or absence of a proton (a conjugate acid-base pair) Buffer Solutions o Made up of a weak acid and its conjugate base or a weak base and its conjugate acid. The two components maintain a pH balance that resists change when strong acids or bases are added to it. Buffer Preparation o The solution is made by taking a weak acid and adding its conjugate base (which is formed by removing a proton from the same type of acid) or by combining a weak base with its conjugate acid. The use of conjugates is what gives a buffer solution its resistance to pH changes; it creates an equilibrium between the acid and the base which is difficult for other acids or bases to overcome. Even when strong acids or bases are added, the equilibrium between the weak acid/base and its conjugate reduces the impact of the addition on overall solution pH. Examples of Buffers o blood - contains a bicarbonate buffer system o TRIS buffer o phosphate buffer pH range of common buffering agents: o • • • • o Buffer pKa pH range citric acid 3.13., 4.76, 6.40 2.1 to 7.4 acetic acid 4.8 3.8 to 5.8 KH2PO4 7.2 6.2 to 8.2 borate 9.24 8.25 to 10.25 CHES 9.3 8.3 to 10.3 When a buffer solution is prepared, the pH of the solution is adjusted to get it within the correct effective range. Typically a strong acid, such as hydrochloric • • acid (HCl) is added to lower the pH of acidic buffers. A strong base, such as sodium hydroxide solution (NaOH), is added to raise the pH of alkaline buffers. Buffering pH o Buffer solutions have a wide range of applications, both in the real world and in the lab. A buffered pH is required for most enzymes to function correctly, and buffering is used to ensure proper color concentration when using dyes. Buffer solutions are also used to calibrate equipment, especially pH meters that might be miscalibrated if a buffer is not present. It's worth noting that buffer solutions do not necessarily have a neutral pH, just a balanced one; buffer solutions made from citric acid, ammonia, acetic acid (which is found in vinegar in low concentrations) and other compounds can have pH values as low as 2 or higher than 10. This allows the use of buffer solutions in work with very strong acids or bases. Buffer Capacity o While buffer solutions are resistant to changes in pH, this doesn't mean that the pH of a buffer solution can't change if enough strong acid or strong base is added. The amount of a strong acid or base that a buffer solution can take before significant pH changes occur is known as the buffer capacity. The capacity differs depending on the core components of the buffer solution and how much of the strong acid or base is added to the solution. If adding a strong acid to the buffer solution, the capacity is equal to the amount of the base in the solution. If adding a strong base, the capacity is equal to the amount of the acid in the solution • • o pH<7 o Turn blue litmus paper to red Base o Properties is due to hydroxyl ion of hydroxide ion OH o pH>7 o Turns red litmus paper to blue Determination of pH o pH paper - precise, show color of different pH levels o pH meter – electric instrument that use electrode ▪ Provides accurate data o Indicator - Complex organic compound that are of one color if hydrogen ion conc. is above a certain value and of different color it is lower. ▪ Phenolphthalein = a-colorless b- pink ▪ Methyl red = a- red b-yellow ▪ Methyl violet = a- yellow b- blue violet LESSON 3 - pH AND BUFFER PREPARATION pH • • • • • • • • • • Power of hydrogen or Puissance d’ hydrogen One of the most important and frequently used analytical procedures in biochemistry It is the logarithm of the reciprocal of the H ion concentration pH= -log [H+] pH for hydrogen group pOH for hydroxyl group pKw for water =14 Soren Sorensen o Danish chemist o Introducted his pH scale in 1909 it is very valuable and is still used today o It gave doctors, chemists, and gardeners a measurement that could tell them exactly how acid, alkaline or neutral a substance o Doctors for diagnosis o Chemist for their chemical reactions or standardization of solution o Gardeners for the chemical composition of soil… usually must be basic Acid o Its properties are due to hydrogen ion or the hydronium ion H3O • • • • Saliva o Values - 6.2 to 7.6 o Acidic ▪ Happens when bacteria in your mouth break down the carbohydrates you consume, releasing lactic acid, butyric acid, and aspartic acid. ▪ Demineralize (break down) tooth enamel. ▪ Generalized chronic periodontitis o Basic - Generalized chronic gingivitis Urine o Values - 6 o Acidic - Dehydration, Diabetic Ketoacidosis, Diarrhea, Starvation o Basic ▪ pH of 8.5 or 9.0 means presence of urea-splitting organism: Proteus, Klebsiella, or Ureaplasma urealyticum: asymptomatic UTI ▪ Also signify struvite kidney stones/“infection stones.” Sweat o Values - 4.5 and 7.0. o Acidic ▪ Happens when the body is unable to break down acids, or even from exercising too frequently • ▪ Diabetic Ketoacidosis ▪ BODY ODOR o Basic - Happen when low in sodium diet, Dehydrated, salt accumulation to the skin. Blood o Values - 7.4 o Acidic - This is associated with high blood pressure, kidney disease, diabetes and other conditions. o Basic ▪ Associated with adrenal disease & alcohol abuse. ▪ Dehydration ▪ Taking Diuretic drugs 8. What is the hydroxide ion concentration in a solution that has a pOH of 5.70? 9. A solution has a pOH of 11.76. What is the pH of this solution? 10. What is the pH of the solution if its pOH is 8.3? CALCULATION OF Ka AND Kb • • pH CALCULATION • • • • • • • • To calculate the pH of an aqueous solution you need to know the concentration of the hydronium ion in moles per liter (molarity). The pH is then calculated using the expression: o pH = - log [H3O+] To calculate the pOH of a solution you need to know the concentration of the hydroxide ion in moles per liter (molarity). The pOH is then calculated using the expression: o pOH = - log [OH-] The hydronium ion concentration can be found from the pH by the reverse of the mathematical operation employed to find the pH. o [H3O+] = 10-pH or [H3O+] = antilog (- pH) The hydroxide ion concentration can be found from the pOH by the reverse mathematical operation employed to find the pOH. o [OH-] = 10-pOH or [OH-] = antilog ( - pOH) The pH and pOH of a water solution at 25oC are related by the following equation. o pH + pOH = 14 Examples 1. Find the pH of a 0.0025 M HCl solution. 2. The H of HCl was 0.05M. What is the pH of the solution? 3. What is the pOH of a solution that has a hydroxide ion concentration of 4.82 x 10-5 M? 4. Find the pOH of a chemical solution if the hydroxide ion concentration is 4.22 x 10-5 M. 5. What is the OH of a solution if it has a pOH of 10.8? 6. What is the pH of the solution if its pOH is 9.5? Formulas o pH = - log [H3O+] o pOH = - log [OH-] o [H3O+] = 10-pH or [H3O+] = antilog (- pH) o [OH-] = 10-pOH or [OH-] = antilog ( - pOH) o pH + pOH = 14 Practice 1. Find the pH of a 0.0055 M HCl solution. 2. Find the pH of a 0.046 M H2SO4 solution. 3. The H of HCl was 0.09M. What is the pH of the solution? 4. What is the pOH of a solution that has a hydroxide ion concentration of 3.83 x 10-5 M? 5. Find the pOH of a chemical solution if the hydroxide ion concentration is 4.55 x 10-5 M. 6. What is the OH of a solution if it has a pOH of 11.6? 7. What is the hydronium ion concentration in a solution that has a pH of 8.34? • • • . The Ka for an acid is calculated from the pKa by performing the reverse of the mathematical operation used to find pKa. o Ka = 10-pKa or Ka = antilog ( - pKa) The Kb for an acid is calculated from the pKb by performing the reverse of the mathematical operation used to find pKb. o Kb = 10-pKb or Kb = antilog ( - pKb) Examples o What is the pKa of acetic acid, if Ka for acetic acid is 1.78 x 10-5? o What is the pKb for methyl amine, if the value of Kb for methyl amine is 4.4 x 10-4? o Calculate the value of the ionization constant for the ammonium ion, Ka, if the pKa is 9.74. o Calculate the value of the ionization constant, Kb, for aniline if the pKb is 9.38. Formulas o Ka = 10-pKa or Ka = antilog ( - pKa) o Kb = 10-pKb or Kb = antilog ( - pKb) Practice 1. What is the pKa of acetic acid, if Ka for acetic acid is 1.89 x 10-5? 2. What is the pKb for methyl amine, if the value of Kb for methyl amine is 5.4 x 10-6? 3. Calculate the value of the ionization constant for the ammonium ion, Ka, if the pKa is 8.34. 4. Calculate the value of the ionization constant, Kb, for aniline if the pKb is 8.78