HomeClass Notes1,200,000AU40,000

HBS1HBA Lecture Notes - Lecture 10: Carbonic Anhydrase, Carboxylic Acid, Proximal Tubule

49 views5 pages

9 Aug 2018

School

La Trobe University

Department

Course

Professor

For unlimited access to Class Notes, a Class+ subscription is required.

Document Summary

>acid-base balance is defined by the concentration of free hydrogen ions (h+) >there must be a balance between the production of hydrogen ions and the net removal of hydrogen ions from the body for homeostasis to be achieved. > a substance that releases h+ when in a solution. > a substance that can accept h+ when in solution. Plays crucial role in body"s ph buffering system. > solutions acidity or alkalinity is expressed on a ph scale. H+ concentration is restricted to between 10-1 10-14 m/l span is 100 million millionth 100,000,000,000,000. >due to large range ph value is expressed in a range of 1 14. >a ph 5 solution is 10 times more acidic than a ph 6 solution. >ph of ecf and icf maintained in narrow range. Acidosis = addition of excess h+ ph<7. 35 a low ph corresponds to a high h+ concentration.

Related Documents

Physiology 2130 Lecture Notes - Lecture 12: Carbonic Anhydrase, Carbonic Acid, Peripheral Chemoreceptors

coffeealligator974

NURS 83PL Lecture Notes - Lecture 5: Hypercapnia, Overproduction, Thiazide

NUR230 Chapter Notes - Chapter 8: Arterial Blood Gas, Pulse Oximetry, Respiratory Alkalosis

Related Questions

Please answer the questions 2-6 on the last 2 pictures by reading the first 5 pictures.

EXPERIMENT 4 ll the Preparation and Properties of Hydrogen t mou MATERIALS AND EQUIPMENT splints. Solids: strips of copper, magnesium, and zine; steel wool; mossy zine; wood Solutions dilute (6 M) acetic acid (HC,HO. 0.1 M coppertID sulfate (Cusdh dilute (6 M) hydrochlorie acid (HCD, dilute (3 AD phosphoric acid (H,PO),9 M sulfuri CHSO) and dilute (3 MD sulfuric acid (HSO), phenolphthalein solution. Pneumatie wide-mouth (gas-collecting) bottles, five glass cover plates, pÅH paper. DISCUSSION Hydrogen, having atomie number 1 and atomic mass 1.008 amu, is the simplest element. It is the ninth most abundant element in the earth's crust (about 0.9 percent by ma88). At ordi- nary temperatures and pressures it is a gas, composed of diatomic molecules, H, and is only very slightly soluble in water. Hydrogen is usually found combined with other elements Water is the most common and probably the most important compound of hydrogen. Hydrogen will not support combustion, but in the presence of oxygen it burns readily to form in air water: in the first l that This reaction is used as a simple test for hydrogen, for mixtures of hydrogen and air (or oxygen) burn explosively with a distinctive "popping" or "barking sound. A. Preparing Hydrogen There are several ways of producing hydrogen gas. The method used in this experiment involves the formation of hydrogen from aqueous solutions containing hydrogen ions (H'). The concentration of hydrogen ions in solution is often expressed as pH. pH is a numerical scale corresponding to the acidity of a solution. Pure water at 25 C is neither acidice nor basic and is given a pH of 7 on the pH scale. pH values less than 7 indicate an acidic solution; the lower the number, the more acidic the solution. pH values greater than 7 indicate a basic solu- tion; the larger the number the more basic the solution. ext ttle ht pH 7 = base solutions It is possible to determine whether a solution is acidic or basic by using acid-base indica- tors such as litmus or phenolphthalein. However, these indicators do not give very precise information. It is also possible to get more precise hydrogen ion concentrations using pH pape In this experiment you will measure the pEH of several acids and compare the relationship of their pH values to the strength of the acids and the rate of hydrogen formation. r strips impregnated with several indicators, or with an instrument called a pH meter -37-

Exp 25. pH Measurements- Buffers and Their Properties

I need help with finding kPa in the calculations part C for 1,2, and 3 Thanks!

Experiment 25 pH Measurements- Buffers and Their Properties ne of the more important propertics of an aqucous solution is its concentration of hydrogen ion. The H ion (or, more precisely, the HO ion) has a great effect on the solubility of many inorganic and organic species, on the nature of complex mctallic cations found in solutions, and on the rales of many chemical reac- tions. It is important that we know how to measure the concentration of hydrogen ion and understand its effect on solution properties. For convenience, the concentration of H' ion is frequendly expressed as the pH of the solution rather than as molarity. The pH of a solution is defined by the following equation: pH =-log where the logarithm is taken to the base 10. If [H] is I x 10-*moles per liter, the pH of the solution is, by the equation, 4. If the [H'1 is 5 x 10 M, the pH is 1.3, Basic solutions can also be described in terms of pH. In aqucous solutions the following equilibrium rela- tion will always be obeyed: In distilled water url equals [ OHÂ·], s, by Equation 2, [HÂ·l must be l x 10" M. Therefore, the pH of distilled water is 7. Solutions in which [H.] > 10é¨are sad to be acidic and will have a pH 7. A solution with a pH of 10 will have a [H"] of 1 x 10 0 M and a (OH] of I x 10M We measure the pH of a solution experimentally in two ways. In the first of these we use a chemical called an indicator, which is sensitive to pH. These substances have colors that change over a relatively short pH range (about two pH units) and can, when properly chosen, be used to delermine roughly the pH of a solu tion. Two very common indicators are litmus, usually used on paper, and phenolphthalein, the most common indicator in acid-base titrations. Litmus changes from red to blue as the pH of a solution goes from about 6 to about 8. Phenolphthalein changes from colorless to red as the pH goes from 8 to 10. A given indicator is useful for determining pH only in the region in which it changes color. Indicators are available for measure- ment of pH in all the important ranges of acidity and basicity. By matching the color of a suitable indicator in a solution of known pH with that in an unknown solution, one can determine the pH of the unknown to within about 0.3 pH units. The other method for finding pH is with a device called a pH meter. In this device two electrodes, one of which is sensitive to [H']. are immersed in a solution. The potential between the two electrodes is related to the pH. The pH meter is designed so that the scale will directly furnish the pH of the solution. A pH meter gives much more precise measurement of pH than does a typical indicator and is ordinarily used when an accurate determination of pH is needed Some acids and bases undergo substantial ionization in water, and are called strong because of their essentially complete ionization in reasonably dilute solutions. Other acids and bases, because of incomplete ionization (often far less than even 1% in 0.1 M solution), are called weak. Hydrochloric acid, HC, and sodium hydroxide, NaOH, are typical examples of a strong acid and a strong base. Acetic acid, HC,H,0, and ammonia, NH, are classic examples of a weak acid and a weak base. A weak acid will ionize according to the Law of Chemical Equilibrium: HBlaq)--H'(aq) + B-(aq)

Exp 25. pH Measurements-Buffers and Their Properties.

I need help in the calculations for part C finding the pKa in 1,2, and 3. Thanks!

Experiment 25 pH Measurements- Buffers and Their Properties ne of the is its concentration of hydrogen ion. The pecies, on they, the H,O ion) has a great effect on the solubility of many isorganic and ora tions. It is important H ion (or. more more important properties of an aqueous solution conplex metallic cations found in solutions, and on the rates of many chemical reac- of hydrogen ion and understand its effect on solution propertics that we know how to measure the concentration For convenience, the concentration of H" as molarity. The frequendly ex ãion is frequently expressed as the pH of the solution rather than pH of a solution is defined by the following equation: the equation, 4. If the [H to the base 10. If (H'l is 1 x 104 moles per liter, the pH of the solution is, by is 10-? M, the pH is 1.3. tion will always be obeyed: described in terms of pH. In aqueous solutions the following equilibrium rela- In distillecd water [H Soler t lequals IOH 1. so, by Equation 2, IH"] must be I X 10 M. Therefore,the pH of distilled will have a pH 7; if H 1 IOH1. the solution is basic and its pH >7. A solution with a pÅ of 10 will have or 1 Ã 10-4 M. We measure the pH of a solution experimentally in two ways. In the first of these we use a chemical an indicator, which is sensitive to pH. These substances have colors that change over a relatively short two pH units) and can, when properly chosen, be used to determine roughly the pH of a solu- called H range on. Two very common indicators are litmus, usually used on paper, and phenolphthalein, the most common indicator in acid-base titrations. Litmus changes from red to blue as the pH of a solution goes from about 6 to about 8. Phenolphthalein changes from colorless to red as the pH goes from 8 to 10. A given indicator is useful for determining pH only in the region in which it changes color. Indicators are available for measure- ment of pH in all the important ranges of acidity and basicity. By matching the color of a suitable indicator in a solution of known pH with that in an unknown solution, one can determine the pH of the unknown to within about 0.3 pH units. The other method for finding pH is with a device called a pH meter. In this device two electrodes, one of which is sensitive to [H'], are immersed in a solution. The potential between the two electrodes is related to the pH. The pH meter is designed so that the scale will directly furnish the pH of the solution. A pH meter gives much more precise measurement of pH than does a typical indicator and is ordinarily used when an accurate determination of pH is needed Some acids and bases undergo substantial ionization in water, and are called strong because of their essentially complete ionization in reasonably dilute solutions. Other acids and bases, because of incomplete ionization (often far less than even 1% in 0.1 M solution), are called weak Hydrochloric acid, HCl, and sodium hydroxide, NaOH, are typical examples of a strong acid and a strong base. Acetic acid, HCH,0, and ammonia, NH, are classic examples of a weak acid and a weak base. A weak acid will ionize according to the Law of Chemical Equilibrium HB(aq)H'(a) B (aq) 3)