6.0 Conceitos sobre pH e soluções tampão

7/30/2019 6.0 Conceitos sobre pH e solues tampo

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CONCEITOS SOBRE pH E SOLUES TAMPO

A buffer solution is an aqueous solution consisting of a mixture of aweak acid and its conjugate base or a weak base and its conjugate acid. It has the property that thepH of the solution changesvery little when a small amount of strong acid or base is added to it. Buffer solutions are used as ameans of keeping pH at a nearly constant value in a wide variety of chemical applications. Manylife forms thrive only in a relatively small pH range; an example of a buffer solution isblood.

Le Chtelier's Principle states that " if a change in conditions ([an external] 'stress') is imposed on a system at equilibrium, the equilibrium position will shift in a direction thattends to reduce that change in conditions "

Principles of Buffering

Buffer solutions achieve their resistance to pH change because of the presence of an equilibriumbetween the acid HA and its conjugate base A -.

HA H+ + A-

When some strong acid is added to an equilibrium mixture of theweak acid and its conjugate base , the equilibrium is shifted to the left, in accordance withLe Chatelier's principle. Becauseof this, the hydrogen ion concentration increases by less than the amount expected for thequantity of strong acid added. Similarly, if strong alkali is added to the mixture the hydrogenion concentration decreases by less than the amount expected for the quantity of alkali added.

The effect is illustrated by the simulated titration of a weak acid with pKa = 4.7. The relativeconcentration of undissociated acid is shown in blue and of its conjugate base in red. The pHchanges relatively slowly in the buffer region, pH = pKa 1, centered at pH = 4.7 where [HA] =

[A-], but once the acid is more than 95% deprotonated the pH rises much more rapidly.

Applications

Buffer solutions are necessary to keep the correct pH for enzymes in many organisms to work.Many enzymes work only under very precise conditions; if the pH moves outside of a narrowrange, the enzymes slow or stop working and can denature, thus permanently disabling their catalytic activity.[2] A buffer of carbonic acid (H2CO3) and bicarbonate (HCO3) is present in blood

plasma, to maintain a pH between 7.35 and 7.45.Industrially, buffer solutions are used infermentation processes and in setting the correctconditions for dyes used in colouring fabrics. They are also used in chemical analysis[1] andcalibration of pH meters.

The majority of biological samples that are used in research are made in buffers,especially phosphate buffered saline (PBS) at pH 7.4.
http://en.wikipedia.org/wiki/Aqueous_solutionhttp://en.wikipedia.org/wiki/Weak_acidhttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Weak_basehttp://en.wikipedia.org/wiki/Conjugate_acidhttp://en.wikipedia.org/wiki/PHhttp://en.wikipedia.org/wiki/Strong_acidhttp://en.wikipedia.org/wiki/Base_(chemistry)#Strong_baseshttp://en.wikipedia.org/wiki/Bloodhttp://en.wikipedia.org/wiki/Strong_acidhttp://en.wikipedia.org/wiki/Weak_acidhttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Le_Chatelier's_principlehttp://en.wikipedia.org/wiki/Enzymehttp://en.wikipedia.org/wiki/Denaturation_(biochemistry)http://en.wikipedia.org/wiki/Buffer_solution#cite_note-Scorpio_2000-1http://en.wikipedia.org/wiki/Buffer_solution#cite_note-Scorpio_2000-1http://en.wikipedia.org/wiki/Carbonic_acidhttp://en.wikipedia.org/wiki/Bicarbonatehttp://en.wikipedia.org/wiki/Blood_plasmahttp://en.wikipedia.org/wiki/Blood_plasmahttp://en.wikipedia.org/wiki/Fermentation_(biochemistry)http://en.wikipedia.org/wiki/Buffer_solution#cite_note-Hulanicki-0http://en.wikipedia.org/wiki/Phosphate_buffered_salinehttp://en.wikipedia.org/wiki/Aqueous_solutionhttp://en.wikipedia.org/wiki/Weak_acidhttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Weak_basehttp://en.wikipedia.org/wiki/Conjugate_acidhttp://en.wikipedia.org/wiki/PHhttp://en.wikipedia.org/wiki/Strong_acidhttp://en.wikipedia.org/wiki/Base_(chemistry)#Strong_baseshttp://en.wikipedia.org/wiki/Bloodhttp://en.wikipedia.org/wiki/Strong_acidhttp://en.wikipedia.org/wiki/Weak_acidhttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Conjugate_basehttp://en.wikipedia.org/wiki/Le_Chatelier's_principlehttp://en.wikipedia.org/wiki/Enzymehttp://en.wikipedia.org/wiki/Denaturation_(biochemistry)http://en.wikipedia.org/wiki/Buffer_solution#cite_note-Scorpio_2000-1http://en.wikipedia.org/wiki/Carbonic_acidhttp://en.wikipedia.org/wiki/Bicarbonatehttp://en.wikipedia.org/wiki/Blood_plasmahttp://en.wikipedia.org/wiki/Blood_plasmahttp://en.wikipedia.org/wiki/Fermentation_(biochemistry)http://en.wikipedia.org/wiki/Buffer_solution#cite_note-Hulanicki-0http://en.wikipedia.org/wiki/Phosphate_buffered_saline


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Example : Removal of Bicarbonate by the Kidneys

Le Chtelier's Principle can be used to explain how the kidneys help prevent excessively high pH(a condition known as alkalosis ). When the pH of the blood is too high, the kidneys removebicarbonate ion (HCO3-) from the blood. As the kidneys decrease the blood's concentration of HCO3-, the equilibrium reaction in Equation 10 is shifted toward the left to compensate for the loss

in HCO3-, according to Le Chtelier's Principle. When the equilibrium shifts to the left, morH+ ions are generated together with HCO3- ions. As a result, the pH decreases.

Useful buffer mixtures

Components pH range

HCl, Sodium citrate 1 - 5

Citric acid, Sodium citrate 2.5 - 5.6

Acetic acid, Sodium acetate 3.7 - 5.6

K2HPO4, KH2PO4 5.8 - 8

Na2HPO4, NaH2PO4 6 - 7.5

Borax, Sodium hydroxide 9.2 - 11

Acidic buffer solutions

An acidic buffer solution is simply one which has a pH less than 7. Acidic buffer solutions arecommonly made from a weak acid and one of its salts - often a sodium salt.

A common example would be a mixture of ethanoic acid and sodium ethanoate in solution. In thiscase, if the solution contained equal molar concentrations of both the acid and the salt, it wouldhave a pH of 4.76. It wouldn't matter what the concentrations were, as long as they were the same.
http://en.wikipedia.org/wiki/Hydrochloric_acidhttp://en.wikipedia.org/wiki/Sodium_citratehttp://en.wikipedia.org/wiki/Citric_acidhttp://en.wikipedia.org/wiki/Sodium_citratehttp://en.wikipedia.org/wiki/Acetic_acidhttp://en.wikipedia.org/wiki/Sodium_acetatehttp://en.wikipedia.org/wiki/Sodium_acetatehttp://en.wikipedia.org/wiki/Dipotassium_phosphatehttp://en.wikipedia.org/wiki/Dipotassium_phosphatehttp://en.wikipedia.org/wiki/Dipotassium_phosphatehttp://en.wikipedia.org/wiki/Dipotassium_phosphatehttp://en.wikipedia.org/wiki/Monopotassium_phosphatehttp://en.wikipedia.org/wiki/Monopotassium_phosphatehttp://en.wikipedia.org/wiki/Monopotassium_phosphatehttp://en.wikipedia.org/wiki/Monopotassium_phosphatehttp://en.wikipedia.org/wiki/Disodium_phosphatehttp://en.wikipedia.org/wiki/Disodium_phosphatehttp://en.wikipedia.org/wiki/Disodium_phosphatehttp://en.wikipedia.org/wiki/Disodium_phosphatehttp://en.wikipedia.org/wiki/Disodium_phosphatehttp://en.wikipedia.org/wiki/Monosodium_phosphatehttp://en.wikipedia.org/wiki/Monosodium_phosphatehttp://en.wikipedia.org/wiki/Monosodium_phosphatehttp://en.wikipedia.org/wiki/Monosodium_phosphatehttp://en.wikipedia.org/wiki/Boraxhttp://en.wikipedia.org/wiki/Sodium_hydroxidehttp://en.wikipedia.org/wiki/Hydrochloric_acidhttp://en.wikipedia.org/wiki/Sodium_citratehttp://en.wikipedia.org/wiki/Citric_acidhttp://en.wikipedia.org/wiki/Sodium_citratehttp://en.wikipedia.org/wiki/Acetic_acidhttp://en.wikipedia.org/wiki/Sodium_acetatehttp://en.wikipedia.org/wiki/Sodium_acetatehttp://en.wikipedia.org/wiki/Dipotassium_phosphatehttp://en.wikipedia.org/wiki/Monopotassium_phosphatehttp://en.wikipedia.org/wiki/Disodium_phosphatehttp://en.wikipedia.org/wiki/Monosodium_phosphatehttp://en.wikipedia.org/wiki/Boraxhttp://en.wikipedia.org/wiki/Sodium_hydroxide


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Ethanoic acid is a weak acid, and the position of this equilibrium will be well to the left:

Alkaline buffer solutions

An alkaline buffer solution has a pH greater than 7. Alkaline buffer solutions are commonly madefrom a weak base and one of its salts.

A frequently used example is a mixture of ammonia solution and ammonium chloride solution. Ithese were mixed in equal molar proportions, the solution would have a pH of 9.25. Again, itdoesn't matter what concentrations you choose as long as they are the same.

Adding an acid to this buffer solution

The buffer solution must remove most of the new hydrogen ions otherwise the pH would dropmarkedly.

Hydrogen ions combine with the ethanoate ions to make ethanoic acid. Although the reaction isreversible, since the ethanoic acid is a weak acid, most of the new hydrogen ions are removed inthis way.

Since most of the new hydrogen ions are removed, the pH won't change very much - but becauseof the equilibria involved, it will fall a little bit.

Adding an alkali to this buffer solution

Alkaline solutions contain hydroxide ions and the buffer solution removes most of these.

This time the situation is a bit more complicated because there are two processes which canremove hydroxide ions.

Removal by reacting with ethanoic acid

The most likely acidic substance which a hydroxide ion is going to collide with is an ethanoic acidmolecule. They will react to form ethanoate ions and water.

Because most of the new hydroxide ions are removed, the pH doesn't increase very much.

How Buffers Work: A Quantitative View

The kidneys and the lungs work together to help maintain a blood pH of 7.4 by affecting thecomponents of the buffers in the blood. Therefore, to understand how these organs help control thepH of the blood, we must first discuss how buffers work in solution.


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Acid-base buffers confer resistance to a change in the pH of a solution when hydrogen ions(protons) or hydroxide ions are added or removed. An acid-base buffer typically consists of aweakacid , and its conjugate base (salt). Buffers work because the concentrations of the weak acid andits salt are large compared to the amount of protons or hydroxide ions added or removed. Whenprotons are added to the solution from an external source, some of the base component of thebuffer is converted to the weak-acid component (thus using up most of the protons added); whenhydroxide ions are added to the solution (or, equivalently, protons are removed from the solution;

see Equations 8-9 in the blue box, below), protons are dissociated from some of the weak-acidmolecules of the buffer, converting them to the base of the buffer (and thus replenishing most of the protons removed). However, the change in acid and base concentrations is small relative to theamounts of these species present in solution. Hence, the ratio of acid to base changes onlyslightly. Thus, the effect on the pH of the solution is small, within certain limitations on the amounof H+ or OH- added or removed.

The Carbonic-Acid-Bicarbonate Buffer in the Blood

By far the most important buffer for maintaining acid-base balance in the blood is the carbonic-acid-bicarbonate buffer. The simultaneous equilibrium reactions of interest are

.

We are interested in the change in the pH of the blood; therefore, we want an expression for theconcentration of H+ in terms of an equilibrium constant (see blue box, below) and theconcentrations of the other species in the reaction (HCO3-, H2CO3, and CO2).

o more clearly show the two equilibrium reactions in the carbonic-acid-bicarbonate buffer, Equation1 is rewritten to show the direct involvement of water:

The equilibrium on the left is an acid-base reaction that is written in the reverse format fromEquation 3. Carbonic acid (H2CO3) is the acid and water is the base. The conjugate base for H2CO3 is HCO3- (bicarbonate ion). Carbonic acid also dissociates rapidly to produce water andcarbon dioxide, as shown in the equilibrium on the right of Equation 10. This second process is notan acid-base reaction, but it is important to the blood's buffering capacity, as we can see fromEquation 11, below.

.

The derivation for this equation is shown in the yellow box, below. Notice that Equation 11 is in asimilar form to theHenderson-Hasselbach equation presented in the introduction to theExperiment (Equation 16 in the lab manual). Equation 11 does not meet the strict definition of a


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Henderson-Hasselbach equation, because this equation takes into account a non-acid-basereaction (i.e., the dissociation of carbonic acid to carbon dioxide and water), and the ratio inparentheses is not the concentration ratio of the acid to the conjugate base. However, therelationship shown in Equation 11 is frequently referred to as the Henderson-Hasselbach equationfor the buffer in physiological applications.

In Equation 11, pK is equal to the negative log of the equilibrium constant, K, for the buffer

(Equation 12).

where K=Ka/K2 (from Equation 10). (12

This quantity provides an indication of the degree to which HCO3- reacts with H+ (or with H3O+ aswritten in Equation 10) to form H2CO3, and subsequently to form CO2 and H2O. In the case of thecarbonic-acid-bicarbonate buffer, pK=6.1 at normal body temperature.

As shown in Equation 11, the pH of the buffered solution (i.e., the blood) is dependent only on

the ratio of the amount of CO2 present in the blood to the amount of HCO3-

(bicarbonate ion)present in the blood (at a given temperature, so that pK remains constant ). This ratio remainsrelatively constant, because the concentrations of both buffer components (HCO3- and CO2) arevery large, compared to the amount of H+ added to the blood during normal activities and moderateexercise. When H+ is added to the blood as a result of metabolic processes, the amount of HCO3- (relative to the amount of CO2) decreases; however, the amount of the change is tinycompared to the amount of HCO3- present in the blood. This optimal buffering occurs when the pHis within approximately 1 pH unit from the pK value for the buffering system,i.e., when the pH isbetween 5.1 and 7.1.

However, the normal blood pH of 7.4 is outside the optimal buffering range; therefore, the addition

of protons to the blood due to strenuous exercise may be too great for the buffer alone toeffectively control the pH of the blood. When this happens, other organs must help control theamounts of CO2 and HCO3- in the blood. The lungs remove excess CO2 from the blood (helping toraise the pH via shifts in the equilibria in Equation 10), and the kidneys remove excess HCO3- fromthe body (helping to lower the pH). The lungs' removal of CO2 from the blood is somewhat impededduring exercise when the heart rate is very rapid; the blood is pumped through the capillaries veryquickly, and so there is little time in the lungs for carbon dioxide to be exchanged for oxygen. Theways in which these three organs help to control the blood pH through the bicarbonate buffer system are highlighted in Figure 3, below.


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Figure 3: This figure shows the major organsthat help control the bloodconcentrations of CO2 and HCO3-, andthus help control the pH of the blood.

Removing CO2 from the blood helps increasethe pH.Removing HCO3- from the blood helps lower thepH.

Other pH-Buffer Systems in the Blood

Other buffers perform a more minor role than the carbonic-acid-bicarbonate buffer in regulating thepH of the blood. The phosphate buffer consists of phosphoric acid (H3PO4) in equilibrium withdihydrogen phosphate ion (H2PO4-) and H+. The pK for the phosphate buffer is 6.8, which allowsthis buffer to function within its optimal buffering range at physiological pH. The phosphate buffeonly plays a minor role in the blood, however, because H3PO4 and H2PO4- are found in very lowconcentration in the blood. Hemoglobin also acts as a pH buffer in the blood. hemoglobin proteincan reversibly bind either H+ (to the protein) or O2 (to the Fe of the heme group), but that when oneof these substances is bound, the other is released (as explained by the Bohr effect). Duringexercise, hemoglobin helps to control the pH of the blood by binding some of the excess protonsthat are generated in the muscles. At the same time, molecular oxygen is released for use by themuscles.

EXERCCIO E TAMPONAMENTO DE PRODUTOS DA RESPIRAO


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The following steps outline the processes that affect the buffers in the blood during exercise.

Hemoglobincarries O2 from the lungs to the muscles through the blood.

The muscles need more O 2 than normal, because their metabolic activity is increased duringexercise. The amount of oxygen in the muscle is therefore depleted in the muscles, setting up aconcentration gradient between the muscle cells and the blood in the capillaries. Oxygen diffusesfrom the blood to the muscles, via this concentration gradient.

The muscles produce CO 2 and H+ as a result of increased metabolism, setting up concentrationgradients in the opposite direction from the O2 gradient.

The CO2 and H+ flow from the muscles to the blood, via these concentration gradients.

The buffering action of hemoglobin picks up the extra H+ and CO2.

If the amounts of H+ and CO2 exceed the capacity of hemoglobin, they affect the carbonic acidequilibrium (Equation 10), as predicted by Le Chtelier's Principle or the quantitative treatment interms of equilibrium constants. As a result, the pH of the blood is lowered, causing acidosis.

The lungs and kidneys respond to pH changes by removing CO2, HCO3-, and H+ from the blood.

Hence, the body has developed finely-tuned chemical processes (based on buffering and acid-base equilibria) that work in combination to handle the changes that exercise produces.
http://www.chemistry.wustl.edu/~edudev/LabTutorials/Hemoglobin/MetalComplexinBlood.htmlhttp://www.chemistry.wustl.edu/~edudev/LabTutorials/Dialysis/Kidneys.htmlhttp://www.chemistry.wustl.edu/~edudev/LabTutorials/Hemoglobin/MetalComplexinBlood.htmlhttp://www.chemistry.wustl.edu/~edudev/LabTutorials/Dialysis/Kidneys.html


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ACIDOSE E ALCALOSE METABLICA

O pH dos fluidos corporais deve permanecer dentro de limites prximos. Morte poderesultar de aumentos ou diminuies relativamente pequenos de pH. A razo para nossasensibilidade ao pH depende das enzimas que catalisam as reaes qumicas do corpo. Sua

atividade cai rapidamente quando o pH muda. Uma vez que as enzimas desempenham papel tocrucial, sua inativao pode ser fatal.

Os tampes do plasma sangneo so as primeiras defesas do corpo contra mudanas dopH interno. Seu papel manter o pH sangneo dentro dos limites 7,35 a 7,45. Se o pH do sanguede uma pessoa cai abaixo de 7,35, diz-se que ela est com acidose, ou baixo pH sangneo.

A acidose causa desorientao, coma e finalmente morte. Se o pH sobe alm de 7,45, diz-se que ela est com alcalose, ou alto pH sangneo.

A alcalose pode causar respirao fraca e irregular, cimbras musculares e convulses. Aacidose e alcalose podem ser "respiratrias", resultando de mudanas na concentrao de H 2CO3,ou "metablicas", resultando de mudanas na concentrao de HCO 3-.

Acidose metablica A acidose metablica a acidez excessiva do sangue caracterizada pela concentrao patologicamentebaixa de ons bicarbonato (HCO3) no sangue. Quando um aumento do cido supera o sistema tampo dopH do corpo, o sangue pode tornar-se realmente cido (pH menor que 7) e provocar um srio desequilbrio

homeosttico. A acidose respiratria ocorre s vezes com pessoas com pneumonia, enfisema, poliomielite, ou empacientes anestesiados, porque estas condies podem interferir na respirao.

Seu nvel de H2CO3 sobe porque CO 2 suficiente no exalado. O sistema nervoso responde acidosetentando aumentar a velocidade e a intensidade da respirao a fim de diminuir a CO 2.

A alcalose respiratria resulta do outro extremo-hiperventilao, ou respirao aumentada. Neste caso,muito CO2 perdido pelos pulmes. A pCO2 cai e a concentrao de H 2CO3 reduzida abaixo do valor normal.

Um ataque de ansiedade ou histeria pode s vezes fazer com que uma pessoa respire muito rapidamente.Nestas condies, o pH sangneo pode subir a 7,6 ou 7,7 dentro de minutos. O sistema nervoso responde alcalose baixando a velocidade de respirao para aumentar a pCO 2 e a concentrao de H 2CO3.

Como mencionado, a acidose e a alcalose podem resultar de mudanas na concentrao de HCO 3-. Acidose metablica, ou baixa concentrao de HCO 3-, pode resultar de desidratao severa ou de diabetes


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mellitus no tratada. Alcalose metablica, ou alta concentrao de HCO3-, pode resultar de vmito ou"overdose" de anticidos.

Para diagnosticar a condio cido-base de um paciente, uma amostra de sangue arterial retirada.Medidas de pH, pCO2 e CO2(H2CO3) total so ento realizadas.

Mecanismos de compensao:

Respirao mais profunda e rpida: o organismo tenta livrar o sangue do excesso de cidoreduzindo a quantidade de dixido de carbono.

Os rins tentam excretar mais cido atravs da urina.

Quando estes dois mecanismos no conseguem estabelecer a homeostase e o corpo continua a produzir cido em demasia, instala-se um quadro de acidose grave e, em ltima instncia, o coma.

Causas

Podem ser agrupadas em trs categorias principais:

1 - Aumento da quantidade de cido no organismo pela ingesto de um cido ou de uma substncia que,ao ser metabolizada, transforma-se em um cido. lcool de madeira (metanol) e os anticongelantes(etileno glicol). Doses muito elevadas de aspirina (cido acetilsaliclico), paraldedo, acetazolamida oucloreto de amnio

2 - Produo de quantidades maiores de cido atravs do metabolismo (aumento da produo endgena,reduo da excreo) Insuficincia renal Acidose tubular renal (uma forma de malformao renal) Cetoacidose diabtica Acidose ltica (acmulo de cido ltico)

3 - Perda de base (bicarbonato). Perda gastrintestinal (fstula digestiva, diarria), perda urinria (acidosetubular

CAUSAS TXICAS

A acidose metablica pode ser a principal manifestao presente nas intoxicaes poretilenoglicol, metanol ou salicilato.

Outros agentes associados com acidose metablica so:

cidos Cloreto de amnio Ibuprofeno Formaldedo Paraldedo


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cido valprico Agentes que causam acidose lctica

Acetaminofeno.Biguanidas (metformina, fenformina).Monxido de Carbono.Cloranfenicol.Cocana.Cianeto.Etanol.Frutose.Sulfeto de Hidrognio.Ferro.Isoniazida.cido nalidxico.Nitroprussiato de Sdio.Estricnina.

A cetoacidose um tipo de acidose metablica que causada por altas concentraesde cetocidos, formados pela desaminao dos aminocidos. Ela mais comum na diabetes mellitus tipo1 no tratada, quando o fgado quebra a gordura e protenas em resposta a uma necessidade percebida desubstrato respiratrio. Ela tambm pode ocorrer em pessoas que passam muita fome, por mais de trsdias, ou pessoas quase morrendo de fome, j que o corpo forado a quebrar a gordura para se manter,devido falta de nutrio externa.

A cetoacidose no deve ser confundida com a cetose , que um dos processos normais do corpopara o metabolismo da gordura corporal. Na cetoacidose o acmulo de cetocidos to severo queo pH do sangue substancialmente reduzido.

uma decorrncia de uma deficincia insulnica grave e de um estado de resistncia a insulina, e temcomo critrios clnicos:

Glicemia maior que 300 mg/dl Acidose metablica = pH < 7,3 e BIC < 15 mEq/l.

Os: Cetocidos so cidos orgnicos que contm um grupo funcional cetona e um grupo cido

carboxlico.
http://pt.wikipedia.org/wiki/Acidosehttp://pt.wikipedia.org/wiki/Ceto%C3%A1cidohttp://pt.wikipedia.org/wiki/Amino%C3%A1cidohttp://pt.wikipedia.org/wiki/Diabetes_mellitus_tipo_1http://pt.wikipedia.org/wiki/Diabetes_mellitus_tipo_1http://pt.wikipedia.org/wiki/F%C3%ADgadohttp://pt.wikipedia.org/wiki/Gordurahttp://pt.wikipedia.org/wiki/Prote%C3%ADnahttp://pt.wikipedia.org/wiki/Cetosehttp://pt.wikipedia.org/wiki/PHhttp://pt.wikipedia.org/wiki/Sanguehttp://pt.wikipedia.org/wiki/Acidose_metab%C3%B3licahttp://pt.wikipedia.org/wiki/PHhttp://pt.wikipedia.org/wiki/%C3%81cidohttp://pt.wikipedia.org/wiki/Composto_org%C3%A2nicohttp://pt.wikipedia.org/wiki/Grupo_funcionalhttp://pt.wikipedia.org/wiki/Cetonahttp://pt.wikipedia.org/wiki/%C3%81cido_carbox%C3%ADlicohttp://pt.wikipedia.org/wiki/%C3%81cido_carbox%C3%ADlicohttp://pt.wikipedia.org/wiki/Acidosehttp://pt.wikipedia.org/wiki/Ceto%C3%A1cidohttp://pt.wikipedia.org/wiki/Amino%C3%A1cidohttp://pt.wikipedia.org/wiki/Diabetes_mellitus_tipo_1http://pt.wikipedia.org/wiki/Diabetes_mellitus_tipo_1http://pt.wikipedia.org/wiki/F%C3%ADgadohttp://pt.wikipedia.org/wiki/Gordurahttp://pt.wikipedia.org/wiki/Prote%C3%ADnahttp://pt.wikipedia.org/wiki/Cetosehttp://pt.wikipedia.org/wiki/PHhttp://pt.wikipedia.org/wiki/Sanguehttp://pt.wikipedia.org/wiki/Acidose_metab%C3%B3licahttp://pt.wikipedia.org/wiki/PHhttp://pt.wikipedia.org/wiki/%C3%81cidohttp://pt.wikipedia.org/wiki/Composto_org%C3%A2nicohttp://pt.wikipedia.org/wiki/Grupo_funcionalhttp://pt.wikipedia.org/wiki/Cetonahttp://pt.wikipedia.org/wiki/%C3%81cido_carbox%C3%ADlicohttp://pt.wikipedia.org/wiki/%C3%81cido_carbox%C3%ADlico


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MANIFESTAES CLNICAS

Na acidose metablica leve, as manifestaes clnicas so aquelas decorrentes daprpria intoxicao.

Nos casos de acidose mais grave (pH < 7.2, bicarbonato < 13 mmol/L), independenteda causa de base, podem ser produzidos efeitos diretos cardiovasculares,respiratrios, gastrointestinais e em SNC. A contratilidade do miocrdio afetada epode progredir para choque circulatrio. A respirao se torna anormal, mais profundae ento mais frequente (respirao de Kussmaul). A depresso de SNC evolui para ocoma. Dor abdominal e nusea podem estar presentes. Hipercalemia umacomplicao da acidose, que resulta em potencial risco de vida.

ALCALOSE

A alcalose um termo clnico que indica umtranstorno no qual h um aumentona alcalinidade dos fluidos do corpo, isto , um excesso de base (alcalina) nos lquidos corporais.Esta condio oposta a produzida por excesso de cidos ( acidose). A alcalose pode se originar por diferentes causas.

O mecanismo subjacente consiste na acumulao de bases ou perda de cidos sem umaperda equivalente de bases nos lquidos do organismo, o que provoca uma reduo naconcentrao de ons hidrognio no plasma sanguneo arterial. Geralmente utiliza-se este termonos casos em que o pH arterial maior que 7,45.

Sendo os pulmes e os rins os que regulam o estado cido/bsico do corpo, a diminuio no nvelde dixido de carbono ou o aumento do nvel de bicarbonato so as causas diretas destefenmeno.
http://pt.wikipedia.org/wiki/Transtornohttp://pt.wikipedia.org/wiki/Alcalinidadehttp://pt.wikipedia.org/wiki/Base_(qu%C3%ADmica)http://pt.wikipedia.org/wiki/Acidosehttp://pt.wikipedia.org/wiki/%C3%81cidohttp://pt.wikipedia.org/wiki/Hidr%C3%B4niohttp://pt.wikipedia.org/wiki/Plasma_sangu%C3%ADneohttp://pt.wikipedia.org/wiki/Art%C3%A9riahttp://pt.wikipedia.org/wiki/PHhttp://pt.wikipedia.org/wiki/Pulm%C3%B5eshttp://pt.wikipedia.org/wiki/Rinshttp://pt.wikipedia.org/wiki/Equil%C3%ADbrio_qu%C3%ADmicohttp://pt.wikipedia.org/wiki/Bicarbonatohttp://pt.wikipedia.org/wiki/Transtornohttp://pt.wikipedia.org/wiki/Alcalinidadehttp://pt.wikipedia.org/wiki/Base_(qu%C3%ADmica)http://pt.wikipedia.org/wiki/Acidosehttp://pt.wikipedia.org/wiki/%C3%81cidohttp://pt.wikipedia.org/wiki/Hidr%C3%B4niohttp://pt.wikipedia.org/wiki/Plasma_sangu%C3%ADneohttp://pt.wikipedia.org/wiki/Art%C3%A9riahttp://pt.wikipedia.org/wiki/PHhttp://pt.wikipedia.org/wiki/Pulm%C3%B5eshttp://pt.wikipedia.org/wiki/Rinshttp://pt.wikipedia.org/wiki/Equil%C3%ADbrio_qu%C3%ADmicohttp://pt.wikipedia.org/wiki/Bicarbonato


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Alcalose respiratria

A alcalose respiratria a ocasionada por nveis baixos de dixido de carbono(CO2). A hiperventilao (frequncia respiratria aumentada) faz com que o corpo perca dixido decarbono. A altitude e, em geral, qualquer enfermidade que produza uma reduode oxigniono sangue obrigam o indivduo a respirar mais rapidamente (como nos quadros de

ansiedade), minguando os nveis de dixido de carbono e ocasionando este tipo de alcalose.

Alcalose metablica

A alcalose metablica, por sua vez, ocasionada por um excesso de bicarbonato no sangue.

A alcalose hipoclormica aquela causada por uma deficincia ou perda extremade cloreto (que pode ser devido a vmitos persistentes). Nesses casos, os rins compensam aperda de cloretos mediante a conservao de bicarbonato.

A alcalose hipocalmica se deve a reao dos rins a uma deficincia ou perda extremade potssio que pode ser provocada pelo uso de alguns medicamentos diurticos.

A alcalose compensada se apresenta quando o corpo tem uma compensao parcial daalcalose, alcanando o equilbrio normal cido/bsico, ou quando os nveis de bicarbonato edixido de carbono permanecem anormais em termos absolutos.

CAUSAS TXICAS:

Bicarbonato Diurtico de Ala (furosemida, cido etacrnico)

Diurticos mercuriais (agora obsoletos)Qualquer intoxicao que resulta em vmitos excessivos pode causar alcalose metablicasecundria.

MANIFESTAES CLNICAS

Histria recente de perda excessiva do contedo gstrico, administrao de altas doses dediurtico de ala ou sobrecarga de lcali em pacientes com falncia renal.

Irritabilidade, hiperexcitabilidade, confuso mental, s vezes semelhante a intoxicao alcolica,bradipnia (mesmo abaixo de 6 a 8 respiraes por minuto), cianose s vezes extrema.

Fraqueza muscular, reduo do peristaltismo gastrointestinal (reteno gstrica, leo paraltico) epoliria, sugerem depleo associada de K+. Tetania pode ocorrer devido a diminuio de clcioionizado no soro.
http://pt.wikipedia.org/wiki/Alcalose_respirat%C3%B3riahttp://pt.wikipedia.org/wiki/Di%C3%B3xido_de_carbonohttp://pt.wikipedia.org/wiki/Hiperventila%C3%A7%C3%A3ohttp://pt.wikipedia.org/wiki/Frequ%C3%AAncia_respirat%C3%B3riahttp://pt.wikipedia.org/wiki/Altitudehttp://pt.wikipedia.org/wiki/Oxig%C3%AAniohttp://pt.wikipedia.org/wiki/Sanguehttp://pt.wikipedia.org/wiki/Alcalose_metab%C3%B3licahttp://pt.wikipedia.org/wiki/Hipocloremiahttp://pt.wikipedia.org/wiki/Cloretohttp://pt.wikipedia.org/wiki/V%C3%B3mitohttp://pt.wikipedia.org/wiki/Hipocalemiahttp://pt.wikipedia.org/wiki/Pot%C3%A1ssiohttp://pt.wikipedia.org/wiki/Medicamentohttp://pt.wikipedia.org/wiki/Diur%C3%A9ticohttp://pt.wikipedia.org/wiki/Alcalose_respirat%C3%B3riahttp://pt.wikipedia.org/wiki/Di%C3%B3xido_de_carbonohttp://pt.wikipedia.org/wiki/Hiperventila%C3%A7%C3%A3ohttp://pt.wikipedia.org/wiki/Frequ%C3%AAncia_respirat%C3%B3riahttp://pt.wikipedia.org/wiki/Altitudehttp://pt.wikipedia.org/wiki/Oxig%C3%AAniohttp://pt.wikipedia.org/wiki/Sanguehttp://pt.wikipedia.org/wiki/Alcalose_metab%C3%B3licahttp://pt.wikipedia.org/wiki/Hipocloremiahttp://pt.wikipedia.org/wiki/Cloretohttp://pt.wikipedia.org/wiki/V%C3%B3mitohttp://pt.wikipedia.org/wiki/Hipocalemiahttp://pt.wikipedia.org/wiki/Pot%C3%A1ssiohttp://pt.wikipedia.org/wiki/Medicamentohttp://pt.wikipedia.org/wiki/Diur%C3%A9tico


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O indivduo com alcalose pode apresentar quadros de confuso, com enjos, nuseas e vmitos,a muitas vezes acompanhados de tremores, espasmos musculares e inchao no rosto ou nasextremidades.
http://pt.wikipedia.org/wiki/Confus%C3%A3ohttp://pt.wikipedia.org/wiki/Enjoohttp://pt.wikipedia.org/wiki/N%C3%A1useahttp://pt.wikipedia.org/wiki/V%C3%B4mitohttp://pt.wikipedia.org/wiki/Tremorhttp://pt.wikipedia.org/wiki/Espasmo_muscularhttp://pt.wikipedia.org/wiki/Incha%C3%A7ohttp://pt.wikipedia.org/w/index.php?title=Extremidades&action=edit&redlink=1http://pt.wikipedia.org/wiki/Confus%C3%A3ohttp://pt.wikipedia.org/wiki/Enjoohttp://pt.wikipedia.org/wiki/N%C3%A1useahttp://pt.wikipedia.org/wiki/V%C3%B4mitohttp://pt.wikipedia.org/wiki/Tremorhttp://pt.wikipedia.org/wiki/Espasmo_muscularhttp://pt.wikipedia.org/wiki/Incha%C3%A7ohttp://pt.wikipedia.org/w/index.php?title=Extremidades&action=edit&redlink=1

6.0 Conceitos sobre pH e soluções tampão

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