Name: buffer capacity and ph stability
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 19 s
Channel: JoVE

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Ohio’s Learning Standards and Model Curriculum

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Chemistry

Interactions Of Matter

Chemical reactions

Buffer Effectiveness

pH 변화를 저애하는 완충제의 효과는 약산과 그 짝 염기 또는 약염기 및 그 짝산의 농도 비율과 절대 농도에 따라 달라집니다. 완충제 범위는 산 또는 염기를 추가할 때 pH의 현저한 증가 또는 감소를 억제하는 pH 범위입니다. 범위는 pKa보다 한 단위 높거나 낮습니다.

따라서 효과적인 완충제가 되려면 약산 대비 염기 또는 약염기 대비 산 비율이 10대 1과 1대 10이 되어야 합니다. 핸더슨-하셀바흐 방정식을 풀어서 이러한 값을 지원할 수 있습니다. 산의 농도가 염기의 10배인 경우 pH는 pKa보다 1단위 적습니다.

이에 반해 염기의 농도가 산의 10배인 경우 pH는 pKa보다 1단위 더 많습니다. 완충제는 약산과 짝염기의 농도가 같고 pH가 pKa와 같을 때 완충 범위의 중간에서 가장 효과적입니다. 약산과 염기의 양의 차이가 증가함에 따라 완충제의 효과도 떨어집니다.

따라서 각각 1몰의 아세트산과 아세테이트를 포함한 완충제 A가 0.1몰의 아세트산과 1몰의 아세테이트를 함유한 완충제 B보다 더 효과적일 것입니다. 약산과 염기의 절대 농도는 또한 완중제의 효과도 결정합니다. 약산과 염기의 농도가 더 클수록 더 강한 산이나 염기를 중화시킬 수 있습니다.

따라서 각각 1몰의 포름산과 포름산염을 포함한 완충제가 각각 0.1 몰의 양을 가진 완충제보다 더 효과적입니다. 완충 용량은 pH가 크게 변경되기 전에 완중제가 중화시킬 수 있는 강산 또는 염기의 양입니다. 따라서, 약산과 그 짝염기의 농도가 높을 수록 그리고 약산과 염기의 비율이 1에 근접할 때 산과 염기의 완충 용량은 증가합니다.

완충 용액 효과

완충 용액은 pH를 상대적으로 일정하게 유지하는 데 무제한의 용량을 갖고 있지 않습니다. 대신, pH 변화에 저항하는 완충 용액의 능력은 상당한 양의 짝산-염기 쌍의 존재에 달려 있습니다. 완충 용액 쌍의 어느 구성원의 농도를 실질적으로 낮추기 위해 충분한 강산 또는 염기가 첨가되면 용액 내의 완충 작용이 손상됩니다.

완충 용량은 pH가 크게 변하기 전에 주어진 부피의 완충 용액에 첨가할 수 있는 산 또는 염기의 양(보통 1단위)입니다. 완충 용량은 완충 혼합물에 있는 약산과 그 짝염기의 양에 따라 달라집니다. 예를 들어, 1.0M의 아세트산과 1.0M의 아세트산나트륨 용액 1L는 0.10M의 아세트산과 0.10M의 아세트산나트륨 용액 1L보다 완충 용량이 더 큽니다. 동일한 pH. 첫 번째 용액은 더 많은 아세트산과 아세테이트 이온을 포함하기 때문에 더 많은 완충 용량을 갖습니다.

적합한 완충 혼합물 선택

완충 용액 혼합물을 선택하는 데에는 두 가지 유용한 경험 법칙이 있습니다.

<ol>
	<li>좋은 완충 용액 혼합물은 두 성분의 농도가 거의 동일해야 합니다. 완충 용액은 일반적으로 완충 용액 쌍의 한 구성 요소가 다른 구성 요소의 약 10% 미만일 때 그 유용성을 잃었습니다.</li>
	<li>약산과 그 염은 pH 7 미만의 완충제로 더 좋습니다. 약염기와 그 염은 7보다 큰 pH에 대한 완충 용액으로 더 좋습니다.</li>
</ol>

혈액은 완충 용액의 중요한 예이며, 완충 작용을 담당하는 주요 산과 이온은 탄산(H_2CO_3)과 중탄산염 이온(HCO_3^−)입니다. 하이드로늄 이온이 혈류에 도입되면 주로 다음 반응에 의해 제거됩니다.

<img alt="Eq1" src="/files/ftp_upload/11412/11412_Equation_1v2.png" style="text-align: center; height: 50px;" />

첨가된 수산화물 이온은 다음 반응에 의해 제거됩니다.

<img alt="Eq1" src="/files/ftp_upload/11412/11412_Equation_2.png" style="text-align: center; height: 50px;" />

따라서 첨가된 강산 또는 염기는 완충제 쌍의 훨씬 더 약한 산 또는 염기로 효과적으로 변환됩니다(H_3O^+는 H_2CO_3로 변환되고 OH^-는 HCO_3^-로 변환됩니다). 따라서 인간 혈액의 pH는 완충 용액 쌍 pK_a에 의해 결정된 값(이 경우 7.35)에 매우 가깝게 유지됩니다. 혈액 pH의 정상적인 변화는 일반적으로 0.1 미만이며, 0.4 이상의 pH 변화는 치명적일 가능성이 높습니다.

이 문서는 에서 발췌되었습니다 <a href="https://openstax.org/books/chemistry-2e/pages/14-6-buffers"> Openstax, Chemistry 2e, Section 14.6: Buffers</a>.

완충 용액은 pH를 비교적 일정하게 유지할 수 있는 무제한의 용량을 가지고 있지 않습니다. 대신, pH 변화에 저항하는 완충 용액의 능력 ...

Education

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JoVE Core

화학

하나

Acid-base and Solubility Equilibria

산-염기 및 용해도 평형

KEY TERMS AND CONCEPTS

주요 용어 및 개념

buffer capacity and ph stability

Buffer solutions help keep pH fairly constant, but they do not work forever. Their resistance to pH change depends on having enough of a weak acid and its conjugate base, or a weak base and its conjugate acid, in the mixture. When too much strong acid or strong base is added, one part of the buffer pair is used up and the buffer action becomes weaker.

Buffer capacity describes how much acid or base can be added to a fixed volume of buffer before the pH changes noticeably, usually by about one pH unit. Capacity depends on the amount of weak acid and conjugate base present. A 1.0 M acetic acid and 1.0 M sodium acetate solution has a greater buffer capacity than a 0.10 M acetic acid and 0.10 M sodium acetate solution, even though both have the same pH. The more buffer components present, the more acid or base the solution can handle.

Good buffer mixtures usually have similar concentrations of both components. A buffer often stops being useful when one component falls below about 10% of the other. Weak acids and their salts work best for buffers with pH values below 7. Weak bases and their salts work best for buffers with pH values above 7.

Blood is a key example of a buffered solution. Its main buffering pair is carbonic acid, H2CO3, and bicarbonate ion, HCO3−. When hydronium ion enters the bloodstream, it is removed by reaction with bicarbonate. When hydroxide ion is added, it is removed by reaction with carbonic acid. In both cases, the strong acid or base is turned into the weaker member of the buffer pair.

This buffering keeps human blood near a pH of 7.35, which matches the pKa of the buffer pair. Normal blood pH changes are usually less than 0.1. A change of 0.4 or more can be fatal, which shows why buffer systems are so important in living things.

Buffer Capacity and pH Stability

common ion effect

Common Ion Effect on Solubility

The common ion effect reduces the solubility of an ionic compound in a solution that already contains one of its ions. In aqueous solutions, this means less of the solid dissolves than it would in pure water. The effect follows the law of mass action and can be explained with Le Châtelier’s principle. Silver iodide shows this equilibrium clearly. When silver(I) ions or iodide ions are added, the dissolution equilibrium shifts left. More AgI precipitates, and the concentrations of dissolved Ag+...

buffers

Buffer Solutions and pH Stability

Buffer solutions are mixtures that help keep pH steady. A buffer contains an appreciable amount of a weak conjugate acid-base pair. It resists a change in pH when small amounts of a strong acid or a strong base are added. One common buffer uses a weak acid and its salt. Acetic acid and sodium acetate form this type of system, written as CH3COOH(aq) + CH3COONa(aq). Another buffer uses a weak base and its salt. A solution of ammonia and ammonium chloride, NH3(aq) + NH4Cl(aq), is an example. The...

calculating ph changes in a buffer solution

calculating-ph-changes-in-a-buffer-solution

Buffer pH After Acid or Base Addition

A buffer solution can resist sudden pH changes when a strong acid or strong base is added. It does not stop the pH from changing completely. The shift is usually small, as long as the addition stays within the buffer’s capacity. To find the new pH, first determine how the added acid or base changes the amounts of the buffer components. A buffer contains a weak acid and its conjugate base, or the reverse pair. These amounts are found with stoichiometric calculations, which track how much of...

buffer effectiveness

Buffer solutions help keep pH fairly constant, but they do not work forever. Their resistance to pH change depends on having enough of a weak acid and its conjugate base, or a weak base and its conjugate acid, in the mixture. When too much strong acid or strong base is added, one part of the buffer pair is used up and the buffer action becomes weaker. Buffer capacity describes how much acid or base can be added to a fixed volume of buffer before the pH changes noticeably, usually by about one...

titration calculations: strong acid - strong base

titration-calculations-strong-acid-strong-base

pH at Key Points in Acid-Base Titration

pH at key points in a strong acid and strong base titration can be calculated from the amounts of acid and base present. The example uses 25.00 mL of 0.100 M HCl, a strong acid, and 0.100 M NaOH, a strong base. The pH changes as base is added, and each stage follows a different calculation. At 0 mL of titrant, the pH comes from the ionization of HCl. Because HCl is a strong acid, its ionization is complete. The hydronium ion molarity is 0.100 M, and the pH is found from that value. At 12.50...

titration calculations: weak acid - strong base

titration-calculations-weak-acid-strong-base

pH at Key Points in a Weak Acid Titration

pH at key points in a weak acid titration can be calculated for an acetic acid solution as sodium hydroxide is added. The example uses 25.00 mL of 0.100 M CH3CO2H titrated with 0.100 M NaOH. The reaction is shown as the weak acid reacts with the strong base. The initial pH is found from the acetic acid solution using the usual ICE approach. ICE means initial, change, and equilibrium, and it is a common way to set up equilibrium problems. The solution pH is calculated before any titrant is...

indicators

Acid-Base Color Changes

Acid-base indicators show color changes that help identify the pH of a solution. These are often weak organic acids or weak organic bases, and they respond to changes in hydronium ion concentration in dilute aqueous solution. Phenolphthalein is one example. It is colorless when the hydronium ion concentration is greater than 5.0 × 10^-9 M, which means the pH is below 8.3. In more basic solutions, where the hydronium ion concentration is lower, it turns red or pink. Methyl orange shows how an...

titration of a polyprotic acid

Polyprotic Acid Equivalence Points

Polyprotic acids have more than one ionizable hydrogen, so they lose protons in steps during titration. When the acid dissociation constants for those protons are far enough apart, the titration curve shows a separate equivalence point for each hydrogen that can ionize. Diprotic acids have two ionizable hydrogens and produce two equivalence points. Triprotic acids have three ionizable hydrogens and produce three equivalence points. Carbonic acid, H2CO3, is a weak diprotic acid. Carbonic acid...

solubility equilibria

Predicting Precipitation with Ksp

Solubility equilibria explain how a dissolved ionic compound and its solid form can exist in balance. In a saturated solution, the solute dissolves and the ions precipitate at equal rates. These ideas matter in real settings such as tooth decay and water purification, where compound solubility affects the outcome. A substance with finite solubility can form a saturated solution when more solute is added than the solvent can dissolve. The result is a heterogeneous mixture of the saturated...

factors affecting solubility

Common Ion Effects on Solubility

The common ion effect lowers the solubility of an ionic compound in water when the solution already contains one of the ions made by that compound. It follows the law of mass action and can be explained with Le Châtelier’s principle. When a common ion is added, the dissolution equilibrium shifts toward the solid, so less of the salt stays dissolved. Silver iodide shows this effect clearly. If silver ions or iodide ions are added to a solution of AgI, the equilibrium shifts left and more AgI...

formation of complex ions

Ligands and Complex Ion Stability

Complex ions form when a central atom, often a transition metal cation, binds to ligands. Ligands are ions or molecules around the metal ion. They can be neutral molecules such as H2O or NH3, or ions such as CN− or OH−. In many cases, the ligands act as Lewis bases by donating a pair of electrons to the central atom. These reactions are part of coordination chemistry, which studies coordination complexes. When a metal ion reacts with one or more ligands to form a complex, the equilibrium...

precipitation of ions

Solubility Limits and Selective Precipitation

Solubility limits and selective precipitation explain when ions form a solid in water. The key idea is the balance between an insoluble compound and its dissolved ions. In aqueous solutions that contain calcium ions and carbonate ions, this equilibrium can exist even if the solution was not made by saturating water with calcium carbonate. When two soluble salts are mixed, such as sodium carbonate and calcium nitrate, the ion concentrations determine whether a precipitate forms. If the reaction...

qualitative analysis

Cation Identification by Precipitation Tests

Cation identification by precipitation tests uses selective reactions to find metal ions in a mixture. In qualitative analysis, different reagents are added step by step. Each reaction can produce a characteristic precipitate for a specific group of cations. Other steps, such as changing pH, heating the mixture to dissolve a precipitate again, or adding reagents that form complex ions, help separate ions within the same group. Group IV cations such as Ba2+, Ca2+, and Mg2+ are identified...

henderson-hasselbalch equation

Buffer pH with the Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation helps relate buffer pH to the weak acid in the solution. It comes from the ionization-constant expression for a weak acid. By rearranging that expression to solve for hydronium ion concentration, [H3O+], and then taking the negative logarithm of both sides, the equation can be written in terms of pH. In this form, pKa is defined as the negative logarithm of the weak acid ionization constant, or pKa = −log Ka. The equation shows the connection between pH, pKa,...

acid-base titration curves

Tracking pH During Acid-Base Titration

Acid-base titration uses pH changes to track how a solution responds as titrant is added. A titration curve is a graph of a solution property versus the amount of added titrant. In acid-base titrations, pH is the main property to follow because it changes in a predictable way with the solution makeup and helps show the endpoint. This pattern can be seen with a strong acid and strong base, a strong acid and weak base, or a strong base and weak acid. For example, 25.00 mL of 0.100 M HCl titrated...

key terms and concepts

Learn how buffer capacity depends on weak acid-conjugate base amounts, why buffer mixtures work best near equal parts, and how blood resists pH change.

Buffer Capacity and pH Stability | JoVE K12

chemistry

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• Equilibrium
• Acids/bases

Content elaboration: In the Physical Science course, coefficients were used to balance simple equations. Other representations, including Lewis structures and three-dimensional models, were also used and manipulated to demonstrate the conservation of matter in chemical reactions. In this course, more complex reactions will be studied, classified and represented with balanced chemical equations and three-dimensional models. Classifying reactions into types can be a helpful organizational tool for recognizing patterns of what may happen when two substances are mixed. Teachers should be aware that the common reaction classifications that are often used in high school chemistry courses may lead to misconceptions because they are not based on the actual chemistry, but on surface features that can be similar from one system to another (e.g., exchanging partners), even though the underlying chemistry is not the same. However, these classifications may be useful in making predictions about what happens when two substances are mixed. Some general types of chemical reactions are oxidation/reduction, synthesis, decomposition, single replacement, double replacement (including precipitation reactions and some acid-base neutralizations) and combustion reactions. Some reactions can fit into more than one category. For example, a single replacement reaction can also be classified as an oxidation/reduction reaction. Identification of reactions involving oxidation and reduction as well as indicating what substance is being oxidized and what is being reduced are appropriate in this course. However, balancing complex oxidation/reduction reactions is reserved for more advanced study. Organic molecules release energy when undergoing combustion reactions and are used to meet the energy needs of society (e.g., oil, gasoline, natural gas) and to provide the energy needs of biological organisms (e.g., cellular respiration). When a reaction between two ionic compounds in aqueous solution results in the formation of a precipitate or molecular compound, the reaction often occurs because the new ionic or covalent bonds are stronger than the original ion-dipole interactions of the ions in solution. Laboratory experiences (3-D or virtual) with different types of chemical reactions should be provided. Reactions occur when reacting particles collide in an appropriate orientation and with sufficient energy. The rate of a chemical reaction is the change in the amount of the reactants or products in a specific period of time. Increasing the probability or effectiveness of the collisions between the particles increases the rate of the reaction. Therefore, changing the concentration of the reactants, changing the temperature or the pressure of gaseous reactants, or using a catalyst, can change the reaction rate. Likewise, the collision theory can be applied to dissolving solids in a liquid solvent and can be used to explain why reactions are more likely to occur between reactants in the aqueous or gaseous state than between solids. The rate at which a substance dissolves should not be confused with the amount of solute that can dissolve in a given amount of solvent (solubility). Mathematical treatment of reaction rates is reserved for more advanced study. Computer simulations can help visualize reactions from the perspective of the kinetic-molecular theory. In middle school, the differences between potential and kinetic energy and the particle nature of thermal energy were introduced. For chemical systems, potential energy is in the form of chemical energy and kinetic energy is in the form of thermal energy. The total amount of chemical energy and/or thermal energy in a system is impossible to measure. However, the energy change of a system can be calculated from measurements (mass and change in temperature) from calorimetry experiments in the laboratory. Conservation of energy is an important component of calorimetry equations. Thermal energy is the energy of a system due to the movement of its particles. The thermal energy of an object depends upon the temperature and chemical composition. Some materials require little energy to change their temperature and other materials require a great deal to change their temperature by the same amount. Specific heat is a measure of how much energy is needed to change the temperature of a specific mass of material a specific amount. Specific heat values can be used to calculate the thermal energy change, the temperature (initial, final or change in) or mass of a material in calorimetry. Water has a particularly high specific heat capacity, which is important in regulating Earth's temperature. As studied in middle school, chemical energy is the potential energy associated with chemical systems. Chemical reactions involve valence electrons forming bonds to yield more stable products with lower energies. Energy is required to break interactions and bonds between the reactant atoms and energy is released when an interaction or bond is formed between the atoms in the products. Molecules with weak bonds (e.g., ATP) are less stable and tend to react to produce more stable products, releasing energy in the process. Generally, energy is transferred out of the system (exothermic) when the products have stronger bonds than the reactants and is transferred into the system (endothermic) when the reactants have stronger bonds than the products. Predictions of the energy requirements (endothermic or exothermic) of a reaction can be made given a table of bond energies. Graphic representations can be drawn and interpreted to represent the energy changes during a reaction. The role of energy in determining the spontaneity of chemical reactions is dealt with conceptually in this course. Entropy and its influence on the spontaneity of reactions are reserved for more advanced study. All reactions are reversible to a degree and many reactions do not proceed completely toward products but appear to stop progressing before the reactants are all used up. At this point, the amounts of the reactants and the products appear to be constant and the reaction can be said to have reached dynamic equilibrium. Dynamic equilibrium means the rate of the reverse reaction is equal to the rate of the forward reaction so there is no apparent change in the reaction. If a chemical system at equilibrium is disturbed by a change in the conditions of the system (e.g., increase or decrease in the temperature, pressure on gaseous equilibrium systems, concentration of a reactant or product), then the equilibrium system will respond by shifting to a new equilibrium state, reducing the effect of the change (Le Chatelier's Principle). If products are removed as they are formed during a reaction, then the equilibrium position of the system is forced to shift to favor the products. In this way, an otherwise unfavorable reaction can be made to occur. Mathematical treatment of equilibrium reactions is reserved for advanced study. Computer simulations can help visualize the progression of a reaction to dynamic equilibrium and the continuation of both the forward and reverse reactions after equilibrium has been attained. Properties of acids and bases and the ranges of the pH scale were introduced in Physical Science. In this course, the structural features of molecules are explored to further understand acids and bases. Acids often result when hydrogen is covalently bonded to an electronegative element and is easily dissociated from the rest of the molecule to bind with water to form a hydronium ion (H3O+). The acidity of an aqueous solution can be expressed as pH, where pH can be calculated from the concentration of the hydronium ion. Bases are likely to dissociate in water to form a hydroxide ion. Acids can react with bases to form a salt and water. Such neutralization reactions can be studied quantitatively by performing titration experiments.

Buffer Capacity and pH Stability

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