Name: buffer solutions and ph stability
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 56 s
Channel: JoVE

K12 Standards

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Ohio’s Learning Standards and Model Curriculum

9-12

Chemistry

Interactions Of Matter

Chemical reactions

Buffers

용액에 소량의 산이나 염기를 첨가하면 pH가 크게 감소하거나 증가할 수 있습니다. 하지만 많은 화학적 및 생화학적 과정들은 기능을 하기 위해 안정적인 pH가 필요합니다. 완충제는 완충 용량이 초과되지 않을 때 용액의 pH의 극적인 변화를 방지합니다.

완충제에는 약산과 그 짝염기 또는 약염기 및 그 짝산이 포함되어 있습니다. 예를 들어 인간의 혈액은 약산인 탄산과 그것의 짝염기인 중탄산 이온으로 구성된 완충제를 통해 pH를 약 7.4로 유지합니다. 짝 산-염기쌍은 짝산 또는 염기를 중화시키지 않기 때문에 완충제로 됩니다.

예를 들어 아세트산과 아세테이트는 반응할 수 없습니다. 하지만 약산인 아세트산과 약염기인 암모니아를 함께 첨가하면 반응하여 염인 아세트산 암모늄을 생성합니다. 완충제에서 약산은 생성된 수산화 이온과 반응함으로써 추가된 염기를 중화시키지만 그것의 짝염기는 하이드로늄 이온과 반응함으로써 추가된 산을 중화시킵니다.

약염기와 그것의 짝산의 경우에 유사한 메커니즘이 작동합니다. 두 비커 X와 Y에는 각각 pH가 7.2인 동일한 양의 서로 다른 용액이 들어 있습니다. 비커 X의 용액은 완충되지 않습니다.

대조적으로 비커 Y의 용액에는 아세트산 아세테이트 완충제가 포함되어 있습니다. 비커 X에 염산을 첨가하면 하이드로늄 이온의 농도가 높아져 용액의 pH가 갑자기 떨어집니다. 대조적으로 비커 Y의 용액은 염산을 첨가했을 때 거의 일정한 pH를 보입니다.

왜냐하면 완충 성분 중 하나인 아세테이트가 염산과 반응하여 염화 이온과 아세트산을 생성하기 때문입니다. 마찬가지로 비커 X에 수산화나트륨을 첨가하면 수산화 이온의 농도가 증가하여 용액의 pH가 갑자기 증가합니다. 반면 비커 Y의 용액은 수산화나트륨이 첨가되었을 때 완충 성분 중 하나인 아세트산이 수산화나트륨과 반응하여 아세트산나트륨과 물 분자를 생성하므로 pH의 변화가 미미합니다.

이 완충제는 용액 내 짝 산-염기쌍의 농도가 첨가된 강산 또는 염기보다 높은 한 용액 pH의 급격한 변화를 방지할 수 있습니다.

완충 용액

상당한 양의 약한 짝산-염기 쌍을 포함하는 용액을 완충 용액 또는 완충 용액이라고 합니다. 완충 용액은 소량의 강산이나 강염기가 추가될 때 pH 변화에 저항합니다. 아세트산과 아세트산 나트륨의 용액은 약산과 그 염인 CH_3COOH(aq) + CH_3COONa(aq)로 구성된 완충 용액의 예입니다. 약염기와 그 염으로 구성된 완충 용액의 예로는 암모니아와 염화암모늄 용액이 있습니다: NH_3(aq) + NH_4Cl(aq).

완충 용액 작동 방식

완충 용액의 기능을 설명하기 위해 대략 같은 양의 아세트산과 아세트산 나트륨의 혼합물을 생각해 보세요. 용액에 약한 짝산-염기 쌍이 존재하면 첨가된 강산이나 염기를 적당한 양으로 중화하는 능력이 부여됩니다. 예를 들어, 이 용액에 강염기를 첨가하면 하이드로늄 이온이 중화되고 아세트산 이온화 평형이 오른쪽으로 이동하여 감소된 H_3O^+ 농도가 부분적으로 회복됩니다.

<img alt="Eq1" src="/files/ftp_upload/11409/11409_Equation_1.png" style="text-align: center; height: 50px;" />

마찬가지로, 이 완충 용액에 강산을 첨가하면 아세트산 이온이 중화되어 위의 이온화 평형이 오른쪽으로 이동하고 [H_3O^+]가 원래 값에 가깝게 돌아갑니다. 그림 1은 강산과 염기를 첨가할 때 완충 용액의 변화를 그래픽으로 보여줍니다. 용액의 완충 작용은 본질적으로 첨가된 강산과 염기가 완충액의 결합 쌍을 구성하는 약산과 염기로 전환된 결과입니다. 강한 산과 염기의 완전한 이온화에 비해 약한 산과 염기는 약간만 이온화됩니다. 따라서 용액 pH는 완충되지 않은 용액에서보다 훨씬 덜 급격하게 변합니다.

<img alt="Image1" src="/files/ftp_upload/11409/11409_Image_1.png" style="text-align: center; height: 400px;" />

Figure 1. 아세트산과 아세트산염 혼합물의 완충 작용.



이 문서는 에서 발췌되었습니다 <a href="https://openstax.org/books/chemistry-2e/pages/14-6-buffers"> Openstax, Chemistry 2e, Section 14.6: Buffers</a>.

상당한 양의 약한 공액 산-염기 쌍을 포함하는 용액을 완충 용액 또는 완충액이라고 합니다. 완충 용액은 변화에 저항합니다 ...

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 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 buffer action comes from the weak conjugate acid-base pair in the solution. In a mixture with roughly equal amounts of acetic acid and sodium acetate, added strong base is neutralized by hydronium ion. This shifts the acetic acid ionization equilibrium to the right and partially restores the lowered H3O+ concentration. Added strong acid has a similar effect in the opposite direction. It neutralizes acetate ion, shifts the same equilibrium to the right, and brings [H3O+] back near its original value.

This buffering action works because the added strong acid and base are converted into the weak acid and weak base that make up the buffer pair. Weak acids and weak bases ionize only slightly compared with strong acids and strong bases, which ionize completely. As a result, the pH changes much less than it would in a solution without a buffer.

Buffer Solutions 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 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 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...

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 solutions resist pH change using weak acid-base pairs like acetic acid/acetate and ammonia/ammonium chloride.

Buffer Solutions and pH Stability | JoVE K12

chemistry

JoVE Core 화학은 간결하고 이해하기 쉬운 애니메이션을 통해 일반 화학 개념을 설명합니다. 또한, 과학자가 직접 고전실험과 연구실험을 진행하는 실사 영상을 제공합니다.

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• 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 Solutions and pH Stability

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