Name: ph at key points in a weak acid titration
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Titration Calculations: Weak Acid - Strong Base

약산이나 염기 적정의 다른 단계에서 pH는 여러 가지 적정 점에서 다른 방법을 사용하여 계산됩니다. 약산 또는 염기가 pH의 주요 결정 인자인 경우 Ka 또는 Kb와 ICE 표 또는 핸더슨-하셀바흐 방정식이 사용됩니다. 중화 반응 후 강산이나 염기가 존재할 경우 남은 하이드로늄 또는 수산화 이온의 농도를 pH를 계산하는 데 사용됩니다.

0.10몰의 아세트산 용액 50밀리리터의 초기 pH는 2.87이며 아세트산이 주 기부자이기 때문에 Ka와 ICE 표를 사용하여 계산합니다. 0.0050몰의 아세테이트가 함유된 이 용액을 0.10몰의 수산화나트륨으로 적정화하면 수산화이온이 아세트산과 반응하여 완충제인 아세테이트를 생성합니다. 따라서 0.0010몰의 수산화 이온을 함유한 수산화나트륨 10밀리리터를 첨가하면 0.0010몰의 아세테이트가 생성되고 0.0040몰의 아세트산이 남습니다.

완충제의 pH는 이러한 값을 핸더슨-하셀바흐 방정식으로 대체하여 계산할 수 있으며 이 값은 4.14와 같습니다. 수산화나트륨 25밀리리터를 첨가하면 아세트산의 초기 몰의 절반이 아세테이트로 전환됩니다. 이 때 아세트산과 아세테이트 이온의 양이 같기 때문에 pH는 pKa와 같습니다.

수산화나트륨을 50 밀리리터까지 첨가하면 모든 아세트산 분자가 아세테이트로 전환되고 당량점에 도달합니다. 아세테이트 이온은 염기이므로 당량점은 염기 영역에 있습니다. 아세테이트 이온의 농도는 몰 수를 용액의 총 부피로 나누어 계산합니다.

아세테이트 이온은 당량점에서 pH의 주요 기부자이기 때문에 pH는 아세테이트 이온 및 ICE 표를 사용하여 결정됩니다. 아세테이트의 Kb는 식 Kw Ka x Kb을 사용하여 계산하며 값은 5.6 10⁻¹⁰입니다. 평형 농도를 Kb에 대한 식에 대입하면 5.3 10⁻⁶몰의 수산화물 농도를 얻을 수 있습니다.

용액의 pOH와 pH는 각각 5.28과 8.72입니다. 용액에 수산화나트륨을 추가하면 아세테이트 이온과 수산화나트륨의 혼합물이 생깁니다. 그러나 수산화나트륨은 아세테이트보다 강한 염기이므로 그것의 최종 농도에 의해 용액의 pH가 결정됩니다.

따라서 수산화나트륨 70밀리리터를 용액에 첨가하면 수산화 이온의 최종 농도는 첨가된 수산화 이온의 총 몰인 0.0070 몰에서 아세트산의 총 몰인 0.0050 몰을 빼고 용액의 총 부피인 120밀리리터 또는 0.12리터로 나누어 계산할 수 있습니다. 수산화물 이온 농도는 0.017 몰이므로 용액의 pOH와 pH는 각각 1.78과 12.22로 계산됩니다.

적정 계산: 약산 - 강염기

적정 용액의 pH 계산: 약산/강염기

0.100 M CH_3CO_2H 25.00 mL를 0.100 M NaOH로 적정하는 경우 반응은 다음과 같이 나타낼 수 있습니다.

<img alt="Eq1" src="/files/ftp_upload/11414/11414_Equation_1.png" style="text-align: center; height: 50px;" />

다양한 부피의 NaOH 적정제를 첨가한 후 적정 용액의 pH는 다음과 같이 계산할 수 있습니다.

(a) 초기 pH는 일반적인 ICE 접근법으로 아세트산 용액에 대해 계산됩니다.

<img alt="Eq2" src="/files/ftp_upload/11414/11414_Equation_2.png" style="text-align: center; height: 155px;" />

(b) 산과 적정제는 모두 단양성이며 시료와 적정제 용액은 동일하게 농축됩니다. 따라서 이 적정제의 부피는 당량점을 나타냅니다. 강산의 예와는 달리, 이 경우 반응 혼합물에는 약한 짝염기(아세트산 이온)가 포함되어 있습니다. 용액 pH는 다음과 같은 농도로 존재하는 아세테이트의 염기 이온화를 고려하여 계산됩니다.

<img alt="Eq3" src="/files/ftp_upload/11414/11414_Equation_3.png" style="text-align: center; height: 72px;" />

아세테이트의 염기 이온화는 다음 방정식으로 표현됩니다.

<img alt="Eq4" src="/files/ftp_upload/11414/11414_Equation_4.png" style="text-align: center; height: 115px;" />

x << 0.0500이라고 가정하면 일반적인 ICE 접근 방식을 통해 pH를 계산할 수 있습니다.

<img alt="Eq5" src="/files/ftp_upload/11414/11414_Equation_5.png" style="text-align: center; height: 150px;" />

이 적정의 당량점에서의 pH는 강염기로 약산을 적정할 때 예상되는 것처럼 7보다 상당히 높다는 점에 유의하십시오.

(c) 적정제 부피 = 12.50mL. 이 부피는 적정제의 화학양론적 양의 1/2을 나타내므로 아세트산의 1/2이 중화되어 동일한 양의 아세트산 이온을 생성합니다. 따라서 이러한 짝산-염기 파트너의 농도는 동일합니다. pH를 계산하는 편리한 접근 방식은 핸더슨-하셀바흐 방정식을 사용하는 것입니다.

<img alt="Eq6" src="/files/ftp_upload/11414/11414_Equation_6.png" style="text-align: center; height: 115px;" />

(pH = 약산 적정의 반당량점에서의 pKa)

(d) 적정제 부피 = 37.50mL. 이 부피는 화학양론적 과량의 적정제와 적정 생성물, 아세테이트 이온 및 과량의 강한 적정제를 모두 포함하는 반응 용액을 나타냅니다. 이러한 용액에서 용액 pH는 주로 과잉 강염기의 양에 따라 결정됩니다.

<img alt="Eq7" src="/files/ftp_upload/11414/11414_Equation_7.png" style="text-align: center; height: 155px;" />

이 문서는 에서 발췌되었습니다 <a href="https://openstax.org/books/chemistry-2e/pages/14-7-acid-base-titrations"> Openstax, Chemistry 2e, Section 14.7: Acid-base Titrations.</a>

적정 용액의 pH 계산: 약산성/강염기
0.100 M NaOH로 0.100 M CH3CO2H의 25.00 mL를 적정하기 위해 반응은 ...

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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 added.

When 25.00 mL of NaOH has been added, the titration reaches the equivalence point. The acid and titrant are both monoprotic, and the sample and titrant have the same concentration. At this point, the mixture contains acetate ion, which is the weak conjugate base of acetic acid.

Because acetate ion is present, the pH at the equivalence point is found from its base ionization. The acetate concentration is used in an ICE setup, and x is assumed to be much less than 0.0500. The pH at this point is significantly greater than 7, which is expected for a weak acid titrated with a strong base.

At 12.50 mL of NaOH, the titration is at the half-equivalence point. This is one-half of the stoichiometric amount of titrant, so one-half of the acetic acid has been neutralized. Equal amounts of acetic acid and acetate are present, and the Henderson-Hasselbalch equation is a convenient way to find the pH. At the half-equivalence point, pH equals pKa.

At 37.50 mL of NaOH, there is a stoichiometric excess of titrant. The reaction mixture contains acetate ion from the titration product and extra strong base. In this case, the pH is determined mainly by the amount of excess base.

pH at Key Points in a Weak Acid Titration

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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+...

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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...

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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...

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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...

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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

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...

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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...

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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...

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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...

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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...

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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...

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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...

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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...

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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,...

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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

Follow how pH changes during a weak acid titration with strong base, including the initial solution, half-equivalence point, equivalence point, and excess base.

pH at Key Points in a Weak Acid Titration | JoVE K12

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For gas-phase equilibria, changes in the concentrations of reactants and products can occur with altered volume and pressure. The partial pressure, P, of ...

Mixtures of Acids

The pH of a solution containing an acid can be determined using its acid dissociation constant and its initial concentration. If a solution contains two ...

le-chatelier-s-principle-changing-temperature

Le Chatelier's Principle: Changing Temperature

Consistent with the law of mass action, an equilibrium stressed by a change in concentration will shift to re-establish equilibrium without any change in ...

Ions as Acids and Bases

Salts with Acidic Ions
Salts are ionic compounds composed of cations and anions, either of which may be capable of undergoing an acid or base ionization ...

Bronsted-Lowry Acids and Bases

The acid-base reaction class has been studied for quite some time. In 1680, Robert Boyle reported traits of acid solutions that included their ability to ...

Determining the pH of Salt Solutions

The pH of a salt solution is determined by its component anions and cations. Salts that contain pH-neutral anions and the hydronium ion-producing cations ...

acidbase-strengths-and-dissociation-constants

Acid/Base Strengths and Dissociation Constants

The relative strength of an acid or base is the extent to which it ionizes when dissolved in water. If the ionization reaction is essentially complete, ...

Polyprotic Acids

Acids are classified by the number of protons per molecule that they can give up in a reaction. Acids such as HCl, HNO3, and HCN that contain one ...

Chemical reactions
• Types of reactions
• Kinetics
• Energy
• 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.

pH at Key Points in a Weak Acid Titration

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