Name: ph at key points in acid-base titration
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 2 min 28 s
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Titration Calculations: Strong Acid - Strong Base

강산을 강염기로 적정화하거나 그 반대인 경우 적정화 전체의 pH는 남은 하이드로늄 또는 수산화 이온의 농도를 측정하여 계산할 수 있습니다. 예를 들어 0.10 몰의 염산 50밀리리터는 0.10 몰 하이드로늄과 염화물 이온으로 해리되며 이 때 초기 pH는 1입니다. 0.10 몰의 수산화나트륨 25 밀리리터를 첨가하면 새로운 수소 이온 농도는 총 몰의 히드로늄이온 이온에서 총 몰의 수산화물을 빼고 75 밀리리터 또는 0.075리터의 총 부피로 나누어 계산할 수 있습니다.

0.10 몰의 염산 50밀리리터에는 0.0050 몰의 하이드로늄이 들어 있고 0.10몰의 수산화나트륨 25밀리리터에는 0.0025 몰의 수산화물 이온이 들어 있습니다. 이 값들을 방정식에 대입하면 하이드로늄 이온 농도는 0.033 몰과 같습니다. 따라서 솔루션의 pH는 1.48로 상승했습니다.

수산화나트륨을 50밀리리터까지 첨가하면 염산의 모든 하이드로늄 이온은 중화되며 pH가 7로 상승할 때 당량점에 도달합니다. 당량점 이상에서 용액의 pH는 염산이 완전히 중화되었기 때문에 수산화 이온에 의해 결정됩니다. 0.0070몰의 수산화물 이온을 포함하는 수산화나트륨 용액 70밀리리터를 용액에 첨가하면 수산화물 이온의 최종 농도는 수산화물 이온의 총 몰에서 히드로늄이온의 초기 몰 수를 빼고 용액의 총 부피로 나누어 계산할 수 있습니다.

최종 수산화물 이온 농도는 0.017 몰이므로 용액의 pOH와 pH는 각각 1.78과 12.22입니다.

적정 계산: 강산 - 강염기

적정 용액의 pH 계산: 강산/강염기

0.100M HCl(강산) 25.00mL에 대해 0.100M 강염기 NaOH를 사용하여 적정을 수행합니다. 첨가된 염기 용액의 다양한 부피에서의 pH는 다음과 같이 계산할 수 있습니다.

(a) 적정제 부피 = 0mL. 용액 pH는 HCl의 산 이온화로 인해 발생합니다. 이는 강산이기 때문에 이온화가 완료되고 하이드로늄 이온 몰농도는 0.100M입니다. 그러면 용액의 pH는 다음과 같습니다.

<img alt="Eq1" src="/files/ftp_upload/11413/11413_Equation_1.png" style="text-align: center; height: 50px;" />

(b) 적정제 부피 = 12.50mL. 산 샘플과 염기 적정제는 둘 다 일양성자이고 동일하게 농축되어 있기 때문에 이 적정제 첨가에는 화학양론적 양보다 적은 양의 염기가 포함되므로 샘플 내 과량의 산과의 반응으로 염기가 완전히 소모됩니다. 남은 산의 농도는 초기 양에서 소비된 양을 뺀 후 용액 부피로 나누어 계산됩니다.

<img alt="Eq2" src="/files/ftp_upload/11413/11413_Equation_2v2.png" style="text-align: center; height: 110px;" />

(c) 적정제 부피 = 25.00mL. 이 적정제 첨가에는 화학양론적 양의 염기(당량점)가 포함되므로 중화 반응의 생성물만 용액(물과 NaCl)에 존재합니다. 이 염의 양이온이나 음이온은 모두 산-염기 이온화를 겪지 않습니다. 하이드로늄 이온을 생성하는 유일한 과정은 물의 자가원생분해입니다. 용액은 중성이며 pH = 7.00입니다.

(d) 적정제 부피 = 37.50mL. 이는 당량점을 초과하여 적정제를 첨가하는 것을 포함합니다. 그런 다음 용액 pH는 수산화물 이온의 농도를 사용하여 계산됩니다.

<img alt="Eq3" src="/files/ftp_upload/11413/11413_Equation_3v2.png" style="text-align: center; height: 170px;" />

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: 강산성/강염기
25.00 mL의 0.100 M HCl(강산)에 대해 0.100 M의 ...

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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 mL of titrant, the base added is less than the stoichiometric amount needed to neutralize the acid. Since the acid sample and the base titrant are both monoprotic and equally concentrated, the added base is fully consumed by the excess acid. The remaining acid concentration is found by subtracting the amount reacted and dividing by the total solution volume.

At 25.00 mL of titrant, the equivalence point is reached. This means a stoichiometric amount of base has been added, so only the neutralization products remain in solution. Those products are water and NaCl, and neither ion from this salt undergoes acid-base ionization. The only source of hydronium ions is the autoprotolysis of water, so the solution is neutral with a pH of 7.00.

At 37.50 mL of titrant, the base is in excess beyond the equivalence point. In this region, the pH is calculated from the hydroxide ion concentration. This example shows how pH calculations shift from excess acid, to equivalence, to excess base during a strong acid-strong base titration.

pH at Key Points in Acid-Base Titration

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

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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...

titration calculations: strong acid - strong base

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...

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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,...

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

Follow pH changes in a strong acid-strong base titration using HCl and NaOH. Learn how to calculate pH before, at, and after the equivalence point.

pH at Key Points in Acid-Base Titration | JoVE K12

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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. 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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 Acid-Base Titration

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