Name: ph at key points in a weak acid titration
Uploaded: 2020-09-24T15:06:51.000Z
Duration: 3 min 55 s
Channel: JoVE

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Titration Calculations: Weak Acid - Strong Base

Il pH in diversi stadi di una titolazione di un acido o di una base debole è calcolato usando metodi diversi durante i vari punti della titolazione. Se un acido o una base debole è il principale determinante del pH, vengono usati la Ka o Kb e una tabella ICE, o l'equazione di Henderson-Hasselbalch. Se dopo la reazione di neutralizzazione, è presente un acido o una base forte, per calcolare il pH viene usata la concentrazione degli ioni idronio o idrossido rimanenti.

Il pH iniziale di 50 millilitri di una soluzione di acido acetico 0, 10 molare è 2, 87 e viene calcolato utilizzando il valore della Ka e una tabella ICE, poiché l'acido acetico è il principale contribuente. Se questa soluzione contenente 0, 0050 moli di acetato viene titolata con idrossido di sodio 0, 10 molare, gli ioni idrossido reagiscono con acido acetico per produrre acetato, risultando in un tampone. Pertanto, quando si aggiungono 10 millilitri di idrossido di sodio contenente 0, 0010 moli di ioni idrossido, si formano 0, 0010 moli di acetato, e rimangono 0, 0040 moli di acido acetico.

Il pH del tampone può essere calcolato sostituendo questi valori nell'equazione di Henderson-Hasselbalch ed è uguale a 4, 14. Quando si aggiungono 25 millilitri di idrossido di sodio, metà delle moli iniziali di acido acetico vengono convertite in acetato. A questo punto, il pH è uguale al pKa, poiché la quantità di acido acetico e lo ione acetato sono uguali.

L'ulteriore aggiunta di idrossido di sodio fino a 50 millilitri converte tutte le molecole di acido acetico in acetato, e il punto di equivalenza viene raggiunto. Poiché gli ioni acetato sono basici, il punto di equivalenza si trova nella regione di base. La concentrazione dello ione acetato è calcolata dividendo il numero di moli per il volume totale della soluzione.

Il pH è determinato usando la Kb per gli ioni acetato e una tabella ICE, poiché lo ione acetato è il principale contributuente del pH nel punto di equivalenza. Il Kb per l'acetato si calcola con la formula:Kw è uguale a Ka per Kb ed è uguale a 5, 6 10⁻¹⁰. Sostituendo le concentrazioni di equilibrio nell'espressione per Kb si ottiene la concentrazione di idrossido, 5, 3 10⁻⁶ molare.

Il pOH e il pH della soluzione sono rispettivamente 5, 28 e 8, 72. Un'ulteriore aggiunta di idrossido di sodio nella soluzione produce una miscela di ioni acetato e idrossido di sodio. Tuttavia, la concentrazione finale di idrossido di sodio determina il pH della soluzione, poiché è una base più forte dell'acetato.

Pertanto, se si aggiungono 70 millilitri di idrossido di sodio nella soluzione, la concentrazione finale degli ioni idrossido può essere calcolata sottraendo le moli totali di acido acetico, 0, 0050 moli, dalle moli totali di ioni idrossido aggiunti, 0, 0070 moli, e dividendo per il volume totale della soluzione, 120 millilitri o 0, 12 litri. Poiché la concentrazione di ioni idrossido è 0, 017 molare, il pOH e il pH della soluzione sono calcolati rispettivamente di 1, 78 e 12, 22.

calcoli di titolazione: acido debole - base forte

Calcolo del pH per soluzioni di titolazione: acido debole/base forte

Per la titolazione di 25,00 mL di CH_3CO_2H 0,100 M con NaOH 0,100 M, la reazione può essere rappresentata come:

<img alt="Eq1" src="/files/ftp_upload/11414/11414_Equation_1.png" style="text-align: center; height: 50px;" />

Il pH della soluzione di titolazione dopo l'aggiunta dei diversi volumi di titolante NaOH può essere calcolato come segue:

(a) Il pH iniziale viene calcolato per la soluzione di acido acetico nel consueto approccio ICE:

<img alt="Eq2" src="/files/ftp_upload/11414/11414_Equation_2.png" style="text-align: center; height: 155px;" />

(b) L'acido e il titolante sono entrambi monoprotici e le soluzioni del campione e del titolante sono ugualmente concentrate; quindi, questo volume di titolante rappresenta il punto equivalente. A differenza dell'esempio dell'acido forte, la miscela di reazione in questo caso contiene una base coniugata debole (ione acetato). Il pH della soluzione viene calcolato considerando la ionizzazione basica dell'acetato, che è presente ad una concentrazione di

<img alt="Eq3" src="/files/ftp_upload/11414/11414_Equation_3.png" style="text-align: center; height: 72px;" />

La ionizzazione della base dell'acetato è rappresentata dall'equazione

<img alt="Eq4" src="/files/ftp_upload/11414/11414_Equation_4.png" style="text-align: center; height: 115px;" />

Supponendo x << 0,0500, il pH può essere calcolato tramite il consueto approccio ICE:

<img alt="Eq5" src="/files/ftp_upload/11414/11414_Equation_5.png" style="text-align: center; height: 150px;" />

Si noti che il pH al punto equivalente di questa titolazione è significativamente maggiore di 7, come previsto quando si titola un acido debole con una base forte.

(c) Volume del titolante = 12,50 ml. Questo volume rappresenta la metà della quantità stechiometrica del titolante, quindi metà dell'acido acetico è stata neutralizzata per produrre una quantità equivalente di ione acetato. Le concentrazioni di questi partner acido-base coniugati, quindi, sono uguali. Un approccio conveniente per calcolare il pH è l'uso dell'equazione di Henderson-Hasselbalch:

<img alt="Eq6" src="/files/ftp_upload/11414/11414_Equation_6.png" style="text-align: center; height: 115px;" />

(pH = pKa al punto di semiequivalenza in una titolazione di un acido debole)

(d) Volume del titolante = 37,50 ml. Questo volume rappresenta un eccesso stechiometrico di titolante e una soluzione di reazione contenente sia il prodotto di titolazione, lo ione acetato, sia il titolante forte in eccesso. In tali soluzioni, il pH della soluzione è determinato principalmente dalla quantità di base forte in eccesso:

<img alt="Eq7" src="/files/ftp_upload/11414/11414_Equation_7.png" style="text-align: center; height: 155px;" />

Questo testo è adattato da<a href="https://openstax.org/books/chemistry-2e/pages/14-7-acid-base-titrations"> Openstax, Chemistry 2e, Section 14.7: Acid-base Titrations.</a>

Calcoli di titolazione: Acido debole - Base forte

Calcolo del pH per le soluzioni di titolazione: acido debole/base forte
Per la titolazione di 25,00 mL di 0,100 M CH3CO2H con 0,100 M NaOH, la reazione può essere ...

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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+...

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

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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...

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...

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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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Le Chatelier's Principle: Changing Concentration

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Le Chatelier's Principle: Changing Volume (Pressure)

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

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Le Chatelier's Principle: Changing Temperature

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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 ...

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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. 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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