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Calculating the pH of a Buffer

Pearson
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Our life depends on the pH of solutions and the existence of buffers. Without the existence of buffers, the pH of our blood would fluctuate with catastrophic consequences. In the absence of a buffer, adding a few drops of a strong acid to water at pH 7.0 will lower the pH significantly, in this case to pH 3.0. However, adding the same number of drops of strong acid to a buffer does not appreciably alter the pH. With the buffer, it only changes from pH 7.0 to 6.9. A buffer solution maintains the pH of a solution by neutralizing small amounts of added acid or base. In a buffer, an acid must be available to react with any OH- that is added, and a base must be available to react with any added H3O+ that is added. However, that acid and base cannot neutralize each other. Therefore, a combination of an acid-base conjugate pair is used in buffers. Most buffer solutions consist of nearly equal concentrations of a weak acid and a salt containing its conjugate base Buffers may also contain a weak base and the salt of the weak base, which contains its conjugate acid. Which of the following represents a buffer system? Is it a, b, c or d? The correct answer is (d). A buffer system consists of a weak acid and a salt containing its conjugate base. H2CO3 is a weak acid and NaHCO3 is the salt containing its conjugate base. A buffer can be made from the weak acid, acetic acid, and its salt, sodium acetate. As a weak acid, acetic acid dissociates slightly in water to form H3O+ and a very small amount of acetate ion, C2H3O2-. The addition of its salt, sodium acetate, provides a much larger concentration of acetate ion, which is necessary for its buffering capability. When a small amount of acid is added to the buffer, the additional H3O+ combines with the acetate ion, causing the equilibrium to shift in the direction of the reactants, acetic acid and water. There will be a slight decrease in the acetate ion and a slight increase in the acetic acid but both the H3O+ concentration and pH are maintained. If a small amount of base is added to this same buffer solution, it is neutralized by the acetic acid, which shifts the equilibrium in the direction of the products, acetate ion and water. The acetic acid decreases slightly, and the acetate ion increases slightly, but again the H3O+ concentration and pH of the solution are maintained. Consider the buffer system of H2PO4- and HPO42-(aq). If a small amount of H3O+ is added, it is neutralized by: a, b, c or d? The correct answer is (a). When a small amount of acid is added, the additional H3O+ combines with the hydrogen phosphate ion, HPO42- , causing the equilibrium to shift to the direction of the reactants. Let's look at a how we can calculate the pH of a buffer if we are given the concentrations of the weak acid and its conjugate base. The problem is what is the pH of a buffer prepared with 2.0 M HF and 2.5 MF-, if the Ka of hydrofluoric acid is 3.5×10-4? Step 1 is to state the given and needed quantities. In this problem, we are given 2.0 M HF and 2.5 MF-. We need to solve for the pH, and our connection between the two is the Ka expression. Step 2 is to write the Ka expression and rearrange for H3O+ concentration. The concentration of liquid water is not included in the Ka expression because it is a constant since it is a pure substance. Step 3 is to substitute [HA] and [A-] into the Ka expression. Step 4 is to use H3O+ concentration to calculate pH. Placing the H3O+ concentration into the pH equation gives the pH of the buffer. What is the pH of a buffer prepared with 0.50 M HNO2 and 0.75 M NO2-, if the Ka of nitrous acid is 4.5 × 10-4? Is it a, b, c or d? The correct answer is (c). The pH of the buffer is 3.52.
Our life depends on the pH of solutions and the existence of buffers. Without the existence of buffers, the pH of our blood would fluctuate with catastrophic consequences. In the absence of a buffer, adding a few drops of a strong acid to water at pH 7.0 will lower the pH significantly, in this case to pH 3.0. However, adding the same number of drops of strong acid to a buffer does not appreciably alter the pH. With the buffer, it only changes from pH 7.0 to 6.9. A buffer solution maintains the pH of a solution by neutralizing small amounts of added acid or base. In a buffer, an acid must be available to react with any OH- that is added, and a base must be available to react with any added H3O+ that is added. However, that acid and base cannot neutralize each other. Therefore, a combination of an acid-base conjugate pair is used in buffers. Most buffer solutions consist of nearly equal concentrations of a weak acid and a salt containing its conjugate base Buffers may also contain a weak base and the salt of the weak base, which contains its conjugate acid. Which of the following represents a buffer system? Is it a, b, c or d? The correct answer is (d). A buffer system consists of a weak acid and a salt containing its conjugate base. H2CO3 is a weak acid and NaHCO3 is the salt containing its conjugate base. A buffer can be made from the weak acid, acetic acid, and its salt, sodium acetate. As a weak acid, acetic acid dissociates slightly in water to form H3O+ and a very small amount of acetate ion, C2H3O2-. The addition of its salt, sodium acetate, provides a much larger concentration of acetate ion, which is necessary for its buffering capability. When a small amount of acid is added to the buffer, the additional H3O+ combines with the acetate ion, causing the equilibrium to shift in the direction of the reactants, acetic acid and water. There will be a slight decrease in the acetate ion and a slight increase in the acetic acid but both the H3O+ concentration and pH are maintained. If a small amount of base is added to this same buffer solution, it is neutralized by the acetic acid, which shifts the equilibrium in the direction of the products, acetate ion and water. The acetic acid decreases slightly, and the acetate ion increases slightly, but again the H3O+ concentration and pH of the solution are maintained. Consider the buffer system of H2PO4- and HPO42-(aq). If a small amount of H3O+ is added, it is neutralized by: a, b, c or d? The correct answer is (a). When a small amount of acid is added, the additional H3O+ combines with the hydrogen phosphate ion, HPO42- , causing the equilibrium to shift to the direction of the reactants. Let's look at a how we can calculate the pH of a buffer if we are given the concentrations of the weak acid and its conjugate base. The problem is what is the pH of a buffer prepared with 2.0 M HF and 2.5 MF-, if the Ka of hydrofluoric acid is 3.5×10-4? Step 1 is to state the given and needed quantities. In this problem, we are given 2.0 M HF and 2.5 MF-. We need to solve for the pH, and our connection between the two is the Ka expression. Step 2 is to write the Ka expression and rearrange for H3O+ concentration. The concentration of liquid water is not included in the Ka expression because it is a constant since it is a pure substance. Step 3 is to substitute [HA] and [A-] into the Ka expression. Step 4 is to use H3O+ concentration to calculate pH. Placing the H3O+ concentration into the pH equation gives the pH of the buffer. What is the pH of a buffer prepared with 0.50 M HNO2 and 0.75 M NO2-, if the Ka of nitrous acid is 4.5 × 10-4? Is it a, b, c or d? The correct answer is (c). The pH of the buffer is 3.52.