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Acids, Bases, and Ph Calculations in Chemistry, Ph Is a Measure of the Acidity Or Basicity of an Aqueous Solution

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Acids, Bases, and Ph Calculations in Chemistry, Ph Is a Measure of the Acidity Or Basicity of an Aqueous Solution ACIDS AND BASES Properties of acids 1. Acids have a sour taste. 2. Acids are corrosive. 3. Acids change the color of certain vegetable dyes, such as litmus, from blue to red. 4. Acids lose their acidity when they are combined with alkalies. The name "acid" comes from the Latin acidus, which means "sour," and refers to the sharp odor and sour taste of many acids. Examples: Vinegar tastes sour because it is a dilute solution of acetic acid in water. Lemon juice tastes sour because it contains citric acid. Milk turns sour when it spoils because lactic acid is formed, and the unpleasant, sour odor of rotten meat or butter can be attributed to compounds such as butyric acid that form when fa`t spoils. Acids, Bases, and pH Calculations In chemistry, pH is a measure of the acidity or basicity of an aqueous solution. Solutions with a pH less than 7 are said to be acidic and solutions with a pH greater than 7 are basic or alkaline. Pure water has a pH very close to 7. There are several ways to define acids and bases, but pH only refers to hydrogen ion concentration and is only meaningful when applied to aqueous (water-based) solutions. When water dissociates it yields a hydrogen ion and a hydroxide. + - H2O ↔ H + OH When calculating pH, remember that [] refers to molarity, M. + - -14 Kw = [H ][OH ] = 1x10 at 25°C for pure water [H+] = [OH-] = 1x10-7 Acidic Solution: [H+] > 1x10-7 Basic Solution: [H+] < 1x10-7 Calculate pH and [H+] + pH = - log10[H ] [H+] = 10-pH Example: Calculate the pH for a specific [H+].C4alculate pH given [H+]= 1.4 x 10-5 M + pH =- log10[H ] -5 pH = log10(1.4 x 10 ) pH = 4.85 Example: Calculate [H+] from a known pH. Find [H+] if pH = 8.5 [H+] = 10-pH [H+] = 10-8.5 [H+] = 3.2 x 10-9 M What is the pH of a solution with [H+] = 1 x 10-6 M. Solution pH is calculated by the formula pH = - log [H+] Substitute [H+] with the concentration in the question. pH = - log (1 x 10-6) pH = -(-6) pH = 6 Answer The pH of the solution is 6. What is the pH of a 0.025 M solution of Hydrobromic Acid? Solution Hydrobromic Acid or HBr, is a strong acid and will dissociate completely in water to H+ and Br-. For every mole of HBr, there will be 1 mole of H+, so the concentration of H+ will be the same as the concentration of HBr. Therefore, [H+] = 0.025 M. pH is calculated by the formula pH = - log [H+] Enter the concentration found before pH = - log (0.025) pH = -(-1.602) pH = 1.602 Answer The pH of a 0.025 M solution of Hydrobromic Acid is 1.602. Calculating pH of a strong base. Question What is the pH of a 0.05 M solution of Potassium Hydroxide? Solution Potassium Hydroxide or KOH, is a strong base and will dissociate completely in water to K+ and OH-. For every mole of KOH, there will be 1 mole of OH-, so the concentration of OH- will be the same as the concentration of KOH. Therefore, [OH-] = 0.05 M. Since the concentration of OH- is known, the pOH value is more useful. pOH is calculated by the formula pOH = - log [OH-] Enter the concentration found before pOH = - log (0.05) pOH = -(-1.3) pOH = 1.3 The value for pH is needed and the relationship between pH and pOH is given by pH + pOH = 14 pH = 14 - pOH pH = 14 - 1.3 pH = 12.7 Answer The pH of a 0.05 M solution of Potassium Hydroxide is 12.7. Exercises CALCULATING THE pH AND pOH OF STRONG ACIDS AND BASES • This is relatively easy because the species have completely dissociated • Only needs to know the original concentration of the acid or base Example 1 Calculate the pH of 0.1M hydrochloric acid. HCl (a strong monoprotic acid) is fully dissociated. HCl ——> H+ (aq) + Cl¯(aq) The [ H+] is therefore the same as the original concentration of HCl i.e. 0.1M. pH = - log10 [H+] = - log10 (10-1) = 1 ANS. 1 Example 2 Calculate the pH of 0.001M sodium hydroxide. Sodium hydroxide (a strong base) is fully dissociated. Na+OH¯ ——> Na+ (aq) + OH¯(aq) The [OH¯] is therefore the same as the original concentration of NaOH i.e. 0.001M. pOH = - log10 [OH¯] = - log10 (10-3) = 3 and pH = 14 - pOH = 14 - 3 = 11 ANS. 11 The dissociation constant for a weak acid (Ka) A weak monobasic acid (HA) dissociates in) H+ Acid-base A4 5 Q.4 Calculate the pH and pOH of the following solutions. a) HCl; 0.1M, 0.5M b) H2SO4; 0.1M, 0.5M c) KOH; 0.1M d) NaOH; 2M, 0.0005M e) The solution remaining when 30 cm3 of 0.100M NaOH has been added to 20 cm3 of 0.200M HCl f) The solution remaining when 24.9 cm3 of 0.100M NaOH has been added to 25 cm3 of 0.100M HCl Acid strength The strength of an acid refers to its ability or tendency to lose a proton (H+). A strong acid is one that completely ionizes (dissociates) in a solution. In water, one mole of a strong acid HA dissolves yielding one mole + + of H (as hydronium ion H3O ) and one mole of the conjugate base, A−. Essentially none of the non-ionized acid HA remains. Examples of strong acids are hydrochloric acid (HCl), hydroiodic acid (HI), hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3) and sulfuric acid (H2SO4). In aqueous solution each of these essentially ionizes 100%. In contrast, a weak acid only partially dissociates. Examples in water include carbonic acid (H2CO3) and acetic acid (CH3COOH). At equilibrium both the acid and the conjugate base are present in solution. Stronger acids have a larger acid dissociation constant (Ka) and a smaller logarithmic constant (pKa = - log Ka) than weaker acids. The stronger an acid is, the more easily it loses a proton, H+. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base. While Ka measures the strength of an acidic molecule, the strength of an aqueous acid solution is measured by pH, which is a function of the concentration of hydronium ion in solution. The pH of a simple solution of an acid in water is determined by both Ka and the acid concentration. For weak acid solutions it depends on the degree of dissociation, which may be determined by an equilibrium calculation. For concentrated solution of strong acids with pH less than about zero, the Hammett acidity function is a better measure of acidity than the pH. Sulfonic acids, which are organic oxyacids, are a class of strong acids. A common example is p-toluenesulfonic acid (tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable. Superacids are acid solutions which are more acidic than 100% sulfuric acid.[1] Examples of superacids are fluoroantimonic acid, magic acid and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations. Strong acids A strong acid is an acid that ionizes completely in an aqueous solution by losing one proton, according to the equation HA(aq) → H+(aq) + A−(aq) For sulfuric acid which is diprotic, the "strong acid" designation refers only to dissociation of the first proton + H2SO4(aq) → H (aq) + − HSO4 (aq) More precisely, the acid must be stronger in aqueous solution than hydronium ion, so strong acids are acids with a pKa < −1.74. An example is HCl for [2] which pKa = -6.3. This generally means that in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution. In all other acid-water reactions, dissociation is not complete, so will be represented as an equilibrium, not a completed reaction. The typical definition of a weak acid is any acid that does not dissociate completely. The difference separating the acid dissociation constants of strong acids from all other acids is so small that this is a reasonable demarcation. Due to the complete dissociation of strong acids in aqueous solution, the concentration of hydronium ions in the water is equal to the total concentration (ionized and un-ionized) of the acid introduced to solution: [H+] = − [A ] = [HA]total and pH = −log[H+]. Determining acid strength The strength of an acid, in comparison to other acids, can be determined without the use of pH calculations by observing the following characteristics: 1. Electronegativity: The higher the electronegativity of a conjugate base in the same period, the more acidic. In other words, the more electronegative A- is, more acidic (where HA → H+ + A−). 2. Atomic Radius: With increasing atomic radius, acidity also increases. For example, HCl and HI, both strong acids, ionize 100% in water to become their respective ionic constituents. However, HI is stronger than HCl. This is because the atomic radius of an atom of iodine is much larger than that of a chlorine atom. As a result, the negative charge over the I− anion is dispersed over a larger electron cloud and its attraction for the proton (H+) is not as strong as the same attraction in HCl.
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