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DEFINITIONS OF ACIDS & BASES

DEFINITIONS OF ACIDS & BASES. ARRHENIUS ACID An Arrhenius acid is any substance that provides hydrogen ions, H + , when dissolved in water. ARRHENIUS BASE An Arrhenius base is any substance that provides hydroxide ions, OH - , when dissolved in water.

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DEFINITIONS OF ACIDS & BASES

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  1. DEFINITIONS OF ACIDS & BASES • ARRHENIUS ACID • An Arrhenius acid is any substance that provides hydrogen ions, H+, when dissolved in water. • ARRHENIUS BASE • An Arrhenius base is any substance that provides hydroxide ions, OH-, when dissolved in water. • EXAMPLES OF ANARRHENIUS ACID AND BASE • HNO3 is an acid: HNO3(aq) H+ (aq) + NO3- (aq) • KOH is a base: KOH(aq) K+ (aq) + OH- (aq)

  2. BRØNSTED ACIDS & BASES • BRØNSTED ACID • A Brønsted acid is any hydrogen-containing substance that is capable of donating a proton (H+) to another substance. • BRØNSTED BASE • A Brønsted base is any substance capable of accepting a proton from another substance. • EXAMPLE OF ABRØNSTED ACID AND BASE • HC2H3O2(aq) + H2O(l) H3O+ (aq) + C2H3O2-(aq) • In this reaction, HC2H3O2 behaves as a Brønsted acid by donating a proton to the H2O. The H2O behaves as a Brønsted base by accepting the proton.

  3. BRØNSTED ACIDS & BASES • NH3(aq) + H2O(l) NH4+ (aq) + OH-(aq) • In this reaction, NH3 behaves as a Brønsted Base by accepting a proton from the H2O. The H2O behaves as a Brønsted acid by donating a proton.

  4. EXAMPLES OF ACIDS Memorize the following acids by name and formula HCl Hydrochloric Acid H2SO4 Sulfuric Acid HNO3 Nitric Acid H3PO4 Phosphoric Acid HC2H3O2 Acetic Acid (an organic acid) H2CO3 Carbonic Acid (carbonated water) [salts of bases – later]

  5. EXAMPLES OF ACIDS A strong acid or basedissociates100%, while a weak or moderately weakone dissociates less than 100%.

  6. ACID DISSOCIATION CONSTANTS • An acid dissociation constant is the equilibrium constant for the dissociation of a weak acid. It is represented by the symbol Ka. [similar to an equilibrium constant covered in Chapter 8] • The dissociation of a weak acid in solution is represented by the following equation in which HB represents the weak acid, and B- is the conjugate base of the acid. • The equilibrium expression for this reaction is: HB (aq) + H2O (l) ⇆ H3O+ (aq) + B− (aq)

  7. EXAMPLES OF BASES Memorize the following bases by name and formula Metal hydroxides, e.g., NaOH (lye), Ca(OH)2, Mg(OH)2 (milk of magnesia) NH3 Ammonia HCO3- Bicarbonate compounds (NaHCO3 – is common baking soda) Strong Bases – NaOH, KOH, Ca(OH)2 Weak Bases – Mg(OH)2, NH3, HCO3-[note: your author describes Mg(OH)2 as a strong base] [salts of acids – later]

  8. PROPERTIES OF ACIDS and BASES • All acids • Taste sour • Turn blue litmus red • Produce H3O+ ions when dissolved in water (or donate Hydrogen ions) • React with oxides and hydroxides to form water and a salt (the reaction with hydroxides is the classic neutralization reaction) • React with carbonates, and bicarbonates to form a salt, water and carbon dioxide • Non-metallic oxides plus water form oxyacids (e.g., SO2 acid rain, CO2) • React with active metals to form hydrogen gas and a salt.

  9. PROPERTIES OF ACIDS and BASES • All bases • Taste bitter • Turn red litmus blue • Feel slippery (due to their caustic nature) • Produce OH- when added to water (or accept hydrogen ions) • Can be formed by a reaction between a metallic oxide and water (e.g., K2O, CaO) • React with acids to form water and a salt (this is the classic neutralization reaction). • Bases also react with fats and oils and convert them into smaller, soluble molecules (soap – related to 3rd bullet above).

  10. PROPERTIES OF ACIDS (continued) • Acids can react with and dissolve active metals to yield hydrogen gas in a redox reaction. • The activity series is a tabular representation of the tendencies of metals to react with H+.

  11. THE SELF-IONIZATION OF WATER • A sample of absolutely pure water does not contain only H2O molecules. In addition, small but equal amounts of H3O+ and OH- ions are also present. • The reason for this is that in one liter of pure water 1.0 x 10-7 moles of water molecules behave as Brønsted acids and donate protons to another 1.0 x 10-7 moles of water molecules, which act as Brønsted bases. The reaction is: • As a result, absolutely pure water contains 1.0 x 10-7 mol/L of both H3O+ and OH-. • The term neutral is used to describe any water solution in which the concentrations of H3O+ and OH- are equal. • Thus, pure water is neutral because each of the ions is present at a concentration of 1.0 x 10-7 M. • Define amphoteric H2O (l) + H2O (l) ⇆ H3O+ (aq) + OH− (aq)

  12. THE ION PRODUCT OF WATER(continued) This gives rise to an expression called the ion product of waterand is termed Kw. Because the molar concentration of both H3O+ and OH- in pure water is 1.0 x 10-7, the numerical value for Kw can be calculated:

  13. THE ION PRODUCT OF WATER(continued) • Even though this equilibrium equation was derived on the basis of pure water, it is true for any solution in which water is the solvent. • ACIDIC SOLUTION • An acidic solution is a solution in which the concentration of H3O+ is greater than the concentration of OH-. It is also a solution in which the pH is less than 7. • BASIC OR ALKALINE SOLUTION • A basic or alkaline solution is a solution in which the concentration of OH- is greater than the concentration of H3O+. It is also a solution in which the pH is greater than 7.

  14. EXAMPLE OF ACID-BASE CALCULATION • Calculate the molar concentration of OH- in a solution that has an H3O+ concentration of 1.0 x 10-5 M. Classify the solution as acidic or basic. • Solution: The molar concentration of H3O+ will be substituted into the equilibrium expression for water, the resulting equation will be solved for [OH-]: • The molar concentration of OH- is seen to be smaller than the molar concentration of H3O+, so the solution is classified as being acidic.

  15. THE pH CONCEPT • It is often the practice to express the concentration of H3O+ in an abbreviated form called the pH rather than to use scientific notation. • It is also a common practice to represent the H3O+ ion by the simpler H+ ion. • The pH notation is defined below, using H+ in place of H3O+: pH = -log[H+], or in alternate form [H+]= 1x10-pH • Thus, the pH is seen to be the negative of the exponent used to express the molar concentration of H+ using scientific notation.

  16. EXAMPLES OF pH CALCULATIONS • Example 1: Calculate the pH of a solution in which [H+]= 1.0x10-9 M. • Solution: Because the pH is the negative of the exponent on 10 used to express [H+] using scientific notation, pH = -log (1.0x10-9) = -(-9) = 9.00.

  17. EXAMPLES OF pH CALCULATIONS (continued) • Example 2: Calculate the [OH-] for a solution with a pH = 4.0. • Solution: Because pH is the negative of the exponent on 10 used to express [H+] in scientific notation, the exponent must be -4. Then, [H+]= 1.0 x 10-4. This value is substituted into the equilibrium expression for water, and the equation is solved for [OH-]:

  18. Hydrolysis of salts • When an simple acid (or base) is added to water, the solution becomes acidic (or basic). • But when a salt is added to water the solution could be neutral, acidic or basic depending on the nature of the salt. Why? • This is related to • Conjugate acids and bases • Hydrolysis of salts • Buffers [the next three topics]

  19. PURE WATER vs. SODIUM ACETATE • Samples of pure water (left) and sodium acetate dissolved in water (right) behave differently when phenolphthalein indicator is added. The acetate ion hydrolyzes in water to form a basic solution that turns phenolphthalein to a pink color.

  20. CONJUGATE ACIDS & BASES • CONJUGATE ACIDS AND BASES • The base formed (NO2-) when a substance (HNO2) acts as a Brønsted acid is called the conjugate base of the acid. Similarly, the acid formed (H3O+) when a substance (H2O) acts as a Brønsted base is called the conjugate acid of the base. • CONJUGATE ACID-BASE PAIRS • A Brønsted acid (such as HNO2) and its conjugate base (NO2-) form what is called a conjugate acid-base pair. • The same name is given to a Brønsted base (such as H2O) and its conjugate acid (H3O+).

  21. MONOPROTIC, DIPROTIC & TRIPROTIC ACIDS • Monoprotic acids give up only one proton per molecule when dissolved in water. • Diprotic acids give up a maximum of two protons per molecule when dissolved in water. • Triprotic acids give up a maximum of three protons per molecule when dissolved in water.

  22. HYDROLYSIS REACTIONS OF SALTS • Salts consist of the cation of a base and the anion of an acid. The cation of a base is the conjugate acid of the base from which it came. Similarly, the anion of an acid is the conjugate base of the acid from which it came. • The strength of a conjugate acid or base depends upon the strength of the base or acid from which they came. The stronger an acid is, the weaker its conjugate base is. Similarly, the stronger a base is, the weaker its conjugate acid is. • The pH of a water solution of a salt depends on the strength of the salt cation as an acid and the strength of the salt anion as a base.

  23. HYDROLYSIS REACTIONS OF SALTS(continued) • Example 1: A solution containing the dissolved salt NaCl has a pH the same as the water used as a solvent for the solution. • This is because the Na+ ion is the conjugate acid of the strong base NaOH and is a very weak acid. • Similarly, the Cl- ion is the conjugate base of the strong acid HCl and is a very weak base. • Neither the Na+ cation nor the Cl- anion will react appreciably with water to produce OH- or H+.

  24. HYDROLYSIS REACTIONS OF SALTS(continued) • Example 2: A solution containing the dissolved salt sodium carbonate, Na2CO3, has a pH significantly higher than that of the water used as a solvent for the solution. • The Na+ ion is a weak acid as was discussed on the previous slide. • The CO32- ion is the conjugate base of the weak acid HCO3- and as a result is a significant base that will react with water as follows: • This reaction, called a salt hydrolysis reaction, is seen to produce OH- ions which causes the pH to be higher than water and the solution is basic. CO32− (aq) + H2O (l) ⇆ HCO3− (aq) + OH− (aq)

  25. BUFFERS • Buffers are solutions with the ability to resist changing pH when acids (H+) or bases (OH-) are added to them. • Many useful buffers consist of a solution containing a mixture of a weak acid and a salt of the acid (e.g. acetic acid and sodium acetate) or in general the conjugates of a weak acid/base pair. • Any added acid (H+ ions) react with the anion from the salt, which also happens to be the conjugate base of the weak acid. • Any added base (OH- ions) react with the nonionized weak acid. • The buffer capacity is the amount of acid (H+) or base (OH-) that can be absorbed by a buffer without causing a significant change in pH. C2H3O2− (aq) + H+ (aq) ⇌ HC2H3O2 (aq) HC2H3O2 (aq) + OH− (aq) ⇌ C2H3O2− (aq) + H2O (l)

  26. UNBUFFERED vs. BUFFERED SOLUTIONS The solution on the left is not buffered; the one on the right is; universal indicator has been added to each solution. Sodium hydroxide has been added to each solution Hydrochloric acid has been added to two fresh samples that originally looked like the first pair of samples.

  27. Common buffers • Acetic acid/acetate • Bicarbonate/carbonate • Carbonic acid(aq. CO2)/bicarbonate (a common buffer in blood) • Dihydrogen phosphate/monohydrogen phosphate Explain/show how a buffer can resist pH change using acid/base theory and Le Chatelier’s principle.

  28. The formal Chapter 9 lecture will stop about here.

  29. ANALYZING ACIDS AND BASES • The analysis of acid solutions to determine the amount of acid they contain is an important procedure done in many laboratories. • An acid-base titration is one commonly-used method of analysis. • When a titration is done, an accurately-measured volume of acid is put into a flask using a pipet. • A few drops of indicator solution is added, then a base solution of known concentration is carefully added from a buret until all the acid has been reacted (equivalence point). • The point at which all the acid has reacted is shown by a color change (endpoint) in the indicator. • The concentration of the base and the volume required in the titration allow the concentration of acid to be determined.

  30. TITRATION TECHNIQUE

  31. TITRATION CALCULATIONS • Titration calculations are dependent upon knowledge of two things: the stoichiometry of the reaction that occurs between the acid and base, and the equation defining molarity. • An example of a reaction equation is: H2SO4(aq) + 2NaOH(aq) →Na2SO4(aq) + 2H2O(l) • Such an equation provides the relationship between the number of moles of acid and base that react. In this reaction it is seen that 1 mole of H2SO4 acid reacts with 2 moles of NaOH base. • The molarity equation may be rearranged to allow the calculation of the number of moles of solute contained in a specific volume of solution or the volume of solution that contains a specific number of moles of solute.

  32. TITRATION CALCULATIONS M x liters of solution = moles of solute

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