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6.2 – Acid & Base Strength

6.2 – Acid & Base Strength. Unit 6 – Acids and Bases. Polyprotic Acids. We have seen examples of acids that contain more than one hydrogen ion that can be lost. Sulfuric acid, H 2 SO 4 , for example, has two hydrogen ions that it can give up. The first hydrogen ion is released as:

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6.2 – Acid & Base Strength

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  1. 6.2 – Acid & Base Strength Unit 6 – Acids and Bases

  2. Polyprotic Acids • We have seen examples of acids that contain more than one hydrogen ion that can be lost. • Sulfuric acid, H2SO4, for example, has two hydrogen ions that it can give up. The first hydrogen ion is released as: H2SO4 (aq) → H+(aq) + HSO4-(aq) • The second hydrogen will be more difficult to remove because it must now be removed from a negative ion, HSO4-. HSO4-(aq) → H+(aq) + SO42-(aq)

  3. Polyprotic Acids • For acids that can donate more than one hydrogen, it will always be easier to donate the first H+ than the second. • If there are more hydrogens to release, each H+ is more difficult to remove because of the increasingly NEGATIVE charge of the rest of the molecule. • Acids that can donate more than one hydrogen ion are called polyprotic

  4. Polyprotic Acids • How many H+ can citric acid, H3C6H5O7, release? • Three (3) • What do you think the equations describing the release of the H+ will be? Make your predictions before you check . . . • The reactions would be: • 1.    H3C6H5O7 (aq) → H+(aq) + H2C6H5O7-(aq) • 2.    H2C6H5O7-(aq) → H+(aq) + HC6H5O72-(aq) • 3. HC6H5O72- H+(aq) + C6H5O73-(aq)

  5. Strong & Weak Acids & Bases • You have undoubtedly heard of the pH scale before and know that it has something to do with indicating how strong or weak an acid is. In this part of the unit we will work towards defining acid and base strength in terms of pH, but there are several important steps along the way.

  6. Acids and Bases and Electricity • When ions break apart from each other in solutions they can form electrolytic solutions which can conduct electricity. • Solutions that conduct electricity well are called strong electrolytes – these are good conductors because the ions break up well and many ions are present in the solution • Weak electrolytes are solutions that do not conduct electricity well because there are few ions in the solution. • Strong electrolytes conduct electricity well and tend to correspond to strong acids and strong bases • Weak electrolytes conduct electricity poorly and tend to correspond to weak acids and weak bases

  7. Strong & Weak Acids & Bases • These terms apply equally well to acids and bases which are, of course, electrolytic solutions:

  8. Strong Acids & Bases • Strong acids and bases are essentially one-way reactions - the acid or base breaks down completely to produce ions. • At equilibrium there are very few reactants left (very low concentration); only products - the ions. • Examples: • Notice how the arrow is one sided, indicating that the products are favored. It is still important to remember that strong acids are still considered to in equilibrium despite the arrow. Confusing and annoying, I know

  9. Weak Acids & Bases • Weak acids and bases, however, do not ionize completely. • For weak electrolytes, equilibrium lies to the left side of the equation (the reactant side) and there will be few ions present. • The double arrow is commonly used to indicate the partial ionization of the solution. Some examples:

  10. Acid & Base Strength vs. Dilution • It is important that you don't confuse acid and base strength will dilution. • Dilute and concentrated, you should remember, refer to the relative amounts of solute and solvent in a solution. By contrast, acid and base strength specially refer to the concentration of ions in the solution. • A strong acid such as HCl is still a strong acid (completely ionizes to produce many ions) even when it is dilute (lots of water and relatively small amounts of HCl). Acetic acid is still a weak acid even when it is concentrated.

  11. Equilibrium Constants for Acids & Bases • Because acid/base solutions are systems at equilibrium, we can write equilibrium constant expressions for these systems. • Notice that we now identify the equilibrium constant as Ka for acids and Kb for bases.

  12. Equilibrium Constants for Acids & Bases - Examples • 1) HCl(aq) → H+(aq) + Cl-(aq) Ka = [H+] [Cl-] = 1.3×106 [HCl] • 2) HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq) Ka = [H3O+] [Cl-] = 1.3×106 [HCl] • 3) HCHO2 (aq) ↔ H+(aq) + CHO2-(aq) Ka = [H+] [CHO2-] = 1.8 × 10-4 [HCHO2] • 4) NH3 (aq) + H2O(l) ↔ NH4+(aq) + OH-(aq) • Kb = [NH4+] [OH-] = 1.8×10-5 [NH3]

  13. Measuring Acid & Base Strengths • Recall that the value of the equilibrium constant indicates which side of a reaction is favored. • When Keq is large the product side is favored; • When Keq is small the reactants are favored. • So we now have a numerical way to indicate acid and base strength - Ka and Kb : • The Table of Acid and Base Strengths gives Ka and Kb values for a number of acids and bases. Ka values for very strong acids are typically not given as there is little use in expressing them numerically.

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