Acids & Bases

1. Acids & Bases Chapter 15 & 16

2. Acids • Have a sour taste • Affect indicators • React with bases to produce salt & water • Conduct an electric current • Examples include salad dressing, rainwater, car batteries, & soft drinks

3. Bases • Taste bitter • Feel slippery • Affect indicators • React with acids to produce salt & water • Conduct an electric current • Examples include soap, baking soda, & drain cleaners

4. Solutions of acids & bases • The amounts of hydrogen ions (H+) & hydroxide ions (OH-) determine whether a solution is acidic, basic, or neutral. • Acidic solutions contain more H+ than OH- • Basic solutions contain more OH- than H+ • Neutral solutions contain equal amounts

5. Arrhenius model • In 1883, Svante Arrhenius proposed the Arrhenius model of acids & bases to explain how pure water could become acidic or basic.

6. Arrhenius model • According to the Arrhenius model • An acid is a substance that contains hydrogen & ionizes to produce H+ in solution HClO4 H+ + ClO4- • A base is a substance that contains a hydroxide group & ionizes to produce OH- in solution NaOH  Na+ + OH-

7. Bronsted-Lowry Model • Johannes Bronsted & Thomas Lowry proposed a better model for acids & bases that focuses on the role of the proton, or H+. • According to the Bronsted-Lowry model: • An acid donates a proton, or H+ • A base accepts a proton, or H+

8. Explanation When an acid, HX, dissolves in water, it donates a H+ to a water molecule. The water acts as a base & accepts the H+. HX + H2O H3O+ + X- On accepting the H+, the water becomes an acid, H3O+. The hydronium is an acid because it has an extra H+ it can donate. On donating its H+, the acid becomes a base, X-. The reaction can also occur in the reverse direction.

9. Conjugates • The products of an acid-base reaction are known as the conjugate acid-base pair. • The conjugate acid is the species produced when a base accepts a H+ from an acid. • The conjugate base is the species that results when an acid donates an H+ to a base.

10. ProblemLabel the acid, base, conjugate acid, & conjugate base in each of the reactions.HF + H2O H3O+ + F-NH3 + H2O NH4+ + OH-

11. Amphoteric:substances, like water, that can act as acids or bases

12. Strong Acids • Strong acids are acids that ionize completely in solution. HCl + H2O  H3O+ + Cl- • Strong acids: HCl HBr HI HClO4 HNO3 H2SO4

13. Weak Acids • Weak acids are acids that ionize only partially in solution. • All other acids are weak acids HC2H3O2 + H2O H3O+ + C2H3O2- • In the reaction, the acetic acid does not ionize completely & some of the acetic acid ions remain in molecular form in the solution.

14. Cont’d… • The double arrow in the reaction means the reaction is an equilibrium reaction. • The equilibrium constant expression is written by placing the concentration of the product in the numerator & the concentration of the reactants in the denominator. • Water is a constant so it is not added to the expression. • Ka is the acid ionization constant. • The weaker the acid, the smaller the Ka. • The stronger the acid, the larger the Ka.

15. Strong & Weak Bases • The same application that apply for acids, apply for bases except using OH-. • Strong bases ionize completely in solution • Weak bases ionize only partially in solution. • The equilibrium equation for bases use the base ionization constant, Kb. • The strong bases are: NaOH KOH Ca(OH)2 RbOH CsOH Ba(OH)2

16. Water • Water always ionizes itself a little. H2O H+ + OH- • This gives us a special constant that applies to the self-ionization of water. Kw= [H+][OH-] Kw is the ion product constant for water.

17. Cont’d… • In pure water at 298K, [H+] & [OH-] always equal 1.0x10-7 M. • Therefore, Kw= (1.0x10-7)(1.0x10-7) Kw= 1.0x10-14 • You can use Kw to calculate the concentration of either ion if you know the concentration of the other.

18. ProblemUsing the ion product constant for water, calculate [OH-] if [H+]= 3.0x10-2 M

19. pH scale • Chemists use a pH scale to express hydrogen ion concentrations. pH= -log [H+] • pH means the power of hydrogen • pH ranges from 0 to 14 • 0 to 7 is acidic • 7 to 14 is basic • 7 is neutral • pH decreases as [H+] increases

21. pOH scale • Chemists can also use the pOH scale to express the basicity, or alkalinity, of a solution. pOH= -log [OH-] • pOH ranges from 0 to 14 • 0 to 7 is basic • 7 to 14 is acidic • 7 is neutral

22. A chemist can calculate pH or pOH if the other is known using:pH + pOH = 14.00

23. Problems Calculate pH. • [H+]= 1.0x10-2 M • [H+]= 3.0x10-6 M • [OH-]= 8.2x10-6 M • [OH-]= 4.3x10-4 M

24. Neutralization • Occurs when an acid and a base react to form a salt & water HCl + NaOH NaCl + H2O

25. Titration • The process of determining the concentration of an unknown solution by dripping, or titrating, a known strong acid or base into it • If the unknown solution is acidic, we titrate the solution with a strong base of a known concentration.

26. Steps to Titration • An indicator is a chemical dye added to the solution whose color is affected by acidic or basic solutions.

27. Cont’d… • As you add the acid or base to the unknown solution, pH slowly rises. • When the acid is neutralized, pH rises rapidly & the indicator in the solution changes color. • The color change is called the end point.