Chapter 14 – Acids, Bases and pH Section 14.2 – Strengths of Acids and Bases

Acids and bases are classified into 2 categories based upon their strength- the degree to which they form ions Strong category- Substance that completely dissociates or ionizes to produce the maximum number of ions when dissolved in H 2 O All other acids and bases are classified as weak- produce few ions when dissolved in water Acid/Base Classification

See table on pg 498 in book. If an acid or base is not listed in this group, it is considered weak The terms strong and weak are not absolute. The strengths of acids and bases cover a wide range The strongest bases are hydroxides in Group 1 (alkali metals) and Group 2 (alkaline earth metals) Acid/Base Classification

All the formula units dissociate into separate ions- Dissociation is complete Examples: NaOH, KOH (group 1 hydroxides) NaOH Na + + OH - STRONG BASE

Completely ionizes in water and no molecules exist in the water solution Examples: HCl, H 2 SO 4 HCl + H 2 O → H 3 O + + Cl - STRONG ACID

In general, the weaker the H-A bond, the stronger the acid Ex. the highly electronegative fluorine creates a weak acid with hydrogen (HF, hydrofluoric acid) because it has such a strong attraction for its shared electrons. Strong Acid (cont)

Acid in which almost all the molecules remain as molecules when placed into a solution Examples: Acetic acid (HC 2 H 3 O 2 ), Phosphoric acid (H 3 PO 4 ), Carbonic acid (H 3 CO 3 ) The molecular structure of a weak acid determines the extent to which the acid ionizes in water WEAK ACID

Base in which most of the molecules do not react with water to form ions Examples: ammonia (NH 3 ), aluminum hydroxide (Al(OH) 3 ), WEAK BASE

 weak does not mean unimportant!  due to the wide variety of structures possible, there is a wide range of "weakness" to this group  they are considered to be weak because there is only partial ionization of these acids and bases into hydronium or hydroxide ions Weak Acids & Bases

Weak and Strong- strengths of acids and bases Dilute and Concentrated- concentrations of solutions Both strength and concentration determine the behavior of the solution Strength and concentration

0.100 M H 2 SO 4 or 1.00 M HF  Which is more concentrated?  Which is more acidic? (The stronger acid?) Strength v. Concentration

Neutral ↓ ← more acidic −− −− more basic → pH ↓ = H + ↑ = OH + ↓ The pH scale

 means "power of hydrogen"  measures the concentration of H + (H 3 O + ) ions, [H + ] in solution  used to make the wide range of possible concentrations easier to work with  pH is a mathematical scale in which the concentration of hydronium ions in a solution are expressed as a number from 0 to 14 (a logarithmic scale of 10 0 to ) The pH scale

The value of pH is the negative of the exponent, ie: 0 to 14 ex.If [H + ] is M, the pH is 6 If pH = 12, then [H + ] = M The pH scale (cont)

pH can be used to determine pOH ("power of hydroxide") ex.If pH = 12, then pOH = 2 **pOH will be the difference between 12 and 14 (the high end of the pH scale) : pH + pOH = 14** If pH = 12, then [OH - ] = M The pH scale (cont)

the pH value can be calculated on calculator using the logarithmic function pH= -log[H + ] [H + ] = hydrogen ion concentration pOH =-log[OH] [OH] = hydroxide ion concentration pH + pOH = 14 (at 25ºC) pH scale

 #1 To find the pH you must take the –log[H+]  #2 To find the pOH you must take the –log[OH-]  #3 To find the [H+] you must take the anti-log of the pH  4# To find the [OH-] you must take the anti-log of the pOH 

1) hydronium ion = M pH= -log[10 -5 ] = 5 pOH = 9 2) hydroxide ion = M pOH= -log[10 -4 ] = 4 pOH = 4 pH Practice problems

3) hydroxide ion =.01 M pH= -log[ ] = 2 pOH = 12 4) hydroxide ion =.056 M pH= -log[ ] = 1.25 pOH = = ) hydroxide ion =.0076 M pH= -log[ ] = 2.11 pOH = = pH Practice problems

Indicators register different colors at different pHs (indicator paper) pH meters measure exact pH pH greater than 7 is basic pH = 7, solution is neutral (not acidic or basic) pH less than 7 is acidic Measuring pH

 As the pH increases above 7, the concentration of hydroxide ion increases, the concentration of hydronium ions decreases pH

Every drop in pH is 10 times increase in [H + ] ex.dropping from pH 5 to pH 4 means the [H + ] of pH 4 is 10 times greater ex: pH = 1 pH=5 5-1 = 4 pH units 10 4 = 10,000 pH (cont)

Every one unit increase in the pH = factor of 10 increase in the OH - concentration Ex: a pH of 11 has 10 times the concentration of hydroxide ions than a pH of 10 In a neutral solution, the concentration of OH - and the concentration of H 3 O + are equal

Lemon juice: pH= ~ 2.3 Milk: pH= ~ 6.4 Ammonia: pH= ~ 11.9 Drain cleaner: pH= ~ 14 pH of common substances

Colored solutions (indicators) have different colors at different pHs – give approximate pH by comparing the color to a standard chart Used in pools, hot tubs, etc. See Figure (p. 504) for acid-base indicator ranges and colors Acid-base indicator ranges

 Occurs when an acid reacts with a base to form a salt and water; salts are crystalline solids that may or may not be soluble in water; Review the solubility guidelines from CH 13 Ex:H 2 SO 4 + 2NaOH Na 2 SO 4 + 2H 2 O Acid + Base Salt + Water  Salts contain the positive ion of a base, and the negative ion of an acid.  Some salts, called acid salts, contain hydrogen, ie: NaHSO 4 or K 2 HPO 4 NEUTRALIZATION REACTION