Presentation on theme: "Chapter 19 Acids and Bases. Properties of Acids n Taste sour (don’t try this at home). n Conduct electricity. –Some are strong, others are weak electrolytes."— Presentation transcript:
Properties of Acids n Taste sour (don’t try this at home). n Conduct electricity. –Some are strong, others are weak electrolytes.
Properties of Acids n React with metals to form hydrogen gas. n Change indicators (blue litmus to red). n React with hydroxides to form water and a salt.
Properties of Bases n React with acids to form water and a salt. n Taste bitter. n Feel slippery (don’t try this either).
Names and Formulas of Acids n An acid is a chemical that produces hydrogen ions (H 1+ ) when dissolved in water n Thus, general formula = HX, where X is a monatomic or polyatomic anion
Names and Formulas of Acids n HCl (g) named hydrogen chloride n HCl (aq) is named as an acid n Name focuses on the anion present
Names and Formulas of Acids 1. When anion ends with -ide, the acid starts with hydro-, and the stem of the anion has the suffix -ic followed by the word acid
Names and Formulas of Acids 2. When anion ends with -ite, the anion has the suffix -ous, then acid 3. When anion ends with -ate, the anion suffix is -ic and then acid n Table 19.1, page 588 for examples
Names and Formulas of Bases n A base produces hydroxide ions (OH 1- ) when dissolved in water. n Named the same way as any other ionic compound –name the cation, followed by anion
Names and Formulas of Bases n To write the formula: write symbols; write charges; then cross (if needed)
Svante Arrhenius n Swedish chemist (1859-1927) - Nobel prize winner in chemistry (1903) n one of the first chemists to explain the chemical theory of the behavior of acids and bases n Dr. Hubert Alyea-last graduate student of Arrhenius. (link below) http://www.woodrow.org/teachers/ci/1992/Arrhenius.html
1. Arrhenius Definition n Acids produce hydrogen ions (H 1+ ) in aqueous solution. n Bases produce hydroxide ions (OH 1- ) when dissolved in water. n Limited to aqueous solutions. n Only one kind of base (hydroxides) n NH 3 (ammonia) could not be an Arrhenius base.
Polyprotic Acids n Some compounds have more than 1 ionizable hydrogen. n HNO 3 nitric acid - monoprotic n H 2 SO 4 sulfuric acid - diprotic - 2 H + n H 3 PO 4 phosphoric acid - triprotic - 3 H + n Having more than one ionizable hydrogen does not mean stronger!
Polyprotic Acids n However, not all compounds that have hydrogen are acids n Also, not all the hydrogen in an acid may be released as ions –only those that have very polar bonds are ionizable - this is when the hydrogen is joined to a very electronegative element
Arrhenius examples... n Consider HCl n What about CH 4 (methane)? n CH 3 COOH (ethanoic acid, or acetic acid) - it has 4 hydrogens like methane does…? n Table 20.4, p. 595 for bases
Johannes Bronsted / Thomas Lowry (1879-1947) (1874-1936)
2. Brønsted-Lowry Definitions n Broader definition than Arrhenius n Acid is hydrogen-ion donor (H + or proton); base is hydrogen-ion acceptor. n Acids and bases always come in pairs. n HCl is an acid. –When it dissolves in water, it gives it’s proton to water. n HCl(g) + H 2 O(l) H 3 O + + Cl - n Water is a base; makes hydronium ion.
Acids and bases come in pairs... n A conjugate base is the remainder of the original acid, after it donates it’s hydrogen ion n A conjugate acid is the particle formed when the original base gains a hydrogen ion n Indicators are weak acids or bases that have a different color from their original acid and base
Acids and bases come in pairs... n General equation is: n HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) n Acid + Base Conjugate acid + Conjugate base n NH 3 + H 2 O NH 4 1+ + OH 1- base acid c.a. c.b. base acid c.a. c.b. n HCl + H 2 O H 3 O 1+ + Cl 1- n acid base c.a. c.b. n Amphoteric - acts as acid or base
3. Lewis Acids and Bases n Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction n Lewis Acid - electron pair acceptor n Lewis Base - electron pair donor n Most general of all 3 definitions; acids don’t even need hydrogen! n Sample Problem 20-7, p.599
Hydrogen Ions from Water n Water ionizes, or falls apart into ions: H 2 O H 1+ + OH 1- n Called the “self ionization” of water n Occurs to a very small extent: [H 1+ ] = [OH 1- ] = 1 x 10 -7 M
Hydrogen Ions from Water n Since they are equal, a neutral solution results from water n K w = [H 1+ ] x [OH 1- ] = 1 x 10 -14 M 2 n K w is called the “ion product constant”
Ion Product Constant H 2 O H + + OH - H 2 O H + + OH - n K w is constant in every aqueous solution: [H + ] x [OH - ] = 1 x 10 -14 M 2 n If [H + ] > 10 -7 then [OH - ] 10 -7 then [OH - ] < 10 -7 n If [H + ] 10 -7
Ion Product Constant n If we know one, other can be determined n If [H + ] > 10 -7, it is acidic and [OH - ] 10 -7, it is acidic and [OH - ] < 10 -7 n If [H + ] 10 -7 n Basic solutions also called “alkaline”
Example [H+] = 1 x 10 -5 M Acid, Base or Neutral 10 -5 > 10 -7 10 -5 > 10 -7 Acidic
The pH concept n Logarithms are powers of ten. n definition: pH = -log[H + ] n in neutral pH = -log(1 x 10 -7 ) = 7 n in acidic solution [H + ] > 10 -7 n in basic solution [H + ] < 10 -7
The pH concept n pH < -log(10 -7 ) n pH < 7 (0 to 7 is the acid range) n in base, pH > 7 (7 to 14 is base range)
pH and pOH n pOH = -log [OH - ] n [H + ] x [OH - ] = 1 x 10 -14 M 2 n pH + pOH = 14 n Thus, a solution with a pOH less than 7 is basic; with a pOH greater than 7 is an acid
What is the pH of a solution with H– ion concentration of 4.2 X 10 -10 M. pH = -log [H+] pH = -log (4.2 X 10 -10 ) pH = -(log 4.2 + log 10 -10 ] pH = -(0.645 + (-10)) pH = 9.38 Alkaline
pH is 6.35, what is the H–ion concentration? pH = -log [H+] -log [H+] = 6.34 log [H+] = -6.34 [H+] = 10 -6.34 [H+] = 4.5 X 10 -7 M
Measuring pH n Why measure pH? –Everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc. n Sometimes we can use indicators, other times we might need a pH meter
Acid-Base Indicators n An indicator is an acid or base that undergoes dissociation in a known pH range, and has different colors in solution (more later in chapter) n Examples: litmus, phenolphthalein, bromthymol blue: Fig 20.8, p.590
Acid-Base Indicators n Although useful, there are limitations to indicators: –usually given for a certain temperature (25 o C), thus may change at different temperatures –what if the solution already has color? –ability of human eye to distinguish colors
Acid-Base Indicators n A pH meter may give more definitive results –some are large, others portable –works by measuring the voltage between two electrodes –needs to be calibrated
Strength n Strong acids and bases are strong electrolytes –They fall apart (ionize) completely. –Weak acids don’t completely ionize. n Strength different from concentration n Strong-forms many ions when dissolved n Mg(OH) 2 is a strong base- it falls completely apart when dissolved. –But, not much dissolves- not concentrated
Measuring strength n Ionization is reversible. n HAH + + A - n This makes an equilibrium n Acid dissociation constant = K a n K a = [H + ][A - ] (water is constant) [HA] n Stronger acid = more products (ions), thus a larger K a (Table 20.8, p.602)
What about bases? n Strong bases dissociate completely. n B + H 2 O BH + + OH - n Base dissociation constant = K b n K b = [BH + ][OH - ] [B] (we ignore the water) n Stronger base = more dissociated, thus a larger K b.
Strength vs. Concentration n The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume n The words strong and weak refer to the extent of ionization of an acid or base n Is concentrated weak acid possible?
Practice n Write the expression for HNO 2 n Write the K b for NH 3 n Sample 20-8, p. 604 n Carefully study Key Terms and equations, p. 608 n Be sure to do the ChemASAP programs, and take all the self-tests that are available!