Acids and Bases. Properties, Reactions, ph, and Titration

Transcription

1 Acids and Bases Properties, Reactions, ph, and Titration C

2 Properties of acids 1. Taste Sour (don t try this except with foods). 2. Are electrolytes (conduct electricity). Some are strong, some are weak. 3. Change indicator colors. (litmus red). 4. React with metals to form hydrogen gas. 5. React with hydroxides to form water and a salt.

3 Acid s Reaction with Metals Metals: Dissolves; Problem: bridges, cars, buildings Magnesium: Iron: 2HCl + Mg ---> H 2 + MgCl 2 2HCl + Fe ---> H 2 + FeCl 2 Copper: 2HCl + Cu ---> H 2 + CuCl 2

4 Common Acids Fruits citric acid Milk lactic acid Vinegar acetic acid Soda pop carbonic and phosphoric acid And lots more!!!!

5 Properties of Bases 1. React with acids to form water and a salt. 2. Taste bitter.(don t try this) 3. Feel slippery (Don t try this either). 4. Can be strong or weak electrolytes. 5. Change indicators (litmus blue).

6 Common Bases Windex ammonia Baking soda sodium bicarbonate Drain cleaner NaOH Milk of Magnesia Mg(OH) 2 And more..

7 Organic acids found in living things (fruits, etc) contain -COOH a carboxyl group weak acids are only slightly ionized to -COO- Called carboxylic acids

8 Mineral acids from inorganic materials (rocks) traditional acids - used industrially

9 Common Industrial Acids Sulfuric acid - H 2 SO 4 petroleum, fertilizer, metallurgy, paper, paints,batteries, etc Nitric acid HNO 3 explosives, rubber, plastics, pharmaceuticals, etc. Phosphoric H 3 PO 4 fertilizer, flavoring agent, detergents, etc. Hydrochloric HCl pickling metal, cleaning, chlorination (pools) Acetic Acid CH 3 COOH plastics, food supplements, etc.

10 Nomenclature Two basic types of acids: binary and oxyacids 1. Binary acids 2 elements only hydro + stem + ic acid HCl hydrochloric acid HI hydroiodic acid

11 Nomenclature 2. Oxyacid names anion stem + ous (ite anions) NO -1 2 (nitrite) HNO 2 nitrous acid Or anion stem + ic (ate anions) NO 3-1 (nitrate) HNO 3 nitric acid HClO 4 perchloric acid

12 Nature: Electrolytes are classified as Acids, Bases, or Salts Acids - react with H 2 O and produce H + The H + ion combines with water and forms H 3 O + called the hydronium ion Bases dissociate with H 2 O and produce OH - Salts - Ionic combinations of metal/nonmetal ions.

13 Strong vs. Weak STRONG electrolytes show complete ionization in water (all ions); good conductors Soluble salts, SA, SB NaCl Na + (aq) + Cl - (aq) WEAK electrolytes show partial ionization in water (mostly molecules); poor conductors WA, WB NH 3 + H + NH 4 +

14 Aqueous acids Arrhenius definition: acids ionize in water to form H + ions - - are polar covalent compounds and all have H. - may ionize in more than 1 step. (ex H 2 SO 4 ) Strong acids show complete ionization (100%) HA H +1 + A -1 Weak acids produce few ions (less than 5%); are dissolved intact as molecules. HA H +1 + A -1

15 Arrhenius Base Bases dissociate and produce OH- ions. Strong bases 100% dissociation Group I and II hydroxides Weak bases less than 5% dissociation Ammonia, aniline, carbonates are not included. All other hydroxides are.

16 Memorize the Strong Acids HCl - hydrochloric HBr hydrobromic HI - hydroiodic H 2 SO 4 - sulfuric HClO 4 perchloric HNO 3 - nitric

17 Memorize the Strong Bases NaOH - sodium hydroxide KOH - potassium hydroxide LiOH lithium hydroxide RbOH - rubidium hydroxide Ba(OH) 2 barium hydroxide Sr(OH) 2 strontium hydroxide Ca(OH) 2 - calcium hydroxide Mg(OH) 2 magnesium hydroxide

18 Acid definitions Bronsted Lowry Acids are proton donors Bases are proton acceptors Acids and bases occur in conjugate pairs

19 Come in Pairs General equation HA(aq) + H 2 O(l) Acid + Base H 3 O + (aq) + A - (aq) Conjugate + Conjugate acid base

20 Conjugate pairs This is an equilibrium. B(aq) + H 2 O(l) BH + (aq) + OH - (aq) Base + Acid Conjugate acid +Conjugate base NH 3 (aq)+h 2 O(l) NH 4 + (aq)+oh - (aq)

21 In Bronsted-Lowry theory, bases do not require OH - Bases are able to accept protons Allows ammonia and carbonate ions to be considered bases, others as well. NH 3 + H + NH 4 + Base + H + Conjugate acid Most accepted theory

22 Acid & Base Reactions Neutralization Reaction: Acid + Base salt + H 2 O (usually) Salt = general term for an ionic compound Example: HCl + NaOH NaCl + H 2 O

23 Acid-Base reactions Are equilibrium reactions (reversible) Compare strength of the two acids (charts) Equilib. shifts away from the stronger acid.

24 HClO 4 + H 2 O H 3 O + + ClO 4 - Acid + base cong.acid + cong. Base HClO 4 is a stronger acid than H 3 O + so. Equilibrium shifts to the right away from HClO 4

25 Protons are Hydrogen ions Monoprotic acids have one proton to donate ex. HCl Diprotic acids have two protons to donate ex. H 2 SO 4 (one step at a time) Polyprotic two or more protons to donate ex. H 3 PO 4

26 Amphoteric substances Substances which can either accept or donate a proton. Water is an example H 2 O + H + H 3 O + (water as a base) H 2 O H + + OH - (water as an acid) Other examples are NH 3 and HSO 4 -

27 Lewis Theory Lewis Acid accepts an electron pair Lewis Base donates an electron pair Not frequently used for chemists Most general definition (same G. Lewis that made e-dot diagrams)

28 Lewis Acids and Bases Lewis Acid A species (atom, ion or molecule) that is an electron pair acceptor. Lewis Base A species that is an electron pair donor. base acid adduct

29 Showing Electron Movement

30 Focus On Acid Rain CO 2 + H 2 O H 2 CO 3 H 2 CO 3 + H 2 O HCO H 3 O + HCO H 3 O + CO 2 + H 2 O 3 NO 2 + H 2 O 2 HNO 3 + NO or SO 2 + H 2 O H 2 SO 3

31 Acid Rain

32 Acid rain Gases like sulfur dioxide and nitrogen dioxide are produced from burning coal, oil, and other fuels. These gases react with water vapor in the atmosphere to form acids. Acid rain can be stopped with govt. regulations. Less in US/Canada now, but more in China/India

33 Acid/Base Titration - a lab process Basic Concepts: 1. Acids & bases neutralize each other 2. From the balanced equation, the number of moles needed of the known reactant & the unknown reactant are given. 3. An indicator is selected based on the strength of the known reactant. 4. The indicator will change color when the known reactant equals the unknown. 5. Concentration of the unknown is calculated.

34 7 Steps 1. Fill Burette with NaOH (known) 2. Place 20ml HCl in flask (unknown) The amount may be different, but record 3. Place indicator in HCl 4. Slowly add NaOH until the endpoint is reached (color change). 5. Record amount of NaOH used (let s pretend 19.9ml) 6. Use the factor label method to find the number of moles of NaOH. 7. Look at the balanced equation to determine the ratio of moles between the Known NaOH & unknown HCl.

35 Titration calculation Use the equation: M a x V a = M b x V b Example: 25 ml of HCl is neutralized by 20 ml of 0.5 M NaOH. Find conc. of HCl. Solution: M a = M b x V b / V a M a = 0.5 M x 20mL / 25mL = 0.4 M HCl

36 Water Self ionization of water. (very small amount) H 2 O H + + OH - [H + ] = [OH - ] = 1 x 10-7 M A neutral solution. In water: Kw = [H + ] x [OH - ] = 1 x Kw is called the ion product constant.

37 Ion Product Constant H 2 O H + + OH - Kw is constant 1 x If [H + ] > 10-7 then [OH - ] < 10-7 (acidic) If [H + ] < 10-7 then [OH - ] > 10-7 (basic) If we know one, we can determine the other. If [H + ] = 1x 10-3 Find [OH-] K w / [H + ] = [OH-] 1 x /1 x 10-3 = [OH-] = 1 x 10-11

38 Logarithms Powers of ten. A shorthand form ph = -log[h + ] in neutral ph = - log(1 x 10-7 ) = 7 in acidic solution [H + ] > 10-7 ph < -log(10-7 ) ph < 7 in base ph > 7

39 ph and poh equations ph = -log[h + ] poh = - log [OH - ] [H + ] x [OH - ] = 1 x ph + poh = 14

40 [H + ] ph Acidic Neutral Basic poh Basic [OH - ]