Notes: Acids and Bases

Transcription

1 Name Chemistry Pre-AP Notes: Acids and Bases Period I. Describing Acids and Bases A. Properties of Acids taste ph 7 Acids change color of an (e.g. blue litmus paper turns in the presence of an acid) React with metals (check activity series) to produce gas React with to produce a salt ( compound) and water. This is known as a reaction. Are (strong or weak) B. Properties of Bases taste Feel ph 7 Bases change color of an (e.g. red litmus paper turns in the presence of a base) React with to produce a salt and water Are (strong or weak, if soluble) C. Review of Acid Nomenclature acids contain only hydrogen and one other element - They are named hydro root- ic acid HCl is named acid acids contain hydrogen and 2 other elements Most of these are also known as because of the presence of oxygen 1

2 Ternary acids are named based on the name -if the polytatomic ends in -ate, the acid name ends in H 2SO 4 is named acid -if the polyatomic ends in -ite, the acid name ends in II. Acid-Base Theories A. Arrhenius (Swedish chemist, ) H 2SO 3 is named acid Arrhenius defined acids as compounds that ionize to yield ions ( ) in aqueous solution Arrhenius defined bases as compounds that ionize to yield ions ( ) in aqueous solution Although Arrhenius described hydrogen ions as bare ions (H + ), we now know that they are hydrated in aqueous solution and exist as H + (H 2O) n, where n is a small integer. For simplicity, we represent the hydrated hydrogen ion as H 3O +. This hydrated form of the hydrogen ion is known as the ion. From now on, whether we use the formula H + (aq) or H 3O + (aq), we are always referring to the hydrated hydrogen ion. The Arrhenius theory is limited in scope, but it led to the development of other acidbase theories. B. Bronsted-Lowry (Bronsted: Danish chemist, /Lowry: English chemist, ) The Bronsted-Lowry theory defines an acid as a hydrogen ion and a base as a hydrogen ion HCl + H 2O H 3O + + Cl NH 3 + H 2O + NH 4 + OH A acid is the particle formed when a base has accepted H +. Every base has a conjugate acid. A base is the particle formed when an acid has donated H +. Every acid has a conjugate base. A conjugate acid-base pair consists of the two substances related by loss or gain of single hydrogen ion. In the above examples, list the conjugate acid-base pairs: 2

3 Practice 1 1. Identify the Bronsted-Lowry acid, the Bronsted-Lowry base, the conjugate acid, and conjugate base in the following. a) HCl + H 2O H 3O + + Cl - b) the ionization of nitric acid in water (similar to a) 2- c) CO 3 + H 2O - HCO 3 + OH - A substance that can act as either an acid or a base is called. The term is used to refer specifically to a substance that can either donate or accept protons (H + ). What substance in the Bronsted-Lowry examples on p. 2 is clearly amphiprotic? III. Strength of Acids and Bases A. Ionization of Acids The ion is responsible for the properties of acids Acids are classified as strong or weak based on the degree to which they in water (how many ions are formed) acids are completely ionized in aqueous solution List of 7 strong acids: HCl, HBr, HI, HClO 3,HClO 4, HNO 3, H 2SO 4 (if it s not in this list, it s a weak acid) acids ionize only slightly in aqueous solution B. Dissociation/Ionization of Bases Strong bases completely into ions and ions in aqueous solution Note: these are all the soluble hydroxides on your solubility chart (alkali metal hydroxides, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 ). NOTE: NH 4OH is produced from ammonia and is not a strong base. 3

4 Weak bases ionize partially/slightly when they react with water to form the ion and the conjugate of the base. You should be familiar with ammonia and related compounds as classic examples of weak bases. NH 3 + H 2O CH 3NH 2 + H 2O Note: strong acids/bases are electrolytes. WHY? weak acids/bases are electrolytes. WHY? IV. Ionization Constant of Water A. Concentration of H + and OH - ions Water self-ionizes to a small degree via this reversible reaction: This reaction happens continually in water. In pure water, the concentration of ions equals the concentration of ions [H + ] = [OH - ] = 1.0 x 10 7 brackets represent concentration in moles/l (M) The product of the concentration of hydronium ions and hydroxide ions in any solution that contains water is always. The value for this constant is 1.0 x It is known as K w, the ionization constant of water. K w = [H + ] x [OH - ] = 1.0 x Since the product of [H + ] and [OH - ] is constant, if [H + ] increases, then [OH - ] (and vice-versa). The [H + ] of a solution is 1 x 10-8 M. What is [OH - ]? Practice 3 1. Determine the [H 3O + ] a) [OH - ] = 1 x M c) 0.1 M HCl b) [OH - ] = 2.0 x 10-3 M 2. Determine the [OH - ] a) [H 3O + ] = 6.0 x 10-4 M c) 0.1 M NaOH b) [H 3O + ] = 5.0 x M d) 0.1 M strontium hydroxide If [H 3O + ] > [OH - ] the solution is. If [H 3O + ] < [OH - ] the solution is. If [H 3O + ] = [OH - ] the solution is. 4

5 Determine if the solution is acidic, basic, or neutral. a) [H 3O + ] = 1.0 x 10-4 M b) [H 3O + ] = 6.0 x M c) [OH - ] = 4.0 x M Practice 4 1. Determine if the solutions in Practice 3 are acidic, basic, or neutral V. ph Expressing hydrogen ion concentration in can be cumbersome. A widely used system for expressing [H 3O + ] is the scale. The of a solution is the negative base 10 logarithm of the ion concentration. The base 10 logarithm of a number is the power to which 10 must be raised to equal that number. log = 2 (because 100 = 10 2 ) What is log ? ph = Remember, for pure water, [H 3O + ] is 1 x 10-7 M. So what s the ph of pure water? What is the ph of a solution with a hydronium ion concentration of 1.0 x M? Practice 5 1. What is the ph of the following solutions? State whether acidic, basic, or neutral. a) [H 3O + ] = 1.0 x 10-2 M e) 1 M HCl b) [H 3O + ] = 1.0 x 10-9 M c) [H 3O + ] = M d) [OH - ] = 1 x 10-3 M The ph scale typically ranges from 0 ([H 3O + ] = 1 M) to 14 ([H 3O + ] = M).Note that as ph increases, [H 3O + ] and [OH - ]. 5

6 Note the relationship between [H 3O + ] and [OH - ]. Remember, the product of these must always equal for aqueous solutions. The poh of a solution equals the negative base 10 logarithm of the ion concentration. poh = ph + poh = Practice 6 1. Calculate the ph of each solution a) [H 3O + ] = 5.0 x 10-6 M d) 0.15 M HClO 4 b) [H 3O + ] = 8.3 x M e) 0.15 M barium hydroxide (tricky!) c) [OH - ] = 4.3 x 10-5 M Practice 7 1. Calculate the [H 3O + ] for each solution. a) ph = 5.0 b) ph = 7.0 c) ph = 12.0 Practice 8 1. Calculate the [H 3O + ] for each of the following solutions a) ph= 7.30 b) ph = 1.80 c) poh = 11.2 d) poh = Determine the molarity of a) a potassium hydroxide solution with a ph of 11.6 b) a hydrochloric acid solution with a poh of 11.6 VI. Net Ionic Equations part 2: Neutralization Reactions A neutralization reaction is the reaction between an acid and a base to yield a salt and water. Neutralization reactions, like precipitation reactions, are double replacement reactions. We will learn to write net ionic equations for neutralization reactions. Helpful hints applicable to neutralization: 1. Remember to always write the balanced molecular equation first. Based upon the molecular equation, you will write the complete ionic and then the net ionic equation. 6

7 2. All soluble ionic compounds are written in their dissociated form, showing the ions and their charges with the symbol (aq) by each. Example: NaCl(aq) is written Na + (aq) + Cl-(aq). 3. All strong acids and bases are written in their ionized form (acids) or their dissociated form (bases). You must memorize a list of strong acids and bases so that you will know which to write ionized/dissociated. (See p. 3 of these notes) Examples: HClO 4(aq) is written H + - (aq) + ClO 4 (aq) H 2SO 4(aq) is written 2H + 2- (aq) + SO 4 (aq) HNO 3(aq) is written H + - (aq) + NO 3 (aq) HI(aq) is written H + (aq) + I - (aq) KOH(aq) is written K + (aq) + OH - (aq) Sr(OH) 2(aq) is written Sr 2+ (aq) + 2OH - (aq) These acids are written ionized because they are considered to be 100% ionized in solution. The bases are strong because they are soluble and 100% dissociated in solution. 4. All other acids and bases are considered weak and written in their unionized or undissociated forms. Examples are: NH 3(aq) written NH 3(aq) H 3PO 4(aq) written H 3PO 4(aq), no charges shown HC 2H 3O 2(aq) written HC 2H 3O 2(aq), no charges shown 5. Water is a pure liquid and is written with its formula, the symbol (l), with no charges shown: H 2O(l) VII. Neutralization and Titrations A. Neutralization A neutralization reaction is one in which an acid and base react to produce a and water. Remember, to a chemist, a salt is any compound that is not a base. HC 2H 3O 2 + NaOH Complete and net ionic equations for this neutralization reaction: Note in the reaction above, the acid:base mole ratio is 1:1. However, in the reaction between H 2SO 4 and NaOH: it takes moles of base to neutralize 1 mole of acid. The reacting ratios of acid and base will be important in solving problems related to neutralization reactions. 7

8 B. Titrations A is a laboratory technique that uses a neutralization reaction to determine the concentration of an unknown acid or base. An is used to show when neutralization has occurred. An indicator is a substance that changes color in response to changes in ph. Indicators are actually weak or that change color in response to ph change. In acidic solutions, indicators act as Bronsted-Lowry. As the indicator molecules accept H +, they change color. In basic solutions, indicators act as Bronsted-Lowry. As the indicator molecules donate H +, they change color. Steps in a titration: 1. A measured volume of a solution of unknown concentration (acid or base) is added to an flask. 2. A solution of known molarity (acid or base) is added to a. 3. Several drops of an are added to the unknown solution. 4. A measured volume of the solution of known molarity (acid or base) is added to the unknown solution until the indicator just barely changes. The point at which all the acid and base have exactly neutralized each other according to the mole:mole ratio from the balanced equation is called the point. The point at which the indicator changes color is the point of the titration. The indicator needs to be chosen wisely so that it changes color as close to the equivalence point of the titration as possible. 8

9 Examples of Common Indicators: Indicator Acid Color Base Color ph range* Titration Type Methyl red Red Yellow Strong acid/weak base Bromthymol Blue Yellow Blue Strong acid/strong base Phenolphthalein Colorless Magenta Weak acid/strong base *where indicator changes color C. Solving Titration Problems To solve titration problems: 1. Write a balanced neutralization reaction. 2. Use the following formula: nbmava = nambvb nb = coefficient of base in reaction na = coefficient of acid in reaction Ma = Molarity of acid Mb = Molarity of base Va = Volume of acid Vb = Volume of base Ex 1: A 25 ml solution of nitric acid is completely neutralized by 18 ml of 1.0 M calcium hydroxide. What is the concentration (M) of the acid solution? Practice 9 1. What is the molarity of sulfuric acid if 15.0 ml of the solution is completely neutralized by 38.5 ml of M sodium hydroxide? 2. What is the ph of a sodium hydroxide solution if 20.0 ml of the solution is neutralized by 28.0 ml of 0.60 M hydrochloric acid? 3. What volume of 0.45M hydrochloric acid must be added to 25.0 ml of 1.00 M barium hydroxide to make a neutral solution? 9