Acids, Bases, and Salts

Acids, Bases, and Salts

Acids, Bases, and Salts Chapter 19 1 1. Properties of Acids (w/ an example of an acid) [green, no dot] 2. Properties of Bases ((w/ an example of a base) [green, w/dot] 3. Similarities between Acid & Bases [yellow] 4. Types of Acids (Monoprotic, diprotic, Triprotic) w/ example of each [red] 5. Arrhenius Acids and Bases [black] 6. Brnsted-Lowry Acids and Bases [white]

7. Lewis Acids and Bases [blue] 2 19.1 Acid- Base Theories Properties of Acids Taste sour React with metals to form H2 gas Will change the color of and acid-base indicator Turns blue litmus paper RED Properties of Bases Taste bitter feel slippery

Will change the color of and acid-base indicator Turns red litmus paper BLUE BOTH Can be strong or weak electrolytes in aqueous solutions React with the other to produce salt (an ionic compound) and water 3 Examples Acid: citric acid Base: milk of magnesia Magnesium hydroxide

4 Types of acids Monoprotic acids are acids that contain one ionizable hydrogen. Nitric acid: HNO3 Diprotice acids are acids that contain two ionizable hydrogens. Sulfuric acid: H2SO4 Triprotic acids are acids that contain three ionizable hydrogens. Phosphoric acid: H3PO4 5

Arrhenius Acids and Bases Svante Arrhenius in1887 posed a way of defining acids and bases Arrhenius Acids Are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) [H+ producer] HCl + H2O H3O+ + Cl NH4+ + H2O H3O+ + NH3 Arrhenius Bases Are compounds that ionize to yield hydroxide ions (OH-) in aqueous solutions. [OH- producer] NH3 + H2O NH4+ + OH CH3COO + H2O CH3COOH + OH

Arrhenius definiton incoporates the FEWEST number of 6 Brnsted-Lowry Acids and Bases In 1923 Johannes Brnsted and Thomas Lowry proposed a new definition Is a more comprehensive definition of acids and bases Brnsted-Lowry Acids Is a hydrogen-ion donor Brnsted-Lowry Bases

Is a hydrogen-ion acceptor 7 Lewis Acids and Bases Gilbert Lewiss theory of acids and bases was an extension of his concept of electron pairs This is the broadest acid/base definition used. Lewis acids accepts a pair of electrons Lewis bases donate a pair of electrons Acid Base 8

Acid-Base Definitions Type Acid Base Arrhenius (smallest definition) H+ producer OH- producer

H+ donor H+ acceptor Brnsted-Lowry Lewis (broadest definition) Electron-pair acceptor Electron pair donor

9 Acid, Bases & Salt Video Pre Test Video Quiz 1. 2. 3. 4. 5. 1.

True False (about equal) True True True* (when in water can form acidic/basic solutions) 6. False 7. True 8. True 9. False (Base has pH above 7) 10. False (stomach acids) 2.

3. 4. 5. 6. 7. 8. 9. 10. True False [acids react w/ metals] Neutralization False [acids produce H+ ion] True

Hydrogen ions True False {only 1 hydrogen ion True Acid rain Conjugate Acids and Bases Conjugate acid is the particle formed when a base gains a hydrogen ion. Conjugate base is the particle formed when an acid has donated a hydrogen ion. Conjugate acid-base pair consists of two substance related by the loss or gain of a single hydrogen ion. When a water molecule gains a hydrogen ion it

becomes a positively charged hydronium ion (H3O+). When water molecule looses a hydrognet ion it becomes the negativly charged hydroxide ion (OH-1) 11 Conjugate Acid and Bases HCl + H2O H3O+ + Cl Acid + Base Conjugate + Conjugate Acid Base NH3 + H2O NH+4 + OH Base + Acid Conjugate + Conjugate Acid Base

A substance that can act as both an acid and a base is said to be amphoteric. (example: Water ) 12 Conjugate acids and base What are the conjugate bases of these acids? a) HClO4 b) HSO4 c) HC2H3O2 d) H2S a) ClO4 b) SO42 c) C2H3O2 d) HS What are the conjugate acids of these bases a) SO42 b) NH3 c) F d) NO3 a) HSO4 b) NH4+ c) HF d) HNO3 13

Acid-Base Definitions Type Arrhenius (smallest definition) Brnsted-Lowry Lewis (broadest definition) Acid Base H+ producer

OH- producer H+ donor H+ acceptor Electron-pair acceptor Electron pair donor 14

19.1 Review 1. What are the properties of acids and bases? 2. How did Arrhenius define and acid and base? 3. How are acids and bases defined by the Brnsted-Lowry theory? 4. What is the Lewis-theory of acids and bases? 5. Identify the following as monoprotic, diprotic or triprotic: a) H2CO3 b) H3PO4 c) HCl d) H2SO4 15 The pH Concept The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

pH = - log [H+] For example a neutral solution has [H+] of 1.0 x 10 -7 so the pH is calculated pH = -log (1.0 x 10 -7 ) = 7.00 The pOH of a solution is the negative logarithm of the hydroxide concentration. pOH = - log [OH-1] 16 pH and Significant Figures A [H+] of 6.0 x 10 -5 has two significant figures The pH is recorded with two decimal places 4.22 A [H+] of 6. x 10 -5 has one significant figures

The pH is recorded with one decimal places 4.2 17 Calculating pH practice 1. [H+] of 1.0 x 10 -11 pH = -log (1.0 x 10 -11 ) = 11.00 2. [H+] of 6.0 x 10 -5 pH = -log (6.0 x 10 -5 ) = 4.22 3. [H+] of 4.0 x 10 -3 pH = -log (4.0 x 10 -3 ) = 2.40

4. [H+] of 9.0 x 10 -9 pH = -log (9.0 x 10 -9 ) = 8.05 18 Calculating pOH pOH = - log [OH-] 1. [OH-] of 1.0 x 10 -3 pOH = -log (1.0 x 10 -3 ) = 3.00 2. [OH-] of 6.0 x 10 -7 pOH = -log (6.0 x 10 -7 ) = 6.22 3. [OH-] of 2.0 x 10 -4

pOH = -log (2.0 x 10 -4 ) = 3.70 4. [OH-] of 7.2 x 10 -9 pOH = -log (7.2 x 10 -9 ) = .14 19 A solution in which [H+] is greater than 1.0 x 10 -7 has a pH less than 7.0 and is acidic. The pH of a neutral solution is 7.0 A solution in which [H+] is less than 1.0 x 10 -7 has a pH greater than 7.0 and is basic. A simple relationship between pH and pOH allows you to calculate the other if one is known.

pH + pOH = 14 20 Calculating concentration from pH/pOH Using pH to find [H+] since pH= -log [H+] Then [H+] = 10(-pH) Calculate the [H+] for the solution pH= 3.00 [H+]= 10(-pH) [H+] = 10(-3.00) = 1.0 x 10-3 M H+ 0.00100 M H+ Using pH to find [OH-1] since pOH= -log [OH-1] Then [OH-1] = 10(-pOH)

Calculate the [OH-1] for the solution pOH= 7.67 [OH-1]= 10(-pOH) [OH-1] = 10(-7.67) = 2.14 x 10-8 21 pH to pOH If [H+1] is 2.5 x 10-4 what is the pOH pH = -log [2.5 x 10-4 ] = 3.60 pOH = 14 3.60 = 10.40 If the [OH-1] is 6.8 x 10-12 what is the pH pOH = -log [6.8 x 10-12 ] = 11.17 pH = 14 11.17 = 2.83

22 19.2 Hydrogen Ions and Acidity The reaction in which water molecules produce ions is called the self-ionization of water H+(aq) + OH(aq) H2O(l) 2H2O(l) H3O+(aq) + OH(aq) 23

Self-ionization of water occurs to a very small extent, in pure water at 25C. Any aqueous solution in which the [H+] and [OH-] are equal is called a neutral solution. In aqueous solution [H+] x [OH-] = 1.0 x 10-14 This is called the ion-product constant for water (Kw). The concentrations may change but the product always equals 1.0 x 10-14 for water. 24 Acidic Solutions

An acidic solution is one in which the [H+] is greater than the [OH-]. The [H+] of an acidic solution is greater than 1.0 x 10-7 M HCl(g) H+ (aq) + Cl (aq) 25 Basic Solution An basic solution is one in which the [H+] is

less than the [OH-]. The [H+] of an acidic solution is less than 1.0 x 10-7 M NaOH Na+ + OH Basic solutions are also known as alkaline solutions. 26 What is the [OH-] if the [H+] is 1.0 x 10 -3? remember: [OH-] x [H+] = 1.0 x 10 -14 [OH-] = 1.0 x 10 -14/ 1.0 x 10 -3= 1.0 x 10 -11 What is the [H+] if the [OH-] is 1.0 x 10 -8 ? [H+] = 1.0 x 10 -14/ 1.0 x 10 -8= 1.0 x 10 -6

Classify each solution as acidic, basic, or neutral: a) [H+] = 6.0 x 10 -10 b) [OH-] = 3.0 x 10 -4 c) [H+] = 6.0 x 10 -7 d) [H+] = 1.0 x 10 -7 27 What is the [H+] of a solution with a pOH of 3.12? Is it acidic, basic, or neutral? pH= 14- pOH pH= 14 3.12 = 10.88 [H+]= 10(-pH) [H+] = 10(-10.88) = 1.3 x 10-11 it is a Basic solution What is the [H+] of a solution with a pOH of 9.18? Is it acidic, basic, or neutral?

pH= 14- pOH pH= 14 9.18 = 4.82 [H+]= 10(-pH) [H+] = 10(-4.82) = 1.5 x 10-5 it is an acidic solution 28 Concentration to pH and back Possible equations pH = -log [H+1] pOH = -log [OH-1] pH + pOH =14 [H+1] = 10-pH [OH-1] = 10-pOH [H+1]x[OH-1]=10-14 If the [OH-1] = 4.68 x 10-3 determine the pOH, pH, [H+1] and if acidic/basic/neutral.

o pOH = - log [4.68 x 10-3 ] = 2.33 o pH = 14-2.33 = 11.67 o [H+1] =10 -11.67 = 2.14 x 10-12 o Substance is Basic because pH is 11.67 29 Strength of Acids & Bases

What does it mean to be a STRONG acid? What are the 6 strong acids? (name and formula) What does it mean to be a STRONG base? What are the 8 strong bases? (name and formula) What does it mean to be a WEAK acid or WEAK base? 19.3 Strengths of Acids and Bases Strong acids are completely ionized in aqueous solutions. HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) 100% ionized Weak acids only slightly ionize in aqueous solutions. CH3COOH(aq) + H2O(l)

H3O+(aq) + CH3COO-(aq) Less than 0.4% ionized 31 Six strong acids!! HCl hydrochloric acid HBr hydrobromic acid

HI hydroiodic acid HNO3 nitric acid H2SO4 sulfuric acid HClO4 perchloric acid All other acids are considered weak acids!! 32 Base Dissociation Constant Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution. Weak bases react with water to form the hydroxide ion and the conjugate acid of the base.

NH3(aq) + H20(l) NH4+(aq) + OH-(aq) 33 Eight strong bases LiOH - lithium hydroxide

NaOH - sodium hydroxide KOH - potassium hydroxide RbOH - rubidium hydroxide CsOH - cesium hydroxide Ca(OH)2 - calcium hydroxide Sr(OH)2 - strontium hydroxide Ba(OH)2 - barium hydroxide 34 Strong Acids or Bases [H+] for strong acids equals molarity of acid [H+]= M of acid [OH-] for strong bases equals molarity of base

[OH-]= M of base What is the [OH-] and [H+] for the following solutions a. 0.275 M HCl b. 0.500 M NaOH c. 0.200 M HNO3 d. 0.375 M Ba(OH)2 35 Remember [H+] x [OH-] = 1.0 x 10-14 a. 0.275 M HCl [H+]= 0.275 [OH-] = 3.64 x 10-14 b. 0.500 M NaOH

[OH-] =0.500 [H+]= 2.00 x 10-14 c. 0.200 M HNO3 [H+]=0.200 [OH-] = 5.00 x 10-14 d. 0.375 M Ba(OH)2 [OH-] =0.375 [H+]= 2.67 x 10-14 36 19.4 Neutralization Reactions When you mix a strong acid with a strong base a neutral solution results. HCl (aq)

+ NaOH (aq) NaCl (aq) + H2O (l) 2 HBr (aq) + Ba(OH)2 (aq) BaBr2(aq) + 2 H2O (l) A neutralization reaction is where an acid and a base react to form a salt and water 37 Titration Acids and bases do not always react in a 1:1 ratio The equivalence point is when the number of

moles of hydrogen ions equals the number of moles of hydroxide ions. How many moles of hydrochloric acid are required to neutralized 0.50 moles of barium hydroxide ? 2HCl (aq) + Ba(OH)2 (aq) BaCl2 (aq) + 2H2O (l) 0.50 mol Ba(OH)2 (aq) x 2 mol HCl = 1.0 mol HCl 1 mol Ba(OH)2 38 The process of adding a known amount of a solution of known concentration to determine the concentration of another solution is called titration. The solution of known concentration is called

the standard solution. We use a buret to add the standard solution. The solution is added until the indicator changes colors. The point at which the indicator changes colors is the end point of the titration. 39 Titration of a strong acid with a strong base 40 Titration Calculations 1st type of problem

Determine the volume of 2.1 M sodium hydroxide needed to titrate 435 mL of 1.75 M sulfuric acid. 1st : write a balanced chemical equation 2nd : determine the number of moles of reactant (sulfuric acid) available 3rd: use stoichiometry to calculate the number of moles of reactant (sodium hydroxide) needed 4th : use molarity to determine the number of mL needed to complete the neutralization. 41 Determine the volume of 2.1 M sodium hydroxide needed to titrate 435 mL of 1.75 M sulfuric acid.

1st : 2NaOH + H2SO4 Na2SO4 +2H2O 2nd : 0.435L x [1.75 mol/1L]= 0.76125 mol H2SO4 3rd: 0.76125mol H2SO4 x 2mol NaOH = 1.5225 mol NaOH 1 mol H2SO4 4th : 1.5225 mol NaOH x __1 L = 0.725 L = 725 ml 2.1 mol 42 Determine the volume of 3.3 M hydrochloric acid needed to titrate 355 mL of 2.50 M magnesium hydroxide . 1st : 2HCl + Mg(OH)2 MgCl2 +2H2O

2nd : 0.355L x [2.50 mol/1L]= 0.8875 mol Mg(OH)2 3rd: 0.8875mol Mg(OH)2 x 2mol HCl = 1.775 mol HCL 1 mol Mg(OH)2 4th : 1.775mol HCl x __1 L = 0.538 L = 538 ml HCl 3.3 mol 43 Titration Calculations 2nd type of problem Titration reveals that 22.5 ml of 2.0 M nitric acid are required to neutralize 20 ml of calcium hydroxide. What is the molarity of the calcium hydroxide? 1st : 2HNO3 + Ca(OH)2 Ca(NO3)2 +2H2O

2nd : 0.0225L x [2.0 mol/1L]= 0.045 mol HNO3 3rd: 0.045mol HNO3 x 1mol Ca(OH)2 = 0.0225 mol Ca(OH)2 2 mol HNO3 4th : 0.0225mol Ca(OH)2 = 1.125 M Ca(OH)2 0.020 L 44 Titration reveals that 11.2 ml of 3.5 M barium hydroxide are required to neutralize 50 ml of perchloric acid. What is the molarity of the perchloric acid? 1st : 2HClO4 + Ba(OH)2 Ba(ClO4)2 +2H2O 2nd : 0.0112L x [3.5 mol/1L]= 0.0392 mol Ba(OH)2 3rd: 0.0392mol Ba(OH)2 x 2 mol HClO4 = 0.0784 mol HClO4

1 mol Ba(OH)2 4th : 0.0784 mol HClO4 = 1.56 M HClO4 0.050 L 45 19.5 Salt hydrolysis the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ions to the water. In general, salts that produce acidic solutions contain positive ions that release protons to water Salts that produce basic solutions contain negative ions that attract protons from water.

46 19.5 Salts in Solution A salt consists of an anion from an acid and a cation from a base. It forms as the result of a neutralization reaction. A buffer is a solution in which the pH remains relatively constant when small amount of acid or base are added. A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts. Buffers are able to resist drastic changes in pH. 47

The buffering capacity is the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs. Buffer systems are crucial in maintaining human blood pH within a narrow range. 48 49

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