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Acids, Bases, and Salts Chapter 19. Acid-Base Theories Essential Question: What are the properties of acids and bases, and what distinguishes the Arrhenius,

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Presentation on theme: "Acids, Bases, and Salts Chapter 19. Acid-Base Theories Essential Question: What are the properties of acids and bases, and what distinguishes the Arrhenius,"— Presentation transcript:

1 Acids, Bases, and Salts Chapter 19

2 Acid-Base Theories Essential Question: What are the properties of acids and bases, and what distinguishes the Arrhenius, Bronsted-Lowry and Lewis theories of acids and bases?

3 Properties of Acids Acids taste tart or sour. Aqueous solutions of acids are electrolytes. Acids change the color of acid-base indicators. Acids react with metals to produce hydrogen gas. Acids react with bases to form water and a salt.

4 Properties of Bases Bases have a bitter taste. Bases have a slippery feel. Aqueous solutions of bases are electrolytes. Bases change the color of acid-base indicators. Bases react with acids to form water and a salt.

5 Three Major Theories Arrhenius Bronsted-Lowry Lewis

6 Arrhenius Acids and Bases Acids are hydrogen-containing compounds that ionize to produce H + ions in aqueous solution. Bases are hydroxide-containing compounds that ionize to produce OH – ions in aqueous solution.

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8 Mono- Di- and Tri-protic Acids HNO 3 is a monoprotic acid. H 2 SO 4 is a diprotic acid. H 3 PO 4 is a triprotic acid. Not all substances that contain hydrogen are acids. Not all hydrogens in acids are necessarily released as H + ions.

9 Hydrochloric Acid

10 Ethanoic Acid (Acetic Acid)

11 Arrhenius Bases The most commonly known is NaOH (lye). Another common base is KOH

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13 Bronsted-Lowry Acids and Bases Defines an acid as a hydrogen-ion donor. Defines a base as a hydrogen-ion acceptor.

14 Why Ammonia is a Base

15 Conjugate Acids and Bases When a substance donates a hydrogen ion, what remains has the ability to accept it back. When a substance accepts a hydrogen ion, what remains has the ability to donate the ion.

16 Conjugate Acids and Bases A conjugate acid is formed when a base gains a hydrogen ion. A conjugate base is remains when an acid has donated a hydrogen ion. A conjugate acid-base pair consists of two substances related by the loss or gain of H +.

17 Consider Ammonia in Water NH 3 + H 2 O NH OH – Base Acid Conjugate Conjugate Acid Base

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19 Hydronium ion ( H 3 O + ) A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion.

20 Amphoteric Sometimes water accepts a hydrogen. HCl + H 2 O H 3 O + + Cl – Other times, water donates a hydrogen. NH 3 + H 2 O NH OH – A substance that can act as either an acid or a base is said to be amphoteric.

21 Lewis Acids and Bases A Lewis acid accepts a pair of electrons during a reaction. A Lewis base donates a pair of electrons during a reaction.

22 Hydrogen Ions and Acidity Essential Question: How are [H + ] and [OH – ] related and how do they relate to acidity?

23 Hydrogen Ions from Water The reaction in which water molecules produce ions is called the self-ionization of water. H 2 O (l) H + (aq) + OH – (aq) In water, H + ions are always joined to water molecules to form H 3 O + hydronium ions.

24 Self-Ionization of Water

25 This happens to a very small extent. In pure water, the concentration of H + and OH – are equal. [H + ] and [OH – ] both equal 1.0 x M. This is called a neutral solution.

26 Ion Product Constant for Water In any aqueous solution, when [H + ] increases, [OH – ] decreases. When [H + ] decreases, [OH – ] increases. The product of the hydrogen-ion concentration and the hydroxide ion concentration always equals 1.0 x

27 Ion Product Constant for Water K w = [H + ] x [OH – ] = 1.0 x Remember, as [H + ] goes up, [OH – ] goes down. But the product of [H + ] x [OH – ] will always be 1.0 x

28 Acidic vs Basic An acidic solution is one in which the [H + ] is greater than the [OH – ]. A basic solution is one in which the [H + ] is less than the [OH – ]. Basic solutions are also called alkaline solutions.

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32 The pH Concept Expressing hydrogen concentration in molarity can be cumbersome. A more popular method is the pH scale. The pH scale ranges from 0 to 14, with the most acidic = 0 and the most basic = 14.

33 Calculating pH pH = the negative of the logarithm of the hydrogen ion concentration. pH = -log [H + ]

34 Calculating pOH pOH is the red-headed stepchild… pOH is the negative logarithm of the [OH – ]. pOH = -log [OH – ]

35 pH and pOH You can calculate the pH or the pOH of a solution using the log function key on a calculator.

36 pH and Acidity Acidic solution:pH < 7.0 and [H+] is greater than 1 × 10 −7 M Neutral solution:pH = 7.0and [H+] is equal to 1 × 10 −7 M Basic solution:pH > 7.0and [H+] is less than 1 × 10 −7 M

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39 pH and Significant Figures Express [H + ] and [OH – ] in scientific notation. Express pH and pOH with the same number of digits to the right of the decimal place as the number of significant digits in the scientific notation.

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41 Sample Problem 19.2

42 The Solution

43 Measuring pH

44 Two common methods are used: –Acid base indicators which change color –pH meters which measure electrical conductivity

45 Acid-Base Indicators Acid-Base indicators have different colors based upon acidity. For each indicator, the change takes place over a relatively narrow range of about 2 pH units.

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48 Effects of Acidity on Plant Color

49 Strengths of Acids and Bases Essential Question: How does the value of an acid dissociation constant relate to the strength of the acid, and how are those values calculated?

50 Strong and Weak Acids Strong acids are completely ionized in aqueous solution. HCl + H 2 O H 3 O + + Cl – (100% ionized) Weak Acids only slightly ionized in aqueous solution. CH 3 COOH + H 2 O H 3 O + CH 3 COO – (partially ionized)

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52 Acid Dissociation – Strong

53 Acid Dissociation – Weak

54 Acid Dissociation – Very Weak

55 Acid Dissociation Constants Takes the same form as the equilibrium-constant expression from a balance chemical equation. For ethanoic acid, for instance, the acid dissociation constant is calculated as follows: [H 3 O + ] x [CH 3 COO – ] K a = [CH 3 COOH] x [H 2 O] These are sometimes called ionization constants.

56 Acid Dissociation Constants Weak acids have small K a values. Strong acids have large K a values. The K a value for HCl (aq) is ∞ (infinite). Why do you think this is the case?

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58 Base Dissociation Constants Strong bases and weak bases refer to the degree of dissociation just like acids. For sodium hydroxide, for instance, the base dissociation constant is calculated as follows: [Na + ] x [OH – ] K b = [NaOH] x [H 2 O]

59 Neutralization Reactions Essential Question: What are the products of the reaction of an acid and a base when the endpoint of a titration is reach?

60 Acid—Base Reactions A mixture of a strong acid with an equal amount of a strong base results in a neutral solution. These reactions are called neutralization reactions.

61 Neutralization Reactions Consider these examples: HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) H 2 SO 4 (aq) + 2KOH(aq) K 2 SO 4 (aq) + H 2 O(l) What would the net ionic equations for each of these reactions be?

62 Titration Acids and bases sometimes react 1:1 HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) However, the ratio can vary H 2 SO 4 (aq) + NaOH(aq) Na 2 SO 4 (aq) + 2H 2 O(aq) 2HCl(aq) + Ca(OH) 2 (aq) CaCl 2 + 2H 2 O(l)

63 Titration When an acid and a base are mixed, the equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

64 Titration You can determine the concentration of an acid in a solution by performing a neutralization reaction. You must select an appropriate acid-base indicator. Phenolphthalein turns from colorless to pink as the pH changes from acidic to basic.

65 Titrations A measured volume of an acid solution of unknown concentration is added to a flask. Several drops of the indicator are added to the solution. Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color.


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