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Chapter 14 Acids and Bases 2006, Prentice hall. 2 CHAPTER OUTLINE  General Properties General Properties  Arrhenius Acids & Bases Arrhenius Acids &

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Presentation on theme: "Chapter 14 Acids and Bases 2006, Prentice hall. 2 CHAPTER OUTLINE  General Properties General Properties  Arrhenius Acids & Bases Arrhenius Acids &"— Presentation transcript:

1 Chapter 14 Acids and Bases 2006, Prentice hall

2 2 CHAPTER OUTLINE  General Properties General Properties  Arrhenius Acids & Bases Arrhenius Acids & Bases  Strength of Acids & Bases Strength of Acids & Bases  Brønsted-Lowery Acids & Bases Brønsted-Lowery Acids & Bases  Ionization of Water Ionization of Water  pH and pOH Scales pH and pOH Scales

3 3 GENRAL PROPERTIES OF ACIDS & BASES  Many common substances in our daily lives are acids and bases.  Oranges, lemons and vinegar are examples of acids. In addition, our stomachs contain acids that help digest foods.  Antacid tablets taken for heartburn and ammonia cleaning solutions are examples of bases.

4 4 GENRAL PROPERTIES OF ACIDS  General properties associated with acids include the following:  sour taste  change color of litmus from blue to red  react with metals to produce H 2 gas  react with bases to produce salt & water

5 binary acids have acid hydrogens attached to a nonmetal atom –HCl, HF Hydrofluoric acid Structure of Acids

6 oxy acids have acid hydrogens attached to an oxygen atom –H 2 SO 4, HNO 3

7 Structure of Acids carboxylic acids have COOH group –HC 2 H 3 O 2, H 3 C 6 H 5 O 3 only the first H in the formula is acidic –the H is on the COOH

8 8 GENRAL PROPERTIES OF BASES  General properties associated with bases include the following:  bitter taste  change color of litmus from red to blue  slippery, soapy feeling  react with acids to produce salt & water

9 Structure of Bases most ionic bases contain OH ions –NaOH, Ca(OH) 2 some contain CO 3 2- ions –CaCO 3 NaHCO 3 molecular bases contain structures that react with H + –mostly amine groups

10 amine groups H – N – H H Ammonia + H + H – N + – H H H Ammonium ion

11 11 ARRHENIUS ACIDS & BASES  The most common definition of acids and bases was formulated by the Swedish chemist Svante Arrhenius in 1884. According to the Arrhenius definition,  Acids are substances that produce hydronium ions (H 3 O + ) in aqueous solution. Commonly written as HCl (g) + H 2 O (l) H 3 O + (aq) + Cl – (aq) HCl (g) H + (aq) + Cl – (aq) H2OH2O

12 12 ARRHENIUS ACIDS & BASES According to the Arrhenius definition,  Bases are substances that produce hydroxide ion (OH - ) in aqueous solution. NaOH (s) Na + (aq) + OH – (aq) H2OH2O NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH – (aq)

13 13 BRØNSTED-LOWRY ACIDS & BASES  The Arrhenius definition of acids and bases is limited to aqueous solutions.  A broader definition of acids and bases was developed by Br ø nsted and Lowry in the early 20 th century.  According to Br ø nsted-Lowry definition, an acid is a proton donor, and a base is a proton acceptor. HCl (g) + H 2 O (l) → H 3 O + (aq) + Cl – (aq) AcidBase NH 3 (aq) + H 2 O (l) → NH 4 + (aq) + OH – (aq) AcidBase A substance that can act as a Brønsted-Lowry acid and base (such as water) is called amphiprotic.

14 14 BRØNSTED-LOWRY ACIDS & BASES  In Brønsted-Lowry definition, any pair of molecules or ions that can be inter-converted by transfer of a proton is called conjugate acid-base pair. HCl (g) + H 2 O (l) → H 3 O + (aq) + Cl – (aq) AcidBase Conjugate acid Conjugate base

15 15 BRØNSTED-LOWRY ACIDS & BASES AcidBase NH 3 (aq) + H 2 O (l) → NH 4 + (aq) + OH – (aq) Conjugate acid Conjugate base

16 16 AcidBase H 2 O + Cl   HCl + OH  Conjugate acid Conjugate base Example 1: Identify the conjugate acid-base pairs for each reaction shown below:

17 17 AcidBase C 6 H 5 OH + C 2 H 5 O   C 6 H 5 O  + C 2 H 5 OH Conjugate acid Conjugate base Example 1: Identify the conjugate acid-base pairs for each reaction shown below:

18 18 HS  Example 2: Write the formula for the conjugate acid for each base shown: + H +  H 2 S NH 3 + H +  NH 4 + CO 3 2  + H +  HCO 3 

19 19 HI Example 3: Write the formula for the conjugate base for each acid shown: - H +  I  CH 3 OH - H +  CH 3 O  HNO 3 - H +  NO 3 

20 20 ACID & BASE STRENGTH  Strong acids and bases are those that ionize completely in water.  Strong acids and bases are strong electrolytes.  According to the Arrhenius definition, the strength of acids and bases is based on the amount of their ionization in water. 100 1M

21 21 ACID & BASE STRENGTH  Weak acids and bases are those that ionize partially in water.  Weak acids and bases are weak electrolytes. 100 ~1~1~1~1 1M~0.01M

22 22 IONIZATION OF STRONG vs. WEAK ACIDS Ionizes completely Ionizes partially

23 23 COMMON ACIDS Strong AcidsWeak Acids HCl Hydrochloric acid HC 2 H 3 O 2 Acetic acid HBr Hydrobromic acid H 2 CO 3 Carbonic acid HI Hydroiodic acid H 3 PO 4 Phosphoric acid HNO 3 Nitric acid HF Hydrofluoric acid H 2 SO 4 Sulfuric acid HCN Hydrocyanic acid H2SH2S Hydrosulfuric acid

24 24 COMMON BASES Strong BasesWeak Bases LiOH Lithium hydroxide NH 3 Ammonia NaOH Sodium hydroxide CO(NH 2 ) 2 Urea Ca(OH) 2 Calcium hydroxide KOH Potassium hydroxide Ba(OH) 2 Barium hydroxide

25 25 COMPARISON OF ACIDS & BASES

26 Titration using reaction stoichiometry to determine the concentration of an unknown solution Titrant (unknown solution) added from a buret indicators are chemicals added to help determine when a reaction is complete the endpoint of the titration occurs when the reaction is complete

27 Titration

28 The base solution is the titrant in the buret. As the base is added to the acid, the H + reacts with the OH – to form water. But there is still excess acid present so the color does not change. At the titration’s endpoint, just enough base has been added to neutralize all the acid. At this point the indicator changes color.

29 Example 1 The titration of 10.00 mL HCl solution of unknown concentration requires 12.54 mL of 0.100 M NaOH solution to reach the end point. What is the concentration of the unknown HCl solution? Find: concentration HCl, M Collect Needed Equations and Conversion Factors: HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l)  1 mole HCl = 1 mole NaOH 0.100 M NaOH  0.100 mol NaOH  1 L sol’n

30 = 1.25 x 10 -3 mol HCl molarity = 0.125 M HCl

31 31 IONIZATION OF WATER  Water can act both as an acid and a base.  In pure water, one water molecule donates a proton to another water molecule to produce ions. H 2 O + H 2 O  H 3 O + + OH – AcidBase Conjugate acid Conjugate base

32 32 IONIZATION OF WATER  In pure water, the transfer of protons between water molecules produces equal numbers of H 3 O + and OH – ions.  The number of ions produced in pure water is very small, as indicated below: [H 3 O + ] = [OH – ] = 1.0 x 10 -7 M  When the concentrations of H 3 O + and OH – are multiplied together, the ion-product constant (K w ) is formed. K w =[H 3 O + ][OH – ]=[1.0x10 -7 ][1.0x10 -7 ]=1.0x10 -14  All aqueous solutions have H 3 O + and OH – ions.  An increase in the concentration of one of the ions will cause an equilibrium shift that causes a decrease in the other one.

33 33 ACIDIC & BASIC SOLUTIONS  When [H 3 O + ] and [OH – ] are equal in a solution, it is neutral.  When [H 3 O + ] is greater than [OH – ] in a solution, it is acidic.  For example, if [H 3 O + ] is 1.0 x 10 –4 M, then [OH – ] would be 1.0 x 10 –10 M.

34 34 ACIDIC & BASIC SOLUTIONS  When [OH - ] is greater than [H 3 O + ] in a solution, it is basic.  For example, if [OH - ] is 1.0 x 10 –6 M, then [H 3 O + ] would be 1.0 x 10 –8 M.

35 35 ACIDIC & BASIC SOLUTIONS [H 3 O + ]=[OH - ] Neutral [H 3 O + ]>[OH - ] Acidic [H 3 O + ]<[OH - ] Basic

36 36 pH SCALE  The acidity of a solution is commonly measured on a pH scale.  The pH scale ranges from 0-14, where acidic solutions are less than 7 and basic solutions are greater than 7. pH = -log [H 3 O + ]

37 37 pH SCALE Acidic solutionspH < 7H 3 O + > 1x10 -7 Neutral solutionspH = 7H 3 O + = 1x10 -7 Basic solutionspH > 7H 3 O + < 1x10 -7

38 38 Example 1: The [H 3 O + ] of a liquid detergent is 1.4x10 –9 M. Calculate its pH. pH = -log [H 3 O + ]= -log [1.4x10 -9 ]= -(-8.85) pH = 8.85 2 significant figures The number of decimal places in a logarithm is equal to the number of significant figures in the measurement. Solution is basic

39 39 Example 2: The pH of black coffee is 5.3. Calculate its [H 3 O + ]. [H 3 O + ] = antilog (-pH)= 10 –pH = 10 -5.3 [H 3 O + ] = 5 x 10 -6 1 significant figure Solution is acidic

40 40 Example 3: The [H 3 O + ] of a solution is 3.5 x 10 –3 M. Calculate its pH. pH = -log [H 3 O + ]= -log [3.5x10 -3 ]= -(-2.46) pH = 2.46 2 significant figures Solution is acidic

41 41 Example 4: The pH of tomato juice is 4.1. Calculate its [H 3 O + ]. [H 3 O + ] = antilog (-pH)= 10 –pH = 10 -4.1 [H 3 O + ] = 8 x 10 -5 1 significant figure Solution is acidic

42 42 Example 5: The [OH  ] of a cleaning solution is 1.0 x 10  5 M. What is the pH of this solution? [H 3 O + ] = = 1.0 x 10  9 M K w =[H 3 O + ][OH – ] pH =  log[H 3 O + ]=  log(1.0x10  9 ) = 9.00 2 sig figs Solution is basic

43 43 Example 6: The pH of a solution is 11.50. Calculate the [H 3 O + ] for this solution. [H 3 O + ] = antilog (-pH)= 10 –pH = 10 -11.50 [H 3 O + ] = 3.2 x 10 -12 Solution is basic

44 44 THE END


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