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Acids & Bases.

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Presentation on theme: "Acids & Bases."— Presentation transcript:

1 Acids & Bases

2 Key Characteristics of Acids & Bases
Taste sour Reacts with alkali metals (forms H2 gas) Litmus paper: Red Neutralizes Bases Bases Tastes bitter Slippery feel Litmus paper: Blue Neutralizes Acids

3 Theories of Acids & Bases
Arrhenius Theory of Acids & Bases Properties of acids are due to the presence of H+ ions Example: HCl  H+ + Cl- Properties of bases are due to the presence of OH- ions NaOH  Na OH-

4 HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)
H+ ions in water H+ ions are bare protons These H+ ions react strongly with the nonbonding pair of electrons in a water molecule This forms the hydronium ion, H3O+ Oftentimes H+ and H3O+ are used interchangeably HCl  H+ + Cl- HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

5 Problems with Arrhenius
Arrhenius theory has limitations: Only deals with aqueous solutions (solutions in water) Not all acids and bases produce H+ and OH- ions NH3 for example is a base Brønsted and Lowry proposed a definition based on acid base reactions transferring H+ ion from one substance to another

6 Brønsted-Lowry Theory

7 Theories of Acids & Bases
Brønsted-Lowry Theory Acids are substances that donate H+ ions Acids are proton donors Bases are substances that accept H+ ions Bases are proton acceptors Example: HBr + H2O  H3O Br- A B

8 Brønsted-Lowry Theory
The behavior of NH3 can now be understood: NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq) Since NH3 becomes NH4+, it is a proton acceptor (or a Brønsted-Lowry base) H2O becomes OH-, which means it is a proton donor (or a Brønsted-Lowry acid)

9 Brønsted-Lowry Theory
Conjugate Acid-Base Pairs An acid and a base that differ only in the presence or absence of H+ are called a conjugate acid-base pair. Every acid has a conjugate base. Every base has a conjugate acid. HX is the conjugate acid of X- H2O is the conjugate base of H3O+

10 Brønsted-Lowry Theory
These pairs differ by only one hydrogen ion Example Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH3 + H2O  NH OH- B A CA CB NH3 acts as a Brønsted base by accepting a proton. Water acts as a Brønsted acid by donating a proton.

11 Brønsted-Lowry Theory
Example HCl (g) + H2O (l) ↔ H3O+(aq) + Cl- (aq) HSO HCO3- ↔ SO H2CO3 A B CA CB A B CB CA

12 Theories of Acids & Bases
Lewis Acids & Bases Acids are electron acceptors Bases are electron donors Example: H2O + NH3  OH- + NH4+ Is really: H2O + :NH3  OH- + H:NH3+ Electron pair donor(NH3) Electron pair acceptor(H+)

13 Summary Of Theories Acids release H+ Bases release OH-
Arrhenius Acids release H+ Bases release OH- Defines acids & bases in H2O Brønsted-Lowry Acids – proton donor Bases – proton acceptor Can define acids & bases in solvents other than H2O Lewis Acids – electron acceptor Bases – electron donor Defines acids & bases without a solvent

14 The Self-Ionization of Water
Even pure water contains a small number of ions: H2O (l) ↔ H3O+ (aq) + OH- (aq) In pure water, the concentrations of the ions (H3O+ and OH-) are equal. [H3O+]=[OH-]= 1x10-7 M

15 The Self-ionization of Water
Writing the equilibrium expression for the self- ionization of water gives: Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x10-14 this is referred to as the ion product constant of water This ion product constant of water is given the symbol Kw

16 The Self-ionization of water
Example #1 What is the H3O+ concentration in a solution with [OH-] = 3.0 x 10-4 M? Kw = [H3O+][OH-] 1x10-14 = [H3O+][3.0x10-4]

17 Example #2 If the hydroxide-ion concentration of an aqueous solution is 1.0 x 10-3 M, what is the [H3O+] in the solution?   Kw = [H3O+][OH-] 1x10-14 = [H3O+][1.0x10-3]

18 The pH scale Developed by Søren Sørensen in order to determine the acidity of ales Used in order to simplify the concept of acids and bases The pH scale goes from 1 to 14 A change in one pH unit corresponds to a power of ten change in the concentration of hydronium (H3O+) ions A pH = 2.0 has 10 times the concentration of H3O+ than a pH = 3.0, and 100 times greater than pH = 4

19 The pH scale pH < 7 Acid pH = 7 Neutral pH > 7 Base

20 pH = -log [H3O+] or [H3O+] = 10-pH
Calculations of pH pH can be expressed using the following equation: pH = -log [H3O+] or [H3O+] = 10-pH Example #1 What is the pH of a solution with M H3O+? Is this solution an acid or a base? Acid

21 pH = -log [H3O+] Kw = [H3O+][OH-]
Calculations of pH Example #2 What is the pH of a solution with the concentration of hydroxide ions M? Is this an acid or a base? pH = -log [H3O+] Kw = [H3O+][OH-] Base

22 Calculations of pH Practice #1 Practice #2

23 Calculations of pH Example #1
What is the hydronium ion concentration in fruit juice that has a pH of 3.3? [H3O+] = 10-pH

24 [H3O+] = 10-pH Kw = [H3O+][OH-]
Calculations of pH What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05? [H3O+] = 10-pH Kw = [H3O+][OH-]

25 Calculation of pH Practice #1 Practice #2

26 Strength of Acids & Bases
When a solution is considered strong, it will completely ionize in a solution Nitric acid is an example of strong acid: HNO3 (l) + H2O (l)  NO3- (aq) + H3O+ (aq) In a solution of nitric acid, no HNO3 molecules are present Strength is NOT equivalent to concentration!

27 Strength of Acids & Bases
Knowing the strength of an acid is important for calculating pH If given concentration of strong acid (such as HNO3) assume it is the same as the concentration of hydronium, H3O+, ions Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH-, ions

28 Strong Acids & Bases Ionize 100%
Example NaOH  Na OH- 1 M 1 M 1 M OH- Na+ Na+ Na+ OH- OH-

29 Weak Acids & Bases Ionize X%
Example HF  H F- 1 M ? M ? M F- HF H+ H+ HF H+ F- F-

30 Strength of Acids & Bases
Stronger the acid Weaker the conj. base Stronger the base Weaker the conj. acid

31 Strength of Acids & Bases

32 Strong Acids Must be memorized! Strong Acids Perchloric acid, HClO4
Hydrochloric acid, HCl Hydrobromic acid, HBr Hydroiodic acid, HI Nitric acid, HNO3 Sulfuric acid, H2SO4

33 Strong Acids 6 of 7 strong acids are monoprotic (HX)
Exists only as H ions and X ions HI(aq) H+(aq) + I-(aq) 2M HI = [H+]= [I-] = 2M Determining pH of Strong Acids For Strong Acids: pH = -log [H+] For monoprotic strong acids: [H+] = [X]

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