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ACIDS AND BASES ChemistryMs. Piela. Key Characteristics of Acids & Bases Acids Taste sour Reacts with alkali metals (forms H2 gas) Forms electrolyte solutions.

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Presentation on theme: "ACIDS AND BASES ChemistryMs. Piela. Key Characteristics of Acids & Bases Acids Taste sour Reacts with alkali metals (forms H2 gas) Forms electrolyte solutions."— Presentation transcript:

1 ACIDS AND BASES ChemistryMs. Piela

2 Key Characteristics of Acids & Bases Acids Taste sour Reacts with alkali metals (forms H2 gas) Forms electrolyte solutions (conducts electricity) pH paper color: Red Neutralizes Bases Bases Tastes bitterSlippery feel Forms electrolyte solutions (conducts electricity) pH paper color: Blue Neutralizes Acids

3 The 3 Main Theories of Acids/Bases This course will mainly deal with BL theory

4 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 Example: NaOH  Na + + OH -

5 What is an H + ?  H + ions are bare protons  These are so reactive that they do not exist naturally, but will bond with water to form a hydronium ion, or H 3 O + ion  Oftentimes H + and H 3 O + are used interchangeably HCl  H + + Cl - HCl (g) + H 2 O (l)  H 3 O + (aq) + Cl - (aq)

6 Problems with the Arrhenius theory  Only deals with aqueous solutions (solutions in water)  Not all acids and bases contain H+ and OH- ions Example: NH 3 is a base Considered the most incomplete theory of acids and bases  Only deals with aqueous solutions (solutions in water)  Not all acids and bases contain H+ and OH- ions Example: NH 3 is a base Considered the most incomplete theory of acids and bases

7 Theories of Acids & Bases  Brønsted-Lowry Theory of Acids & Bases  Acids are substances that donate H + ions Acids are proton (H + ) donors  Bases are substances that accept H + ions Bases are proton (H + ) acceptors  Example: HBr + H 2 O  H 3 O + + Br - A B

8 Brønsted-Lowry Theory  The behavior of NH 3 can be understood now: NH 3 (aq) + H 2 O (l) ↔ NH 4 + (aq) + OH - (aq)  NH 3 becomes NH 4 +, so NH 3 is a proton acceptor (or a Brønsted-Lowry base)  H 2 O becomes OH -, so H 2 O is a proton donor (or a Brønsted-Lowry acid)

9 Brønsted-Lowry Theory

10

11  Example Problems  Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH 3 + H 2 O  NH OH -

12 Brønsted-Lowry Theory  Example HCl (g) + H 2 O (l) ↔ H 3 O + (aq) + Cl - (aq) HSO HCO 3 - ↔ SO H 2 CO 3

13 Theories of Acids & Bases  Lewis Acids & Bases  Acids are electron acceptors  Bases are electron donors  Amphoteric – substances that can act as both an acid and a base  Examples: H 2 O, HCO 3 -

14 Summary Of Theories Acids release H + Bases release OH- Arrhenius Acids – proton donor Bases – proton acceptor Brønsted- Lowry Acids – electron acceptor Bases – electron donor Lewis

15 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 for his workers  The pH scale goes from 0 to 14  The acidity/basicity of the solutions depends on the concentration of H + (or H 3 O + )

16 The pH scale

17 pH scale  Low pH values means a high concentration of H + (acidic)  High pH values means a low concentration of H + (basic)

18 Calculations of pH  The Self Ionization of Water  In pure water (pH = 7), the concentrations of the ions (H 3 O + and OH - ) are equal. [H 3 O + ]=[OH - ]= 1x10 -7  This is because water will spontaneously dissociate naturally: H 2 O (l) ↔ H 3 O + (aq) + OH - (aq)  Writing the equilibrium expression for the self-ionization of water gives:

19 The Self-ionization of Water  Plugging in the concentrations in pure water, this gives an equilibrium constant of 1x  This is referred to as the ion product constant of water  The ion product constant of water has its own symbol: K w Unlike other equilibrium constants, the K w will always be the same value

20 Calculations of H 3 O + /OH -  Example #1  What is the H 3 O + concentration in a solution with [OH - ] = 3.0 x M ? K w = [H 3 O + ][OH - ] 1 x = [H 3 O + ][3.0 x ] _________________________ 3.0 x 10 -4

21 Calculations of H 3 O + /OH -  If the hydronium-ion concentration of an aqueous solution is 1.0 x M, what is the hydroxide ion concentration in the solution? K w = [H 3 O + ][OH - ] 1 x = [1 x ][OH - ] _________________________ 1.0 x 10 -3

22 Calculations of pH  pH can be expressed using the following equation: pH = -log [H 3 O + ] or [H 3 O + ] = 10 -pH  Example #1  What is the pH of a solution with M H 3 O + ? Is this solution an acid or a base? Acid!

23 Calculating pH of a solution  Example #2  What is the pH of a solution where the concentration of hydroxide ions is M? Is this an acid or a base? K w = [H 3 O + ][OH - ]pH = -log [H 3 O + ] Base!

24 Calculating pH of a solution  Practice #1  Practice #2

25 Calculating H 3 O + /OH - from pH  Example #1  What is the hydronium ion concentration in fruit juice that has a pH of 3.3? [H 3 O + ] = 10 -pH

26 Calculating H 3 O + /OH - from pH  What are the concentrations of the hydronium and hydroxide ions in a sample of rain that has a pH of 5.05? [H 3 O + ] = 10 -pH K w = [H 3 O + ][OH - ]

27 Calculating H 3 O + /OH - from pH  Practice #1  Practice #2

28 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: HNO 3 (l) + H 2 O (l) ⇋ NO 3 - (aq) + H 3 O + (aq)  In a solution of nitric acid, no HNO 3 molecules are present!  Strength is NOT equivalent to concentration!

29 Strength of Acids & Bases  Knowing the strength of an acid is important for calculating pH  If given concentration of strong acid (such as HNO 3 ) assume it is the same as the concentration of hydronium, H 3 O +, ions  Given concentration of a strong base, assume it has the same concentration as the hydroxide, OH -, ions

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

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

32 Naming Bases  Bases are soluble metal hydroxides  Follow identical naming rules for ionic compounds  Examples  NaOH  Ba(OH) 2  NH 3  NH 4 +

33 Naming Acids  Binary Acids (HX)  If the acid has an anion that ends in “-ide” use the following basic format to name the acid: “Hydro – root – ic acid”  Example HCl

34 Naming Acids  Example  HBr  Practice  HI H2SH2S

35 Naming Acids PPolyatomic acids (aka oxoacids, H x A y O z ) NName depends on the polyatomic used: If polyatomic ends in “-ite”, replace with “ous acid” If polyatomic ends in “-ate”, replace with “ic acid” Trick: “I ate something icky”

36  Examples  HClO 4  HClO 2  Sulfuric acid


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