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Acids & Bases. Key Characteristics of Acids & Bases Acids Taste sour Reacts with alkali metals (forms H2 gas) Litmus paper: RedNeutralizes Bases Bases.

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Presentation on theme: "Acids & Bases. Key Characteristics of Acids & Bases Acids Taste sour Reacts with alkali metals (forms H2 gas) Litmus paper: RedNeutralizes Bases Bases."— Presentation transcript:

1 Acids & Bases

2 Key Characteristics of Acids & Bases Acids Taste sour Reacts with alkali metals (forms H2 gas) Litmus paper: RedNeutralizes Bases Bases Tastes bitterSlippery feelLitmus paper: BlueNeutralizes 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 Example: NaOH  Na + + OH -

4 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, H 3 O +  Oftentimes H + and H 3 O + are used interchangeably HCl  H + + Cl - HCl (g) + H 2 O (l)  H 3 O + (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 NH 3 for example is a base  Brønsted and Lowry proposed a definition based on acid base reactions transferring H + ion from one substance to another  Arrhenius theory has limitations:  Only deals with aqueous solutions (solutions in water)  Not all acids and bases produce H+ and OH- ions NH 3 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  Br ø nsted-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 + H 2 O  H 3 O + + Br - A B

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

10 Brønsted-Lowry Theory  These pairs differ by only one hydrogen ion  Example  Identify the Brønsted-Lowry acid, base, conjugate acid and conjugate base NH 3 + H 2 O  NH OH - B A CA CB  NH 3 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) + H 2 O (l) ↔ H 3 O + (aq) + Cl - (aq) HSO HCO 3 - ↔ SO H 2 CO 3

12 Theories of Acids & Bases  Lewis Acids & Bases  Acids are electron acceptors  Bases are electron donors  Example: H 2 O + NH 3  OH - + NH 4 +  Is really: H 2 O + :NH 3  OH - + H:NH 3 + Electron pair donor(NH 3 ) Electron pair acceptor(H + )

13 Summary Of Theories Acids release H + Bases release OH- Defines acids & bases in H 2 O Arrhenius Acids – proton donor Bases – proton acceptor Can define acids & bases in solvents other than H 2 O Brønsted- Lowry Acids – electron acceptor Bases – electron donor Defines acids & bases without a solvent Lewis

14 The Self-Ionization of Water  Even pure water contains a small number of ions: H 2 O (l) ↔ H 3 O + (aq) + OH - (aq)  In pure water, the concentrations of the ions (H 3 O + and OH - ) are equal. [H 3 O + ]=[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 1x  this is referred to as the ion product constant of water  This ion product constant of water is given the symbol K w

16 The Self-ionization of water  Example #1  What is the H 3 O + concentration in a solution with [OH - ] = 3.0 x M ? K w = [H 3 O + ][OH - ] 1x = [H 3 O + ][3.0x10 -4 ]

17 Example #2  If the hydroxide-ion concentration of an aqueous solution is 1.0 x M, what is the [H 3 O + ] in the solution? K w = [H 3 O + ][OH - ] 1x = [H 3 O + ][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 (H 3 O + ) ions  A pH = 2.0 has 10 times the concentration of H 3 O + 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 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

21 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 [H 3 O + ] K w = [H 3 O + ][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? [H 3 O + ] = 10 -pH

24 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? [H 3 O + ] = 10 -pH K w = [H 3 O + ][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: 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!

27 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

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

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

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 Perchloric acid, HClO4 Chloric acid, HClO 3 Hydrochloric acid, HCl Hydrobromic acid, HBr Hydroiodic acid, HI Nitric acid, HNO 3 Sulfuric acid, H 2 SO 4 Strong Acids  Must be memorized!

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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