1422-Chapt-18-Acids-Bases

1422-Chapt-18-Acids-Bases - Chapters 18 & 19 Acids...

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Chapters 18 & 19 Acids & Bases H 2 O + H F ( aq ) H 3 O + ( aq ) + F ( aq )
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Acids & Bases 2 Acid and Base Definitions Arrhenius Acid Î increases H + concentration Base Î increases OH - concentration Brønsted-Lowry (1923) Acid Î donates a H + Base Î combines with or accepts H + Lewis Acid Î electron acceptor Base Î electron donor
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Acids & Bases 3 Water itself has some ionic character : H 2 O ( l ) + H 2 O ( l ) H 3 O + ( aq ) + OH ( aq ) Shorthand: H 2 O ( l ) H + ( aq ) + OH ( aq ) This is called self-ionization . Although most chemists simply write H + , it is important to realize that H + by itself represents a naked proton. In water the H + is hydrated, that is, it forms ionic-diople interactions with other water molecules. A common way of more accurately representing the fact that it is hydrated (interacting with waters) is to write H 3 O + .
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Acids & Bases 4 H 2 O ( l ) H + ( aq ) + OH ( aq ) + 2 eq OH H [[ K O ]] [ H ] = But H 2 O is a pure liquid and its concentration is constant ([ H 2 O ] = 55.6 M , activity = 1), so it is not included in the equilibrium expression: 14 + @ 20 C K[ ] [ ] K1 . 0 10 H × H O = = D w w [ H + ] [ OH ] = 1 × 10 14 [ H + ] = [ OH ] = 1 × 10 7 M So the [ H + ] and [ OH ] concentration in pure water is 1 × 10 7 M We use K w to indicate the water self- or auto- ionization . This is still a K eq or K c . Chemists use many subscripts on the equilibrium constant K to indicate specific types of equilibria: K a = acid equilibria K b = base equilibria K sp = slightly soluble (solubility product) equilibria Because of the often very small nature of H + concentrations, chemists (and others) have devised
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Acids & Bases 5 a logarithmic scale to simplify expressing these values: p H = log [ H + ] The negative sign in front of the log makes sure that most small concentrations of acid are given by positive values. [ H + ] = 1 × 10 7 M p H = log(1 × 10 7 ) = 7.0 The greater the [ H + ] concentration, the LOWER the p H value!!! p H < 7.0 Acidic p H = 7.0 Neutral p H > 7.0 Basic (Alkaline)
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Acids & Bases 6 Substance p H 10 M HCl 1.0 1 M HCl 0.0 Stomach Acid (HCl) 1.4 Lemon Juice 2.1 Orange Juice 2.8 Wine 3.5 Black Coffee 5.0 Urine 6.0 Pure Water 7.0 Blood 7.4 Baking Soda Solution 8.5 Ammonia Solution 11.9 1 M NaOH 14.0 10 M NaOH 15.0
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Acids & Bases 7 We can also define: p OH = log [ OH ] Although most chemists mainly use p H , p OH can be useful in base equilibrium numerical problems that we will run into later. Another definition we use is: p K w = log K w = 14 So, for a given water solution: p OH + p H = p K w = 14 or: p OH = 14 p H p H = 14 p OH
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Acids & Bases 8 Example: The p H of wine is 3.5. What is the [ H + ] ? What is the [ OH ] ? What is the p OH ? p H = log [ H + ] = 3.5 [ H + ] = antilog( 3.5) = 10 3.5 = 3.16 × 10 4 M H OH K [] 14 11 4 11 0 3.1 10 3.16 10 × == = × × + w p OH = log [ OH ] = log(3.1 × 10 11 ) = 10.5 ---- or ---- p OH = p K w p H = 14 3.5 = 10.5 Problem: The p H in your stomach is around 1. What is the [ H + ] ? What is the [ OH ] ? What is the p OH ?
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Acids & Bases 9 Dissociation Equilibrium Constants An acid is a compound that will ionize in solution (usually water) to form a H + ( aq ) and a counter- anion. This can be writen in a general fashion as: H A ( aq ) H + ( aq ) + A ( aq ) The equilibrium expression for this reaction is: [] [ A ] H K H [ ] A = + eq
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1422-Chapt-18-Acids-Bases - Chapters 18 &amp; 19 Acids...

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