VI Solutions D Acid Base

VI Solutions D Acid Base - Solutions D Acids and Bases...

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Solutions D Acids and Bases Bronsted-Lowry concept
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Bronsted-Lowry concept of acids and bases more general definition of acids and bases Acid any substance that can donate an H + Base any substance that can accept an H +
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HA --> H + + A - A - + H + --> HA B + H + --> HB + HB + --> H + + B So: Bronsted acids and bases occur in pairs Acid - Base pairs are called Conjugate Pairs Acid(HA) and conjugate base(A - ) Base(B) and conjugate acid(HB + ) When acid and base react, they react in pairs. Note: H + and H 3 O + mean the same thing
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CH 3 COOH + H 2 O CH 3 COO - + H 3 O + A 1 A 2 B 1 B 2 B 1 B 2 A 1 A 2 NH 3 + H 2 O NH 4 + + OH - Note that H 2 O can act as an acid or base depending on the other reactant. H 2 O is AMPHIPROTIC
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Strong Acids(Acid Strength) the stronger the acid, the easier it gives up H + . Reactions proceed strong to weak. Strong Bases(Base Strength) the stronger the acid, the weaker the conjugate base. So, base strength is usually determined by the strength of the conjugate acid.
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The reaction: HNO 2 + CH 3 COO - CH 3 COOH+ NO 2 - is displaced to the right (that is more products than reactants) Since reaction proceed strong to weak, HNO 2 is stronger than CH 3 COOH And CH 3 COO - is stronger than NO 2 - .
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pH and pOH pH scale In aqueous solution, [H 3 O + ] and [OH - ] are generally small, less than 1 M. Since most of the time scientists were interested in how acidic the solution was, and wanted a way to easily assess the information, the pH scale was developed. (power of hydrogen) pH = - log[H + ] pH = - log[H 3 O + ]
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Since we now have a defined function (-log[X]) we can now use it for other terms pOH = - log[OH - ] [H + ] = 4 x 10 -6 pH = - log(4 x 10 -6 ) pH = 5.6
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Equilibrium the reaction proceed to a point where the concentration of reactants and products do not change we can set up a relationship showing the relationship. K is called an equilibrium constant, and
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VI Solutions D Acid Base - Solutions D Acids and Bases...

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