Properties of Acids. Base Chemistry. Properties of Bases. Three Acid and Base Theories. 1) Arrhenius Theory. May 09, Naming Acids Review

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1 May 09, 2013 Properties of Acids AcidAcid Base Chemistry Base Chemistry Taste sour Are strong or weak electrolytes React with bases to form water and salts React with active metals to produce H2 Turn litmus (and cabbage) red Low ph Acid Rxns with Metals 1. Al, Mg and Zn form hydrogen gas 2Al + 6HCl 2AlCl3 + 3H2 2. Metal carbonates form carbon dioxide CaCO3 + 2HCl CaCl2 + H2O + CO2 Chapter 19 Properties of Bases Naming Acids Review Anion Ending Example Acid Name -ide chloride hydro-(stem)-ic -ite sulfite (stem)-ous -ate nitrate (stem)-ic Example Three Acid and Base Theories 1. Arrhenius Theory 2. Bronsted-Lowry Theory 3. Lewis Theory Taste bitter Are strong or weak electrolytes React with acids to form water and salts Feel slippery Turn litmus (and cabbage) blue High ph 1) Arrhenius Theory Simplest definition and most restrictive Acids ionized to produce H and Bases ionized to produce hydroxide ions OH + -

2 Examples of Arrhenius An acid in water ionizing to form a H+ ion 2) Bronsted-Lowry Theory Added to Arrhenius definition of bases HCl (aq) H + (aq) + Cl - (aq) Bronsted-Lowry acid: molecule or ion that is a proton donor A base in water dissociating to form hydroxide ion, OH - NaOH Na + (aq) + OH - (aq) Bronsted-Lowry base: molecule or ion that is a proton acceptor. Bases don't necessarily have to supply OH - NH 3, ammonia, can now be recognized as a base Bronsted-Lowry Theory: General acid/base reaction Acid + Base Conjugate base + Conjugate Acid Conjugate base: particle that remains after a proton that is released by the acid. Conjugate acid: particle that remains after a base has acquired a proton from the acid. Examples Identify the conjugate pairs. a) H2SO4(aq) + H2O(l) HSO4 - (aq) + H3O + (aq) b) H2O(l) + F - (aq) OH - (aq) + HF(aq) 3) Lewis Theory The broadest definition of acids and bases Definitions: Acid: electron-pair acceptor Doesn't necessarily have to supply (H + ) or be a proton donor Base: electron-pair donor Doesn't necessarily have to supply (OH-) or be a proton acceptor

3 Type Acid Base Arrhenius H + or H 3 O + producer OH- producer Example Bronsted- Lowry proton (H+) donor proton (H+) acceptor BF 3(g) + NH3(g) F3BNH3(g) Lewis electron-pair acceptor electron-pair donor Amphoteric Compounds Any substance that can behave as an acid or as a base. The most common example is water. Monoprotic and Polyprotic Acids Monoprotic: an acid that can only donate one proton per molecule. Polyprotic: is an acid that can donate more than one proton per molecule. (Diprotic can donate two protons.) (Triprotic can donate three protons.) Ionization of polyprotic acids Protons can only be donated one at a time to water. As a result each hydrogen ion will leave the acid in its own reaction step. The first proton leaves in the first ionization, the second proton is removed by the second ionization, etc.) One proton=one reaction step Two protons=two reaction steps Three protons=three reaction steps HCl H 2 SO 4 Monoprotic, Diprotic and Triprotic An acid that contains "x" number of hydrogens/molecule able to dissociate in water.

4 Strong Acids and Bases Strength in acids and bases does NOT refer to how corrosive or dangerous they are Strong acids and bases dissociate (break apart) completely in water. At equilibrium, there is no acid or base left that has not ionized (separated into its ions). In contrast, one of the most corrosive and dangerous acids, HF, is actually a weak acid. Acid Base Strength Strong acids and bases will dissociate completely in aqueous solution. Reaction equations will show a one way arrow and there will be no equilibrium Good conductors of electricity Weak acids and bases do no ionize completely in aqeous solution. Reaction will show a two way arrow and equilibrium will be established. Not good conductors of electricity. Acid Base Concentration Concentration refers to the amount of solute in a given amount of solution Concentrated Acid-more solute in a given amount of solution than another solution of the same volume Dilute Acid-less solute in a given amount of solution than another solution of the same volume. Equilibrium Constants Law of Chemical Equilibrium-at a given temperature, at equilibrium, a ratio of reactant to product concentrations has a constant value. Equilibrium constants (K eq )- the numeric value that compares reactant and product concentrations Equilibrium constant expressions can only be written for reactions at equilibrium and solids and liquids are always excluded from the equilibrium constant expression.

5 Equilibrium Constants If the value of K eq is <1, there is more reactant at equilibrium If the value of K eq is >1, there is more product at equilibrium. Equilibrium Constant Expression: aa + bb cc + dd > A, B - reactants > C, D - products > a, b, c, d- coefficients of the balanced equation 2H 2 S (g) Example 2H 2 (g) + S 2 (g) Write the equilibrium constant expression. If [H 2 S]= 0.184M, [H 2 ]=0.0377M and [S 2 ]=0.0540M, calculate the value of K eq. Does the reaction favor the production of reactants or products? Acid Ionization Constant (K a ) K a is the value of the equilibrium constant expression for the ionization of a weak acid. The base ionization constant (K b ) is the value of the equilibrium constant expression for a weak base. Large value of K a or K b =stronger acid or base Small value of K a or K b =weaker acid or base Ionization Constants K a : Quantitative expression for the strength of an acid K b : Quantitative expression for the strength of a base K w : Ionization constant for water Kw=[H3O + ][OH - ]=1.0Χ10-14 Ion Product Constant for Water (K w ) We can relate the three values to find either the Ka or Kb value. K w =K a K b K w =[H + ][OH - ]

6 Self Ionization of Water Hydronium Ion The H + is so small it immediately attaches to one of the unshared pair of electrons on a water molecule forming H 3 O + Studies have shown that hydrogen "bonds" form with other water molecules to form even larger complexes. As far as we're concerned, we'll symbolize the hydronium ion Example: at 298K an aqueous solution has a [H + ] of 1.0 x 10-5 M. What is the [OH - ] concentration? > H + or H 3 O + > We will use these interchangeably ph and poh calculations

7 ph values into perspective Because the ph scale is a log scale, changes by one ph unit are changes by a factor of 10. Assume that a ph=1 solution, which [H + ] =1x10-1 is similar in size as the length of a sports field. As ph increases, the [H + ] decreases by the same factor as the comparative scale. ph [H+] Comparison 1 1 x yards 2 1 x yards 3 1 x yard 4 1 x inches 5 1 x inches 6 1 x inches 7 1 x inches ph ph is a number scale from 0-14 that represents the acidity or basicity of a solution ph is a measure of the hydrogen/hydronium ion concentration of a solution ph=-log[h + ] the lowercase p stands for power ph and poh table with H + and OH - values ph H + OH - 1x10-7 1x10-5 poh ph + poh=14 [H + ][OH - ]=1x x10-5 Example ph Honors: See weak acid/base ice table! 1. Find the ph of a solution with a 7.01Χ10-6 M H3O + concentration. [H3O + ]= 7.01Χ10-6 M ph= -log[h3o + ] ph= -log[7.01χ10-6 M] ph= Find the ph of a solution with Χ10-4 M H3O + concentration.

8 Measuring ph: Indicators 1. Indicators are organic bases and acids whose colors differ from their conjugate acids or bases. 2. Many organic substances can be made into solutions that can be used to identify acids and bases. 3. Example: Red cabbage will be red if it is acidic and green if it is basic. Transition Interval A transition interval is the ph range at which the indicator is changing color. Not every indicator will work in the same ph range. ph meter: Makes a rapid, accurate ph measurement. This instrument is usually easier to read than liquid or ph indicator strips. Neutralization/Titration Titration: a way to determine the amount of acid/base needed to completely neutralize a substance. (Honors: More notes to follow) Buffers A buffer system is a solution that can absorb moderate amounts of acid or base without a significant change in its ph Example: Enzymes Blood: buffer of carbonic acid (H 2CO 3) and bicarbonate (HCO 3 -) to maintain a ph between 7.35 and 7.45.

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