Acids Bases and Salts Acid ph less than 7.0 Sour taste Electrolyte
Names of Acids Binary acids Contain only 2 elements Begin with hydro; end with ic Ternary acids Ex: H 2 S = hydrosulfuric Contain a polyatomic ion End with ic if ion ends with ate H 2 SO 4 = sulfuric End with ous if ion ends with ite Ex: H 2 SO 3 = sulfurous
Base ph greater than 7.0 Slippery Less common electrolytes bitter
Neutralization The reaction between an acid and a base Produces water and a salt Water can act as an acid or a base Amphoteric or amphiprotic
Arrhenius Definitions Acid substance that produces hydrogen ions when in solution HCl H 3 O + + Cl - Base substance that produces hydroxide ions when in solution NaOH Na + + OH -
Bronsted Lowrey definitions Acid any substance that donates a proton in a chemical reaction HCL + H 2 O Cl - + H 3 O + the HCL molecule donates a hydrogen ion or free proton that combines with a nearby water molecule to form the hydronium ion H 3 O + Base any substance that accepts a proton as a result of a chemical reaction H 3 O + + Cl - H 2 O + HCl the chloride ion accepts the proton to form HCl
Conjugate acid the particle formed when the base acquires a proton from the acid Base + H+ = conjugate acid Conjugate base the particle that remains after an acid donates a proton Acid H+ = conjugate base
Typical reaction H 3 O + + Cl - H 2 O + HCl Hydronium is the acid Chloride is the base Water is the conjugate base HCl is the conjugate acid
Lewis Definitions Acid an electron pair acceptor Base an electron pair donor H 3 N: + H 2 O H 4 N + + OH - The electron pair is accepted by the hydrogen ion from the water molecule water is the acid the electron pair is donated by the ammonia molecule making ammonia the base
NH 3 + H 2 O NH 4 + + OH - Notice that ammonia is a base by all three definitions It forms hydroxide ions in the reaction It is a proton acceptor It is an electron pair donor
Anhydrides Anhydrous means without water anhydrides are simply acids and bases that have had water removed Acidic anhydrides (acid without water) Nonmetal oxides; SO 3, N 2 O5 SO 2 + H 2 O H 2 SO 3 ; Sulfur dioxide is the acidic anhydride of sulfurous acid Basic anhydrides (base without water) Metal oxides; MgO, Cr 2 O 3, Na 2 O Na 2 O + H 2 O 2NaOH; sodium oxide is the basic anhydride of sodium hydroxide
Salts Ionically bonded compounds Produced by a neutralization reaction Positive ion from the base, negative ion from the acid Dissociates in water to form electrolyte Also produced in reactions between anhydrides and acids or bases
Hydrolysis Reaction of a salt with water to form an acidic or basic solution Salts do not have to be neutral If it has an ionizable H+, it will be acidic Strong acid + strong base = neutral salt ph 7; [H + ]= [OH - ] Strong acid + weak base = acidic salt ph < 7; [H+] >[OH-] Weak acid + strong base = basic salt ph > 7; [H+] < [OH-] Weak acid + weak base =??
Is the Salt acidic or basic? Look at the parent acid and parent base Determine strength of parents Determine acidity/basicity of the salt CaCl 2 Parent acid HCl (add H to ion) strong Parent base Ca(OH) 2 (add OH to + ion) strong Salt is neutral
Strengths of acids The strength of an acid or a base is related to the degree to which the ions dissociate in solution Strong acids/bases are 100% ionized no equilibrium Acid strength increases across a period as electronegativity increases; decreases down a family as bond strength decreases The concentration of the acid solution you have is not related to its strength Concentration is measure of molar volume Weak acids/bases dissociate only slightly and set up an equilibrium
Binary acids The more electronegative the anion, the stronger the acid Strong binary acids HI HBr HCl
Ternary acids The strength of ternary acids is dependant upon the polyatomic ion in the acid Find the difference between the number of oxygens in the polyatomic and the number of hydrogens If the difference is 2 or greater it is a strong acid If the difference is <2 it is a weak acid
Bases Bases containing hydroxide and a group one metal are strong Calcium hydroxide and magnesium hydroxide are also strong bases All others are weak
More examples Na 3 PO 4 Parent acid H 3 PO 4 weak Parent base NaOH strong Salt is basic CuCl 2 Parent acid HCl strong Parent base Cu(OH) 2 weak Salt is acidic
Net ionic equations Reactions that take place in aqueous solution react more often as individual ions rather than as the molecules depicted in the chemical equation Ex: HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l) An ionic equation better represents what is actually occurring H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) Na + (aq) + Cl - (aq) + H 2 O (l)
Net Ionic cont. Inspection of the equation H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) Na + (aq) + Cl - (aq) + H 2 O (l) Shows that only some ions participate directly in the reaction Na+ and Cl- are present both before and after We call these spectator ions Only H+ and OH- ions partcipate directly, forming water H + (aq) + OH - (aq) H 2 O (l) This is the net ionic equation
Precipitation Reactions Many reactions will produce a precipitate. The net ionic reaction can help you determine what the precipitate is Solubility review Soluble nitrates, salts of group 1 and NH 4+, chlorides (except Ag, PbII, Hg); sulfates (except Ba, PbII Ca) Not soluble hydroxides (except Na, K, Ca), sulfides, carbonates, phosphates
Ex: reactions Silver nitrate and sodium chloride are mixed in aqueous solution forming a white precipitate. Write the net ionic equation and predict what the precipitate is. AgNO 3(aq) + NaCl (aq) AgCl + NaNO 3 Ag + (aq) + NO 3(aq) + Na + (aq) + Cl - (aq)?? Which product is insoluble?
Ag + (aq) +NO 3(aq) +Na (aq) + Cl - (aq) Na + (aq) + NO 3 - (aq) + AgCl (s) Na + and NO 3 - are present on both sides so Net ionic equation is. Ag + (aq) +Cl - (aq) AgCl (s)
Self ionization of water H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH - (aq) This leads to the equilibrium expression Kw = [H 3 O+][OH-] Experiments show that at 25 C [H + ] = [OH - ] = 1.0 x 10-7 mol/l K w = [H 3 O + ][OH - ] K w = (1.0 x 10-7 mol/l)(1.0 x 10-7 mol/l) K w = (1.0 x 10-14 mol 2 /L 2 ) The units are customarily omitted
Acidity of solutions Kw tells us that the product of [H 3 O+] and [OH-] must always be 1.0 x 10-14 This is the ion product constant for all dilute aqueous solutions These solutions exist in equilibrium Adding acid to the solution raises the [H 3 O+] thus lowering the [OH-] Adding base to the solution raises the [OH-] thus lowering the [H 3 O+]
Examples What is the [OH - ] if the [H 3 O + ] is 1.0 x 10-4 K w = [H 3 O + ] [OH - ] 1.0 x 10-14 = [OH - ] (1.0 x 10-4 ) 1.0 x 10-14 = [OH - ] (1.0 x 10-4 ) (1.0 x 10-10 ) = [OH - ]
Calculating ph The ph scale is based on the % of hydrogen (or hydronium) ions in a solution Since the concentration can vary over a wide scale of values (from 1 x 10-1 to 1 x 10-14 ), chemists report the concentration in terms of ph The negative logarithm (to the base 10) of the [H 3 O + ] ph = -log [H 3 O + ]
ph of Solutions ph = 7 neutral; [H 3 O + ] = [OH - ] ph < 7 acidic; [H 3 O + ] > [OH - ] ph > 7 basic; [H 3 O + ] < [OH - ] (alkaline) [H 3 O + ] = 1.0 x 10-5 ph = -log (1.0 x 10-5 ) ph = 5
Concentration of Solutions We can use the definition of log to reverse the calculation if we know the ph and want to find the [H 3 O + ] 10 ph = [H 3 O + ] If ph = 3.2 10-3.2 = [H 3 O + ] 6.3 x 10-4 = [H 3 O + ]
Equilibrium of Acid/Base Solutions Consider the dissociation of HCl HCl + H 2 O H 3 O + + Cl - The equilibrium expression for this would be K eq = [H 3 O + ][Cl - ] [HCl ] [H 2 O] Since water is a constant it can be combined with K eq to give the acid dissociation constant K a K a = [H + ][Cl - ] [HCl ]
Equilibrium of Acid/Base Solutions The equilibrium of the dissociated ions to the parent acid is expressed as the K a and for the base K b K a = [H + ] [A - ] [HA] K b = [M + ] [OH - ] [MOH] The value of K a /K b like K eq, tells us which is favored, reactants or products
Ionization of acids/bases Ka/Kb reflect the fraction of the acid or base that has been ionized The larger the value the greater the dissociation, the stronger the acid/base Polyprotic acids (more than one H) lose their hydrogens one at a time and have a separate dissociation constant for each
Equilibrium of a weak acid Consider a 1.00 M solution of HF K a for HF is 7.2 x 10-4 7.2 x 10-4 = [H + ] [F - ] [HF] 7.2 x 10-4 = [x] [x] 1.00 M 7.2 x 10-4 = x 2 X = 2.7 x 10-2 1.00 M
Titration Method to experimentally determine the concentration of an acidic or basic solution An acid of unknown concentration can be determined from a base with a known concentration A base of unknown concentration can be determined from an acid with a known concentration
Indicators Substance that is one color in an acidic solution and a different color in a basic solution Litmus Base blue Acid - red Phenolphthalein Base pink/purple Acid - colorless Bromthymol blue Base blue/green Acid - yellow
Titration Calculations (#H + ) (C a ) (V a ) = (C b ) (V b ) (#OH - ) EX: If 25.0 ml of a sulfuric acid solution is titrated to a neutral endpoint using 53.2 ml of 1.2M sodium hydroxide, what is the concentration of the sulfuric acid? (2)(C a )(25.0ml) = (1.2M)(53.2 ml) (1) (C a )(50.0) = 63.84 (C a ) = 1.3M