2 Topics we ll be looking at in this chapter Arrhenius theory of acids and bases Bronsted-Lowry acid-base theory Mono-, di- and tri-protic acids Strengths of acids and bases Ionization constants for acids and bases Salts Acid-base neutralization reactions Self-ionization of water ph pk a and acid strength ph of aqueous salt solutions Buffers The Henderson-Hasselbach equation Electrolytes Equivalents and milliequivalents of electrolytes Acid-base titrations
3 Arrhenius theory of acids and bases Arrhenius acids are substances that increase the concentration of H + (or H 3 O + ) when dissolved in water H 2 O HCl (g) H + (aq) + Cl - (aq) H 2 O HNO 3(l) H + (aq) + NO 3 - (aq) Recognize acid formulas: H is at the beginning of the formula
4 Arrhenius theory of acids and bases When acids and bases are dissolved in water, they ionize (break apart into their constituent ions) Ionization is a process in which individual positive and negative ions are produced from a molecular compound that is dissolved in solution The acids listed are all molecular compounds. Acids ionize when they are dissolved in water like ionic compounds Most molecular compounds don t ionize. The exceptions are acids and bases.
5 Arrhenius theory of acids and bases Arrhenius bases are hydroxide (OH - ) containing substances that increase the concentration of OH when dissolved in water H 2 O NaOH Na + + OH - H 2 O Ca(OH) 2 Ca OH - Arrhenius bases contain hydroxide (OH - ) in their formulas
6 Arrhenius theory of acids and bases In contrast to Arrhenius acids, Arrhenius bases are ionic compounds. When bases (and salts) are dissolved in water, they dissociate. Dissociation is the process by which individual positive and negative ions are released from an ionic compound that is dissolved in water
7 Bronsted-Lowry acid-base theory Arrhenius theory is limited to aqueous solutions. Bases are limited to hydroxidecontaining compounds which ionize in water NH 3 also produces OH - ions when dissolved in water but by the Arrhenius definition, it is not a base Bronsted and Lowry defined bases as H + (proton) acceptors Acids are H + (proton) donors NH 3(aq) + H 2 O (l) D NH 4 + (aq) + OH - (aq) H + Arrhenius acid/base: H + (proton) transfer
8 Bronsted-Lowry acid-base theory In Bronsted-Lowry theory, H + ions do not exist in the free state in aqueous solutions, but instead, as H 3 O + ions In this reaction, the acid (HCl) has donated a proton to H 2 O. Water is acting as a B.L. base, since it accepts the proton hydrochloric acid chloride ion hydronium
9 Bronsted-Lowry acid-base theory acid base When water takes a proton (H + ) from hydrochloric acid, two new things are formed in solution (Cl - and H 3 O + ) The products are related (by their formulas) to a reactant each differing by one H + ion from one of the reactants HCl (aq) + H 2 O (l) Cl - (aq) + H 3 O + (aq)
10 Bronsted-Lowry acid-base theory acid base When water takes a proton (H + ) from hydrochloric acid, two new things are formed in solution (Cl - and H 3 O + ) The products are related (by their formulas) to a reactant each differing by one H + ion from one of the reactants HCl (aq) + H 2 O (l) Cl - (aq) + H 3 O + (aq)
11 Bronsted-Lowry acid-base theory acid base When water takes a proton (H + ) from hydrochloric acid, two new things are formed in solution (Cl - and H 3 O + ) The products are related (by their formulas) to a reactant each differing by one H + ion from one of the reactants HCl (aq) + H 2 O (l) Cl - (aq) + H 3 O + (aq)
12 Bronsted-Lowry acid-base theory acid base Two species that differ from each other by one H + are called conjugate pairs The partner that has the extra H + is called the acid and the other is called the base conjugate pair HCl (aq) + H 2 O (l) Cl - (aq) + H 3 O + (aq) chloride ion is the conjugate base of HCl H 3 O + is the conjugate acid of water conjugate pair Conjugate acid/base pairs always differ by one proton in their formulas
13 Bronsted-Lowry acid-base theory Some practice problems: What are the conjugate acids of these? NO 3 - What are the conjugate bases of these? HF OH - H 2 SO 4 C 2 H 3 O - 2 H 2 O NH 3 H 3 PO 4
14 Bronsted-Lowry acid-base theory Amphiprotic substances Some substances can either gain or lose protons, depending on their environment. When water encounters something that is a better proton donor than itself, it acts as a B.L. base H 2 O(l) + H 2 SO 4 (aq) H 3 O + (aq) + HSO 4 -(aq) When water encounters something that is a better base than itself, it acts as a B.L. acid H 2 O(l) + NH 3 (aq) D OH - (aq) + NH 4+ (aq) Water can act as with an acid or a base it is amphiprotic
15 Bronsted-Lowry acid-base theory Many acids are capable of donating more than one proton during acid-base reactions: Carbonic acid: H 2 CO 3 (aq) + H 2 O(l) D HCO 3- (aq) + H 3 O + (aq) HCO 3- (aq) + H 2 O(l) D CO 2-3 (aq) + H 3 O + (aq) Phosphoric acid Mono-, di-, and triprotic acids H 3 PO 4 (aq) + H 2 O(l) D H 2 PO 4- (aq) + H 3 O + (aq) H 2 PO 4- (aq) + H 2 O(l) D HPO 4 2- (aq) + H 3 O + (aq) HPO 4 2- (aq) + H 2 O(l) D PO 4 3- (aq) + H 3 O + (aq) H 2 CO 3 is diprotic H 3 PO 4 is triprotic Just because a molecule has hydrogen in its formula does not mean that compound is an acid. Need to look at the molecule s Lewis structure to see if any H-atoms are acidic.
16 Strengths of acids and bases Some acids (e.g. HCl) ionize almost completely when they are dissolved into water. These acids transfer essentially 100% of their protons to water: HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl - (aq) This equilibrium lies far to the right HCl is a strong acid For many acids, only a small portion of the acid transfers protons to water. For example, in vinegar, acetic acid (HC 2 H 3 O 2 ) is 95% nonionized: HC 2 H 3 O 2(aq) + H 2 O (l) D H 3 O + (aq) + C 2 H 3 O 2 - (aq) Hydrochloric acid in water looks like this Acetic acid in water mostly looks like this This equilibrium lies far to the left HC 2 H 3 O 2 is a weak acid
17 Strengths of acids and bases (e.g. HCl) (e.g. HC 2 H 3 O 2 )
18 Strengths of acids and bases memorize these There are only seven strong acids: Hydrochloric (HCl) Hydrobromic (HBr) Hydroiodic (HI) Nitric (HNO 3 ) Sulfuric (H 2 SO 4 ) Chloric (HClO 3 ) Perchloric (HClO 4 )
19 Strengths of acids and bases Some bases dissociate almost completely. For example, when NaOH is dissolved in water, essentially all of the NaOH is transformed into Na + (aq) + OH - (aq) Others, like ammonia, react only partially: NH 3(aq) + H 2 O (l) D OH - (aq) + NH 4 + (aq) NH 3 is a weak base This equilibrium lies far to the left
20 Strengths of acids and bases The strong bases are the soluble salts of hydroxide ion: LiOH NaOH memorize KOH RbOH CsOH and Ca(OH) 2 Sr(OH) 2 Ba(OH) 2 All group I hydroxides Certain group II hydroxides Need to memorize these
21 Strengths of acids and bases An acid s strength can be reported in terms of an equilibrium constant. The acid ionization constant, K a, is calculated as follows: HA (aq) + H 2 O (l) D H 3 O + (aq) + A - (aq) K a [ H 3 O ][ A [ HA] ] - The size of K a depends on the ratio of [products]/[reactants]. The more an acid ionizes, the higher will be [products] and the lower will be [reactants] - Acids that only weakly ionize will have small K a values - Strong acids will have very large K a values
22 Acid ionization constants Acid strength decreasing All of the acids shown in this table are considered to be weak acids
23 Strengths of acids and bases It s possible to determine the value of an acid ionization equilibrium constant if you know the amounts of products and reactants at equilibrium: Data: for a weak acid (HA), the equilibrium concentrations of products and reactants are: [HA] = M [A - ] = M [H 3 O + ] = M HA (aq) + H 2 O (l) D H 3 O + (aq) + A - (aq) K a [ H 3 O ][ A [ HA] ]
24 Strengths of acids and bases K K K a a a [ H3O ][ A [ HA] [0.0015][0.0015] [0.0085] 2.6x10 4 ]
25 Strengths of acids and bases You can also determine an acid ionization constant if you know the extent to which an acid of a given concentration ionizes A M solution of an acid is 6.0% ionized. Calculate the acid ionization constant. HA (aq) + H 2 O (l) D H 3 O + (aq) + A - (aq)
26 Strengths of acids and bases To solve this problem: Determine what amount of A - is formed when 0.100M HA ionizes by 6.0% (this is the amount of A - that is formed when the reaction reaches equilibrium) Determine the amount of HA that is left over after equilibrium is established The amount of H 3 O + that is formed when the reaction reaches equilibrium will be the same amount as A - Knowing these three quantities, solve for K a HA (aq) + H 2 O (l) D H 3 O + (aq) + A - (aq) K a [ H 3 O ][ A [ HA] ]
27 Strengths of acids and bases Base ionization constants (K b ) can be determined similarly: B (aq) + H 2 O (l) D BH + (aq) + OH - (aq) Example: NH 3(aq) + H 2 O (l) D NH 4 + (aq) + OH - (aq) K b [ BH ][ OH [ B] ]
28 Salts Salts can often influence the acidity/basicity of a solution. Salt is a term that means an ionic compound that consists of a metal or polyatomic positive ion and a nonmetal or polyatomic ion as the negative ion. e.g. NaCl, KC 2 H 3 O 2, MgCO 3, NH 4 NO 3 Salts are not always water-soluble, but the amount that does dissolve will always dissociate (will always break apart and form ions) We ll look at these in examples later in this chapter
29 Acid-base neutralization reactions When an acid and a hydroxide base react, the products are a salt and water: acid hydroxide base HCl (aq) + KOH (aq) KCl (aq) + H 2 O (l) When an acid is completely reacted by a base, a neutralization reaction occurs. In many neutralization reactions the resulting solution is not neutral (i.e. some will result in acidic solutions and some in basic solutions)
30 Acid-base neutralization reactions When a diprotic acid is involved, two equivalent amounts of NaOH are needed for the neutralization: H 2 SO 4(aq) + 2NaOH (aq) Na 2 SO 4(aq) + 2H 2 O (l) Triprotic acid: H 3 PO 4(aq) + 3NaOH (aq) Na 3 PO 4(aq) + 3H 2 O (l) Basically, the hydroxide formed by the base is what reacts with the H 3 O + /H + formed by the acid: For H 2 SO 4 reacting with NaOH: 2H + (aq) + 2OH - (aq) 2H 2 O (l) the real reaction for H 2 SO 4 + 2NaOH
31 Neutralization Reactions An example of acidbase neutralization in the body: antacids in water: Mg(OH) 2(s) D Mg 2+ (aq) + 2OH - (aq) Mg(OH) 2(s) + 2HCl (aq) MgCl 2(aq) + 2H 2 O (l) Mg(OH) 2 is almost insoluble in water (i.e. in the body), but in the presence of acid, it reacts in an acid-base neutralization reaction.
32 Self-ionization of water As we have seen, water is amphiprotic. Even in the presence of other water molecules, water can accept or donate protons. this one acts as a base H 2 O(l) + H 2 O(l) this one acts as an acid H 3 O + (aq) + OH (aq) This is referred to as self-ionization. autoionization The concentration of H 3 O + and OH - ions in water is very small; at 25 o C, in pure water, [H 3 O + ] = [OH - ] = 1.00 x 10-7 M
33 Self-ionization of water this one acts as a base H 2 O(l) + H 2 O(l) this one acts as an acid H 3 O + (aq) + OH (aq) The H 2 O on the left uses an e - pair to form a new bond to H + ion from the H 2 O on the right.
34 Ion-product constant for water The self-ionization reaction of water occurs to a very small extent (equilibrium lies far to the left). We can calculate a value for the equilibrium constant (K w ) using the following: K w [ H3 O ][ OH ] Ion-product constant for H 2 O
35 Ion-product constant for water At 25 o C, [H 3 O + ] = [OH - ] = 1.00 x 10-7 M, thus the value of K w at 25 o C is: K w [ H O 3 ][ OH ] K w [1.00x10 7 ][1.00x10 7 ] K w 1.00x10 14
36 Ion-product constant for water At 25 o C, the product of the concentrations of H 3 O + and OH - must be 1.00 x This is true even if some solute is added which changes the amount of H 3 O + or OH -. K w [ H3 O acid species ][ OH ] base species This means that as [H 3 O + ] increases, [OH - ] decreases (as a solution becomes more acidic, it becomes less basic) Because K w is a constant, as [H 3 O + ] increases, [OH - ] decreases
37 Ion-product constant for water Example: An acidic solute is added to water in an amount that increases [H 3 O + ] to 5.7 x 10-6 M. What is [OH - ] in this solution? K w K w [ H O 3 [ H 1.00x10 5.7x x10 3 [ OH ] 14 6 O ][ OH [ OH [ OH ] ] ] ] Notice: we ve made [H 3 O + ] greater than what it would be under neutral conditions (1.0 x 10-7 M). This makes [OH - ] less than it would be under neutral conditions 1.0 x 10-7 M)
38 Ion-product constant for water
39 Acidic, basic, and neutral solutions The relative amounts of H 3 O + and OH - in a solution determine whether the solution is acidic, basic, or neutral. An acidic solution has a higher concentration of H 3 O + than OH -. A basic solution has a higher concentration of OH - than H 3 O +. K w [ H3 O ][ OH ] Neutral solutions have equal concentrations of H 3 O + and OH -
40 ph Because H 3 O + and OH - concentrations occur over such a large range (typically between 10-1 to M in water), it is more convenient to report [H 3 O + ] as a logarithmic value. ph = -log[h 3 O + ] To calculate the ph of a solution for a known [H 3 O + ], take the logarithm of [H 3 O + ] and multiply the answer by -1
41 ph For cases where the concentration of H 3 O + expressed in scientific notation has a coefficient of 1.0, the ph is just the negative value of the exponent (integral ph value) Example: a solution has [H 3 O + ] = 1.0 x 10-6 M. What is the ph of this solution? [H 3 O + ] = 1.0 x 10-6 ph = 6.00 # of sig figs in concentration = # of decimal places in ph figure
42 ph Calculate the ph of a solution whose [OH - ] is 1.0 x 10-4 M. Use K w to get [H 3 O + ], then get ph: or, use ph + poh = pk w = (at 25 o C)
43 ph When [H 3 O + ] (or [OH - ] values) don t have coefficients of 1.0, the ph values calculated are non-integral. Calculate the ph of a solution that has [H 3 O + ] = 7.23 x 10-8 M ph ph ph log[ H O log( 7.23x ] 8 )
44 ph values and [H 3 O + ] If a solution s ph is known, [H 3 O + ] can be calculated by taking the antilog of the ph (antilog x = 10 x ) ph 10 [ H3O Example, a solution has a ph of What is [H 3 O + ] in this solution? x10 [ H O 4 3 ] [ H O 3 ] ]
45 Interpreting ph values Aqueous solutions that are acidic have [H 3 O + ] > 10-7 M. These solutions have a ph lower than 7 A neutral solution has [H 3 O + ] = 10-7 M, so it has a ph of 7 Aqueous solutions that are basic have [H 3 O + ] < 10-7 M. These solutions have a ph higher than 7 A change of one ph unit corresponds to a change in [H 3 O + ] by a factor of ten
46 Interpreting ph values
47 Interpreting ph values
48 pk a and acid strength We know that an acid s strength can be reported by means of the acid ionization constant, K a. The stronger the acid, the greater the value of K a. Can also report K a like we do for ph (as pk a ), since K a values are often very small. pk log a K a The weaker the acid, the greater will be the value of pk a.
49 Acid ionization constants pk a Acid strength decreasing
50 The ph of aqueous salt solutions Sometimes (most times), the salt of an acid-base neutralization reaction can influence the acid/base properties of water. NaCl dissolved in water: ph = 7 NaC 2 H 3 O 2 dissolved in water: ph > 7 (basic) NH 4 Cl dissolved in water: ph < 7 (acidic) To determine whether a salt will make water acid, basic, or not influence the ph at all, we need to look at the type of reactions that make them.
51 The ph of aqueous salt solutions When an acid-base neutralization reaction occurs, a salt and water are produced: HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O (l) NaCl in water = Na + (aq) + Cl - (aq) The reaction above shows what happens when a strong acid and strong base react. The salt of a strong acid-strong base neutralization reaction has no acid/base properties (the resulting solution would have a ph of 7) When a strong acid/strong base reacts with any other base/acid, reaction goes all the way to the right
52 The ph of aqueous salt solutions When a weak acid is reacted with a strong base, a salt and water are produced: HC 2 H 3 O 2(aq) + NaOH (aq) NaC 2 H 3 O 2(aq) + H 2 O (l) NaC 2 H 3 O 2 in water = Na + (aq) + C 2 H 3 O 2 - (aq) C 2 H 3 O 2 - is the conjugate base of a weak acid (means that C 2 H 3 O 2 - is a weak base) The resulting solution would be basic (ph > 7), even though a neutralization reaction has occurred. If you made up a solution by dissolving NaC 2 H 3 O 2 in water, the solution would be basic.
53 The ph of aqueous salt solutions When a strong acid and a weak base are reacted in a neutralization reaction, the resulting solution is acidic (ph < 7): HCl (aq) + NH 3(aq) NH 4 Cl (aq) NH 4 Cl in water = NH 4 + (aq) + Cl - (aq) NH 4 + is the conjugate acid of a weak base (means NH 4 + is a weak acid) If you made up a solution by dissolving NH 4 Cl in water, the solution would be acidic.
54 The ph of aqueous salt solutions Let s look at why this is so 1. Salt of a strong acid and a strong base: Both the strong acid and strong base would ionize/dissociate completely if put in water HCl (aq) H + (aq) + Cl - (aq) NaOH (aq) Na + (aq) + OH - (aq) The one-way arrows here imply that the reverse reactions do not occur to any significant extent (Cl - (aq) is a really bad base and Na + (aq) is a really bad acid) The conjugate base of a strong acid has no base properties in water The conjugate acid of a strong base has no acid properties in water
55 The ph of aqueous salt solutions Salt of a strong acid and a weak base: HCl (aq) + NH 3(aq) NH 4 Cl (aq) Resulting solution is acidic When NH 3 (a weak base) is dissolved in water, an equilibrium results: Conjugate acid of a weak base has some acid properties in water NH 3(aq) + H 2 O (l) D NH 4 + (aq) + OH - (aq) If NH 4 + is a weak acid, then when a salt containing NH 4 + (e.g. NH 4 Cl) is dissolved in water, the resulting solution will be acidic Conjugate acid of a weak base has acid properties in water
56 The ph of aqueous salt solutions Salt of a weak acid and a strong base: HC 2 H 3 O 2aq) + NaOH (aq) NaC 2 H 3 O 2(aq) + H 2 O (l) When HC 2 H 3 O 2 (a weak acid) is dissolved in water, an equilibrium results: Resulting solution is basic HC 2 H 3 O 2(aq) + H 2 O (l) D H 3 O + aq) + C 2 H 3 O 2 - (aq) Conjugate base of a weak acid has some base properties in water If C 2 H 3 O 2 - is a weak base, then when a salt containing C 2 H 3 O 2 - (e.g. NaC 2 H 3 O 2 ) is dissolved in water, the resulting solution will be basic Conjugate base of a weak acid has base properties in water
57 Chemical equations for salt hydrolysis reactions You can recognize salts that will influence the ph of water from the positive and negative ions in the formula for the salt: + NaCl (Na +, Cl - ) NaC 2 H 3 O 2 (Na +, C 2 H 3 O 2- ) - KF (K +, F - ) NH 4 Cl (NH 4+, Cl - )
58 Chemical equations for salt hydrolysis reactions The positive ion might be acidic and the negative ion might be basic. + - NaCl (Na +, Cl - ) NaC 2 H 3 O 2 (Na +, C 2 H 3 O 2- ) KF (K +, F - ) NH 4 Cl (NH 4+, Cl - )
59 Chemical equations for salt hydrolysis reactions If the cation (positive ion) of the salt is NH 4+, dissolving the salt into water will produce an acidic solution If the anion (negative ion) is the conjugate base of a weak acid, the salt will make the solution basic Cases involving NH 4 + with weak base anions won t be considered conjugate The strong acids bases Hydrochloric (HCl) Cl - Hydrobromic (HBr) Br - Hydroiodic (HI) I - Nitric (HNO 3 ) NO - 3 Sulfuric (H 2 SO 4 ) SO 2-* 4 Chloric (HClO 3 ) ClO - 3 Perchloric (HClO 4 ) ClO - 4 * Conjugate base of H 2 SO 4 is HSO 4-, but SO 4 2- is not basic; HSO 4 - is basic these anions are not basic
60 Chemical equations for salt hydrolysis reactions So, for example: Will NH 4 NO 3 make a solution acidic, basic, or have no effect? NH 4 + will make the solution acidic. NO 3 - is the conjugate base of a strong acid (HNO 3 ), so it is not basic. The resulting solution will be acidic, according to the following chemical equation: NH H 2 O D NH 3 + H 3 O + NH 4 + is acidic a H + donor
61 Chemical equations for salt Another example: hydrolysis reactions Will LiF make a solution acidic, basic, or neutral? The cation isn t NH 4+, so it s not acidic. The anion is F -. The conjugate acid is HF (not one of the strong acids, so F- is a weak base) F - + H 2 O D HF + OH - F - is basic a H + acceptor
62 Buffers Buffers are mixtures of weak acid/conjugate base pairs that are able to resist significant changes in ph when small quantities of acids or bases are added. They are particularly resistant to changes in ph, when small amounts of a strong acid or base is added. Example When 0.02 mol of NaOH is added to 1L of water, the ph jumps from 7.0 to 12.3 (5.3 units) When 0.02 mol of NaOH is added to 1L of 0.3 M HC 2 H 3 O 2 /0.3 M NaC 2 H 3 O 2 buffer, the ph jumps just 0.06 units Buffers resist changes in ph
63 Buffers in everyday life Because so many chemical reactions (including ones that occur in our body) produce/consume H +, ph regulation is essential Our blood is buffered (H 2 CO 3 /HCO 3- /CO 3 2- ) to a ph of 7.4. Many metabolic reactions produce H + and CO 2. ph is extremely important in cellular reaction (e.g. many enzymes will work only near ph = 7.4) The body needs to regulate ph within a narrow range (keep it very close to 7.4). Below ph = 6.8 and above ph = 7.8, cell death occurs
64 Buffers How do buffers work? Since buffers contain both acid and base components, they are able to offset small quantities of another acid or base added to them. The addition of an acid to a buffer consumes some of the base that is already present in the buffer Acid + Buffer weak acid + conj. base neutralization products
65 Buffers How do buffers work? Since buffers contain both acid and base components, they are able to offset small quantities of another acid or base added to them. The addition of a base consumes some of the acid that is already present in the buffer Base + Buffer weak acid + conj. base neutralization products
66 Buffers As an example, consider a buffer that is made up from the following weak acid/conjugate base pair: Acid = HF Conj. base = F - (in the form of NaF) Buffer mixture
67 Buffers Addition of a acid to a buffer If HCl is added to this mixture, it will react with the base component of the buffer: H + + F - HF Remember, Cl - (from HCl) has no influence on the ph of solutions, so it is not shown in this reaction. Reaction of an acid with the buffer consumes a bit of the buffer s base and makes more of the buffer s acid.
68 Buffers Addition of a base to a buffer If NaOH is added to this mixture, it will react with the acid component of the buffer: OH - + HF F - + H 2 O Na + (in NaOH) has no influence on ph, so it is not included in this reaction. The reaction of a base with the buffer consumes a bit of the buffer s acid and makes more of the buffer s base.
69 The Henderson-Hasselbalch equation The Henderson-Hasselbalch equation provides a means of calculating the ph of a buffer, provided the amounts of weak acid and conjugate base are known (or, more importantly, the ratio of their concentrations) concentration of weak base in buffer ph pk a [ A ] log [ HA ] K a is the acid ionization constant for the weak acid/base pair concentration of weak acid in buffer
70 The Henderson-Hasselbalch equation For example, a buffer is made up by adding 2.0 mol of HC 2 H 3 O 2 and 1.0 mol of NaC 2 H 3 O 2 to enough water to make up 1L of solution. If K a for HC 2 H 3 O 2 is 1.8 x 10-5, what is the ph of the resulting solution? ph pk a [ A ] log [ HA ]
71 The Henderson-Hasselbalch equation ph ph ph ph 4.44 [ A ] log [ HA] log 1.8x pk a 5 log [1.0] [2.0]
72 The Henderson-Hasselbalch equation It can be seen that if the amounts of weak acid and conjugate base in the buffer are equal, the ph will be pk a ph ph ph ph pk pk pk pk a a a a [ A ] log [ HA] log 1 0
73 The Henderson-Hasselbalch equation Some hints on the use of logarithms and the H.H. equation: log of a ratio less than 1 is a negative # log of a ratio greater than 1 is a positive # log of 1 = 0
74 The Henderson-Hasselbalch equation If a buffer contained 1000 times as much base as conjugate acid, the ph of the buffer would be pk a + 3 (for the acetic acid/acetate buffer we just looked at, ph would be 7.74 ( ) ph pk a [ A ] log [ HA ] ratio log(ratio)
75 Electrolytes An electrolyte is a substance whose solution conducts electricity. Electrolytes produce ions (by dissociation of an ionic compound or ionization of an acid) in water. Salts are a typical example of an electrolyte. Non-electrolytes do not ionize when put into water. Glucose and isopropyl alcohol are examples of non-electrolytes. Example of an electrolyte: Example of a non-electrolyte: H 2 O NaCl(s) Na + (aq) + Cl - (aq) H 2 O C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) (ions present in solution) (no ions)
76 Electrolytes Some electrolytes are able to (essentially) completely ionize/dissociate in water. Strong acids Strong bases Soluble salts Some electrolytes produce equilibrium mixtures of ionized and non-ionized forms Weak acids Weak bases called strong electrolytes Example: HCl + H 2 O H 3 O + + Cl - called weak electrolytes Example: HC 2 H 3 O 2 + H 2 O D H 3 O + + C 2 H 3 O 2 -
78 Equivalents and milliequivalents One equivalent (1 Eq) is the molar amount of an ion that is needed to supply one mole of positive (or negative) charge. One mole of NaCl supplies one mole of + charge ions (Na + ) one mole of charge ions (Cl - ) 1 mole of Cl - = 1 Eq 1 mole of Ca 2+ = 2 Eq 1 mole of PO 4 3- = 3 Eq Each of these is considered to be one equivalent (1 Eq) Equivalents are units used like moles. They express the amount of ions (charge). Just like M (mol/l), concentrations can be expressed with equivalents, as Eq/L
79 Equivalents and milliequivalents Because the concentrations of ions in body fluids is usually low, the term, milliequivalents, is often seen. 1 meq = Eq 1000 meq = 1 Eq
80 Equivalents and milliequivalents Example problem: the concentration of Na + in blood is 141 meq/l. How many moles of Na + are present in 1 L of blood? 1 mol Na + = 1 Eq = 1000 meq 141mEq 1mol _ Na L 1000mEq 1 L 0.141mol _ Na volume of blood concentration of Na + Eq mol
81 Equivalents and milliequivalents Another example: The concentration of Ca 2+ ion present in a blood sample is found to be 4.3 meq/l. How many mg of Ca 2+ are present in 500 ml of blood? 1 mol Ca 2+ = 2 Eq = 2000 meq 1L 4.3mEq 1mol _ Ca 1000mL L 2000mEq 40.08g _ Ca 1mol _ Ca 1000mg 1g mL 43mg _ Ca volume of blood ml L conc. of Ca 2+ Eq mol mol g g mg
82 Acid-Base Titrations titrant Titrations are experiments in which two solutions are made to react together (a balanced equation for the reaction must be known). analyte Using C 1 V 1 = C 2 V 2, the concentration of the analyte can be determined (also need to consider coefficients) In an acid-base titration, a known volume and concentration of base (or acid) is slowly added to a known volume of acid (or base). An indicator is often used to find the endpoint in acid-base titration experiments.
83 Acid-Base Titrations Before endpoint (acidic) After endpoint (basic) acid has been neutralized
84 Acid-Base Titrations Know the concentration of this solution and can measure the volume needed to reach endpoint (example, this could be M NaOH) Know the volume of this solution, but not the concentration (example, this could be ml of HNO 3 )
85 Acid-Base Titrations Example: It takes ml of M NaOH to neutralize ml of HNO 3. What is the concentration of HNO 3? NaOH + HNO 3 NaNO 3 + H 2 O At endpoint: # of moles of NaOH added = # of moles of HNO 3 C NaOH V NaOH = C HNO3 V HNO3
86 Acid-Base Titrations C 0.100M 21.09mL C 25.00mL C C NaOH HNO3 HNO3 V NaOH C HNO M M V HNO3 HNO3
87 Acid-Base Titrations In a sulfuric acid-sodium hydroxide titration, 17.3 ml of M NaOH is needed to neutralize 25.0 ml of H 2 SO 4 of unknown concentration. Find the molarity of the H 2 SO 4 solution.
Mr. B s Chemistry Acids and Bases Unit 11 Name Block Let s start our discussion of acids and bases by defining some terms that are essential to the topics that follow. Arrhenius acids and bases are: acid
Unit Nine Notes N C U9 I. AcidBase Theories A. Arrhenius Acids and Bases 1. Acids contain hydronium ions (H O ) commonly referred to as hydrogen ions (H ) that dissociate in water a. Different acids release
Acids And Bases A. Characteristics of Acids and Bases 1. Acids and bases are both ionic compounds that are dissolved in water. Since acids and bases both form ionic solutions, their solutions conduct electricity
ACIDS AND BASES A. CHARACTERISTICS OF ACIDS AND BASES 1. Acids and bases are both ionic compounds that are dissolved in water. Since acids and bases both form ionic solutions, their solutions conduct electricity
Chapter 10 Acids and Bases 1 Properties of Aqueous Solutions of Acids and Bases Aqueous acidic solutions have the following properties: 1. They have a sour taste.. They change the colors of many indicators.
Aqueous Reactions and Solution Stoichiometry (continuation) 1. Electrolytes and non-electrolytes 2. Determining Moles of Ions in Aqueous Solutions of Ionic Compounds 3. Acids and Bases 4. Acid Strength
Periodic Table Name Academic Chemistry Acids & Bases Notes Unit #14 Test Date: 20 cincochem.pbworks.com Acid Base cincochem.pbworks.com Notes Find ph To go from [H 3 O + ] to ph EXAMPLE: [H 3 O + ] = 3.23
Acids and Bases Properties, Reactions, ph, and Titration C-19 2017 Properties of acids 1. Taste Sour (don t try this except with foods). 2. Are electrolytes (conduct electricity). Some are strong, some
Acids and Bases Acids and bases, as we use them in the lab, are usually aqueous solutions. Ex: when we talk about hydrochloric acid, it is actually hydrogen chloride gas dissolved in water HCl (aq) Concentrated
Unit 9 Acids, Bases, & Salts Acid/Base Equilibrium Properties of Acids sour or tart taste strong acids burn; weak acids feel similar to H 2 O acid solutions are electrolytes acids react with most metals
Chapter 16 1 Learning Objectives Acid Base Concepts Arrhenius Concept of Acids and Base a. Define acid and base according to the Arrhenius concept. Brønsted Lowry Concept of Acids and Bases a. Define acid
Chapter 7 Acids and Bases 7.1 The Nature of Acids and Bases 7.2 Acid Strength 7.3 The ph Scale 7.4 Calculating the ph of Strong Acid Solutions 7.5 Calculating the ph of Weak Acid Solutions 7.6 Bases 7.7
Section 32 Acids and Bases 1 Copyright (c) 2011 by Michael A. Janusa, PhD. All rights reserved. Acid-Base Concepts Acids and bases are among the most familiar and important of all chemical compounds. You
Chapter 14 Acid- Base Equilibria Study Guide This chapter will illustrate the chemistry of acid- base reactions and equilibria, and provide you with tools for quantifying the concentrations of acids and
Physical Properties Acid and Bases Chemistry 30 Acids Corrosive when concentrated Have a sour taste Bases Corrosive when concentrated Have a bitter taste Often have a sharp odour Chemical Properties Indicators
Chemistry I Notes Unit 10: Acids and Bases Acids 1. Sour taste. 2. Acids change the color of acid- base indicators (turn blue litmus red). 3. Some acids react with active metals and release hydrogen gas,
Acids and Bases Chapter 11 Acids and Bases in our Lives Acids and bases are important substance in health, industry, and the environment. One of the most common characteristics of acids is their sour taste.
As you work through the chapter, you should be able to: Advanced Placement Chemistry Chapters 14 16 Syllabus Chapter 14 Acids and Bases 1. Describe acid and bases using the Bronsted-Lowry, Arrhenius, and
Acids and Bases Acid A compound that produces H + ions when dissolved in water. Examples! Vinegar Acetic acid Lemon Juice Citric acid Sour Candy Malic acid (and others) Milk Lactic acid HCl(aq) Acid Properties
Unit 4a Acids, Bases, and Salts Theory Chemistry 12 Arrhenius Theory of Acids and Bases The first theory that was proposed to explain the actions of acids and bases was by Svante Arrhenius. It is still
You are already familiar with some acid and base chemistry. According to the Arrhenius model, acids are substances that when dissolved in water ionize to yield hydrogen ion (H + ) and a negative ion. e.g.
ACID BASE EQUILIBRIUM Part one: Acid/Base Theories Learning Goals: to identify acids and bases and their conjugates according to Arrhenius and Bronstead Lowry Theories. to be able to identify amphoteric
Chem 105 Tuesday March 8, 2011 Chapter 17. Acids and Bases 1) Define Brønsted Acid and Brønsted Base 2) Proton (H + ) transfer reactions: conjugate acid-base pairs 3) Water and other amphiprotic substances
Chapter 15 Aqueous Equilibria: Acids and Bases Properties of Acids and Bases Generally, an acid is a compound that releases hydrogen ions, H +, into water. Blue litmus is used to test for acids. Blue litmus
There are three definitions for acids and bases we will need to understand. Arrhenius Concept: an acid supplies H + to an aqueous solution. A base supplies OH to an aqueous solution. This is the oldest
Acids, Bases, & Neutralization Name Warm-Ups (Show your work for credit) Date 1. Date 2. Date 3. Date 4. Date 5. Date 6. Date 7. Date 8. Acids, Bases, & Neutralization 2 Study Guide: Things You Must Know
10.1 Acids and Bases in Aqueous Solution Arrhenius Definition of Acids and Bases An acid is a substance that gives hydrogen ions, H +, when dissolved in water. In fact, H + reacts with water and produces
NAME: UNIT #11: Acids and Bases ph and poh Neutralization Reactions Oxidation and Reduction 1. SELF-IONIZATION OF WATER a) Water molecules collide, causing a very small number to ionize in a reversible
Acids and Bases Unit 10 1 Properties of Acids and Bases Acids Bases Taste Sour Turns Litmus Dye Red Reacts with Metals to give H 2 (g) Taste Bitter Turns Litmus Dye Blue Do Not React with Metals Reacts
10 Reactions in Aqueous Solutions I: Acids, Bases & Salts CHAPTER GOALS 1. Properties of Aqueous Solutions of Acids and Bases 2. The Arrhenius Theory 3. The Hydronium Ion (Hydrated Hydrogen Ion) 4. The
Chapter 14 Acids and Bases General Properties of Acids 1. An acid tastes sour - acidus = Latin, sour; acetum= Latin, vinegar 2. An acid turns indicator dye litmus from blue to red. 3. An acid reacts with
Acids Definition of Acid Acids are substances that contain H + ions that ionize when dissolved in water. Arrhenius acid: a compound that increases the concentration of H + ions that are present when added
STUDENT VERSION Unit 9: Acids, Bases, & Salts Unit Vocabulary: Arrhenius acid Arrhenius base Bronsted-Lowry acid Bronsted-Lowry base Electrolyte hydronium ion hydroxide ion indicator (acid/base) neutralization
Chapter 10 - Acids & Bases 10.1-Acids & Bases: Definitions Arrhenius Definitions Acids: substances that produce hydrogen ions when dissolved in H 2 O Common Strong Acids: Common Weak acids: Organic carboxylic
Name: Per: Date: Unit 11 - Acids, Bases and Salts Chemistry Accelerated Chemistry I Define each of the following: 1. Acidic hydrogens 2. Binary acids 3. Oxyacids 4. Carboxylic acid 5. Amines Name the following
ACID-BASE TITRATION AND PH Section 1 Aqueous Solutions and the Concept of ph Hydronium and Hydroxide Ions Acids and bases form hydroxide and hydronium ions These ions are not the only ones in an aqueous
Acids and Bases Properties of Acids and Bases Acids taste. Lemon juice and, for example, are both aqueous solutions of acids. Acids conduct electricity; they are. Some are strong electrolytes, while others
Acids and Bases Acid/Base Definitions Arrhenius Model Acids produce hydrogen ions in aqueous solutions Bases produce hydroxide ions in aqueous solutions Bronsted-Lowry Model Acids are proton donors Bases
Acids and Bases Chapter 11 Acids and Bases in our Lives We produce lactic acid in our muscles when we exercise. Acid from bacteria turns milks sour in the products of yogurt and cottage cheese. We have
Worksheet 4.1 Conjugate AcidBase Pairs 1. List five properties of acids that are in your textbook. Acids conduct electricity, taste sour, neutralize bases, change the color of indicators, and react with
Chapter 17 Acids and Bases - we are all familiar with 'acids' - depicted on television as burning liquids - from foods (i.e. vinegar) - taste "sour" or "tart' - less familiar with 'bases' - taste "bitter"
(Hebden Unit 4 page 109 182) 182) We will cover the following topics: 1. Definition of Acids and Bases 2. Bronsted-Lowry Acids and Bases 2 1 Arrhenius Definition of Acids and Bases An acid is a substance
CHAPTER 15: ACIDS AND BASES Part One: Acid-Base Concepts A. Properties of Aqueous Solutions of Acids. 1. Sour taste. (Examples: vinegar = acetic acid; lemons - citric acid) 2. Change the colors of many
Unit 4: Acid/Base I I) Introduction to Acids and Bases What is an acid? http://www.kidsknowit.com/flash/animations/acidsbases.swf What are properties of acids? 1) Acids react with. 2) Acids create when
INTRODUCTION TO ACIDS AND BASES ALIGNED STANDARDS S.C. 912.P.8.11 Relate acidity and basicity to hydronium and hydroxide concentration and ph. S.C.912.N.1.2 Describe and explain what characterizes science
Acid-Base Equilibria You have just completed a chapter on equilibrium. That chapter focused primarily on gas phase reactions (with a few exceptions). This section on Acid-Base equilibria (along with the
Section 1 Properties of Acids and Bases Objectives List five general properties of aqueous acids and bases. Name common binary acids and oxyacids, given their chemical formulas. List five acids commonly
19.1 Acids & Bases 1. Compare and contrast the properties of acids & bases. 2. Describe the self-ionization of water & the concept of K w. 3. Differentiate between the Arhennius & Bronsted-Lowry models
ph calculations MUDr. Jan Pláteník, PhD Brønsted-Lowry concept of acids and bases Acid is a proton donor Base is a proton acceptor HCl(aq) + H 2 O(l) H 3 O + (aq) + Cl - (aq) Acid Base Conjugate acid Conjugate
Unit 10: Acids and Bases PROPERTIES OF ACIDS & BASES Properties of an Acid: a Tastes sour substance which dissociates (ionizes, breaks apart in solution) in water to form hydrogen ions Turns blue litmus
Chemistry 12 Acid-Base Equilibrium II Name: Date: Block: 1. Strengths of Acids and Bases 2. K a, K b 3. Ionization of Water 4. Relative Strengths of Brønsted-Lowry Acids and Bases Strengths of Acids and
Properties of Acids and Bases Chapter 14 Acids and Bases Svante Arrhenius (1859-1927) First to develop a theory for acids and bases in aqueous solution Arrhenius Acids Compounds which dissolve (dissociate)
Lecture INTRODUCTORY CHEMISTRY Concepts and Critical Thinking Seventh Edition by Charles H. Corwin Acids and Bases Properties of Acids An acid is any substance that releases hydrogen ions, H +, into water.
Acid-Ionization Equilibria Acid-Base Equilibria Acid ionization (or acid dissociation) is the reaction of an acid with water to produce hydronium ion (hydrogen ion) and the conjugate base anion. (See Animation:
CHEM 120 Online Chapter 10. Date: 1. Which of the following statements concerning Arrhenius acids and Arrhenius bases is incorrect? A) In the pure state, Arrhenius acids are covalent compounds. B) In the
The Chemistry of Acids and Bases 1 Acid and Bases 4 Acid and Bases 2 Acids Have a sour taste. Vinegar is a solution of acetic acid. Citrus fruits contain citric acid. React with certain metals to produce
Chapter 14 Acids and Bases I. Bronsted Lowry Acids and Bases a. According to Brønsted- Lowry, an acid is a proton donor and a base is a proton acceptor. Therefore, in an acid- base reaction, a proton (H
Name Chemistry Pre-AP Notes: Acids and Bases Period I. Describing Acids and Bases A. Properties of Acids taste ph 7 Acids change color of an (e.g. blue litmus paper turns in the presence of an acid) React
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
CHAPTER 14 ACIDS AND BASES Topics Definition of acids and bases Bronsted-Lowry Concept Dissociation constant of weak acids Acid strength Calculating ph for strong and weak acids and bases Polyprotic acids
NOTES Acids, Bases & Salts Arrhenius Theory of Acids & Bases: an acid contains hydrogen and ionizes in solutions to produce H+ ions: a base contains an OH group and ionizes in solutions to produce OH ions:
Acids and Bases Chapter 15 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain
I. A/B General A. Definitions 1. Arrhenius 2. B-L 3. Lewis *H+=H3O+ B. Properties 1. Acid 2. Base II. ACIDS A. Types 1. Binary 2. Ternary 3. Acid Anhydride B. Protic 1. mono 2. di 3. tri C. Strengths 1.
Section 14.1 Defining Acids and Bases Properties of acids and bases Chapter 14 Properties of Acids and Bases taste sour Acids taste bitter Bases conduct electricity no characteristic feel react with metals
3 ACID AND BASE THEORIES: A) Arrhenius Acids produce H+ and bases produce OH not always used because it only IDs X OH as basic species B) Bronsted and Lowry Acid = H + donor > CB = formed after H + dissociates
CHAPTER 19 Acids, Bases, and Salts 19.1 Acid Base Theories ACIDS tart or sour taste Electrolytes Strong acids are corrosive Acid Facts... indicators will change color Blue litmus paper turns pink react
Unit 9: Acids and Bases Chapter 19 I. Introduction In aqueous solutions, the solvent is. Aqueous solutions contain. In the self-ionization of water, the hydrogen ion (H+) exists in solution as the ion.
Chapter 16 Acid-Base Equilibria Arrhenius Definition Acids produce hydrogen ions in aqueous solution. Bases produce hydroxide ions when dissolved in water. Limits to aqueous solutions. Only one kind of
Talk n Acids & Bases... Lady Dog! Definitions So far in this course, we have looked at processes in chemistry that deal with, or are best explained by, ionic salts or molecules. Now we will turn our attention
Arrhenius Definition: Classic Definition of Acids and Bases Acid: A substance that increases the hydrogen ion concetration, [H ], (also thought of as hydronium ion, H O ) when dissolved in water. Acids
Understanding the shapes of acidbase titration curves AP Chemistry Neutralization Reactions go to Completion Every acidbase reaction produces another acid and another base. A neutralization reaction is
Chapter 16 Acid-Base Equilibria Learning goals and key skills: Understand the nature of the hydrated proton, represented as either H + (aq) or H 3 O + (aq) Define and identify Arrhenuis acids and bases.
Principles of Reactivity: The Chemistry of Acids and Bases **a lot of calculations in this chapter will be done on the chalkboard Do not rely on these notes for all the material** Acids, Bases and Arrhenius
Acids Definition of Acid Acids are substances that contain H + ions that ionize when dissolved in water. Arrhenius acid: a compound that increases the concentration of H + ions that are present when added
William L Masterton Cecile N. Hurley http://academic.cengage.com/chemistry/masterton Chapter 13 Acids and Bases Edward J. Neth University of Connecticut Outline 1. Brønsted-Lowry acid-base model 2. The
Chapter 4 (Hill/Petrucci/McCreary/Perry Chemical Reactions in Aqueous Solutions This chapter deals with reactions that occur in aqueous solution these solutions all use water as the solvent. We will look
The Chemistry of Acids and Bases 1 Acid and Bases 2 Acid and Bases 3 Acid and Bases 4 Acids 5 Have a sour taste. Vinegar is a solution of acetic acid. Citrus fruits contain citric acid. React with certain
19.1 ACID-BASE THEORIES Section Review Objectives Define the properties of acids and bases Compare and contrast acids and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis Vocabulary
Acid-Base Equilibria 1 Will the following salts be acidic, basic or neutral in aqueous solution? 1.NH 4 Cl.NaCl.KC H O 4.NaNO A = acidic B = basic C = neutral Solutions of a Weak Acid or Base The simplest