Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Topic 12 – Acids & Bases3.2.1 – Classify substances using the hydronium and hydroxide concentrations.3.2.2 – Summarize the properties of acids and bases.3.2.3 – Infer the quantitative nature of a solution (molarity, dilution, and titration with a 1:1 molar ratio).

I. Theories of Acids & Bases There are two acid-base theories with which you need to be familiar:

1. Arrhenius acids & bases2. Br Ø nsted-Lowry acids & bases

(We will not consider Lewis acids & bases here…)

ARRHENIUS ACIDS & BASES

♦ The Arrhenius theory of acids and bases is the more traditional theory of the three. It was postulated before the other theories, and thus is more specific (less inclusive) than either BrØnsted-Lowry or Lewis acids and bases.

Arrhenius acid – a chemical compound that contains hydrogen and ionizes in H2O to form hydrogen ions (a.k.a. hydronium ions = H3O+)

HCl(aq) + H2O(l) H3O+(aq) + Cl–

(aq)

Arrhenius base – a chemical compound that contains hydroxide ions (OH–) and dissociates in H2O to form hydroxide ions

NaOH(s) Na+(aq) + OH–

(aq)

A solution containing OH- ions is basic, or “alkaline”.

BRNSTED-LOWRY ACIDS & BASES

The BrØnsted-Lowry theory of acids and bases (often called simply the BrØnsted theory) was developed later, and broadens the definition of what qualifies as an acid or base.

BrØnsted-Lowry acid – a molecule or ion that is a proton donor In other words, a BrØnsted acid is any molecule or ion that can donate a H+ ion.

BrØnsted-Lowry base – a molecule or ion that is a proton acceptor In other words, a BrØnsted base is any molecule or ion that can receive a H+ ion.

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In a Bronsted acid-base rxn., protons are transferred from the acid to the base.

Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

You need to be able to identify conjugate acids and bases of , as well: conjugate base – what remains after an acid has given up a proton

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–

(aq)

acid conjugate base

In compounds where there are several H atoms (as with acetic acid above), the H which is held most weakly is generally the one which is lost, which must be reflected in the writing of the compound.

conjugate acid – formed by adding a proton to a base

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–

(aq)

base conjugate acid

II. Properties of Acids & Bases Properties of acids :

sour taste change the color of dyes known as “acid-base indicators”

acids turn litmus paper RED acids are clear in phenolphthalein

react with bases to produce a salt and water (acids neutralize bases, and vice versa):

NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

Acid-base neutralization rxns. are noticeably exothermic.

react with carbonates & bicarbonates to produce CO2(g) & H2O(l): This reaction is driven to products by formation of a gas (which results in an increase in

entropy). Acids react with carbonates to produce carbonic acid, which is unstable and

decomposes into water and carbon dioxide gas:

Na2CO3(aq) + 2HCl(aq) 2NaCl(aq) + H2CO3(aq) H2CO3(aq) H 2O(l) + CO2(g) .

Na2CO3(aq) + 2HCl(aq) 2NaCl(aq) + H2O(l) + CO2(g)

Acids likewise react with bicarbonates (such as NaHCO3) to produce a salt, water, and carbon dioxide:

HC2H3O2(aq) + NaHCO3(aq) H2O(l) + CO2(g) + NaC2H3O2(aq)

CO2(g) produced from the reaction of acids with carbonates will turn limewater (saturated CaO(aq)) a milky white when bubbled through it.

Some acids react with metals to produce H2 gas:

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Reaction of Zn with HCl in a test tube

Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Any metal which is above H2 on the activity series will displace hydrogen in a single replacement reaction. The higher up the metal is on the activity series, the faster it will react to displace hydrogen (you will see more vigorous bubbling):

Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

You can prove it is hydrogen (and not some other gas) by capturing the hydrogen in an upside down test tube and igniting the H2 gas with a match or splint.

electrolytes (acids conduct electricity in aqueous solution)

Properties of Bases :

bitter taste feel slippery to the skin change the color of dyes known as “acid-base indicators”

bases turn litmus paper BLUE bases turn phenolphthalein pink

alkalis are bases which dissolve in water electrolytes (bases conduct electricity in aqueous solution)

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Titration of H2CO3 with phenolphthalein indicator; note that

the acid is clear, before titration

Bitter tasting medicines often contain alkaloid bases such as codeine

Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

III. Strong and Weak Acids & Bases strong acid – acids that ionize 100% in aqueous solution:

HCl, HBr, HI, H2SO4, HNO3, HClO4

For hydrohalic acids, a decrease in bond strength results in a stronger acid (greater ease of ionization):

Consider a 0.5M solution of HCl as an example of a strong acid: When placed in water, the hydrochloric acid will completely ionize / dissociate into

aqueous hydronium and chloride ions:

HCl(aq) + H2O(l) H3O+(aq) + Cl–

(aq)

Because strong acids such as HCl ionize completely, producing many ions in aqueous solution, they are considered strong electrolytes. They conduct electricity very well in aqueous solution.

For the same reason (they produce so many H3O+ ions), strong acids have a low pH.

weak acid – acids that ionize less than 100%, and are weaker electrolytes ex: HF, HC2H3O2 or CH3COOH (acetic or ethanoic acid), H2CO3 (carbonic acid) Consider a 0.1M solution of ethanoic acid:

When placed in water, the acetic acid will only ionize approximately 1% into aqueous hydrogen (hydronium) and acetate (ethanoate) ions:

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO–

(aq)

The forward reaction proceeds only slightly before equilibrium is reached – few products are created because only 1% of the ethanoic acid ever ionizes.

Since so few ions are produced in solution, the conductivity of weak acids is much lower than strong acids (they are weaker electrolytes).

For the same reason, the pH of weak acid solutions is higher (due to the relatively low [H3O+] ) – the pH is still below 7, but probably greater than 2.

Equimolar solutions of strong and weak acids (or bases, for that matter) could thus be differentiated on the basis of either pH or conductivity.

♦ 3 animations of strong vs. weak acids and bases:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/017_AQUEOUSACIDS.MOV http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/016_STRONGWEAK.MOV http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/018_AQUEOUSBASE.MOV

strong base – completely dissociates in solution, yielding a high concentration of OH– ions: all Group I hydroxides (KOH, NaOH, etc.) & barium hydroxide (Ba(OH)2)

Consider a 0.5M solution of NaOH as an example:

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

When placed in water, the sodium hydroxide will completely dissociate into aqueous sodium and hydroxide ions:

NaOH(aq) Na+(aq) + OH–

(aq)

Because strong bases such as NaOH dissociate completely, producing many ions in aqueous solution, they are considered strong electrolytes. They conduct electricity very well in aqueous solution.

For the same reason, the pH of strong base solutions is very high (due to the very low [H3O+] and high [OH–] ).

weak base – bases which dissociate less than 100%, and are weaker electrolytes ex: NH3 (ammonia), Al(OH)3, etc. Consider a 0.1M solution of ammonia (NH3):

When placed in water, the ammonia will only dissociate slightly into aqueous ammonium and hydroxide ions:

Consequently, the forward reaction proceeds only slightly before equilibrium is reached – few products are created because most of the weak base never dissociates.

Since so few ions are produced in solution, the conductivity of weak bases is much lower than strong bases (they are weaker electrolytes).

For the same reason, the pH of weak base solutions is lower (a low [H3O+], but not as low as strong bases) – the pH is still above 7, but probably not as high as 12 – 14, which you might expect for strong bases.

STRENGTH VS. CONCENTRATION

Remember that there is a difference between strength and concentration:

Concentration has to do with dilution (amount of water, relative to the amount of acid); Strength deals with the extent of dissociation.

A concentrated acid contains more acid molecules relative to the amount of water.(by “concentration”, we mean molarity – moles of acid per liter of solution) A 1M solution of HCl has 1 mole of acid dissolved per liter of water, whereas a 6M

solution of HCl has 6 moles of acid per liter of water.

A strong acid ionizes to a greater extent in water than a weak acid.

So, a strong acid that is more dilute (less concentrated), may still have more H+ ions in solution than a concentrated weak acid. (ex: 0.50 M HNO3 vs. 1.50 M HC2H3O2)

Equimolar solutions of strong and weak acids (or bases, for that matter) could thus be differentiated on the basis of either pH or conductivity.

Relative acidities can also be found by neutralizing two acids with a strong base in the presence of an indicator, to determine which is the stronger acid.

H2SO4(aq) + NH3(aq) NH4+

(aq) + HSO4–

(aq) Bronsted Bronsted conjugate conjugate

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

acid base acid base stronger stronger weaker weaker acid base acid base

The stronger an acid, the weaker its conjugate base. The stronger a base, the weaker its conjugate acid. The reaction will go in the direction that produces the weaker acid and base.

HSO4-(aq) + H2O(l) H3O+

(aq) + SO4 2–

(aq)

acid base conjugate conjugate acid base

NH3(aq) + H2O(l) OH–(aq) + NH4

+(aq)

base acid conjugate conjugate base acid

Anything that can react as both an acid and a base is called amphoteric.(in this case, H2O is amphoteric)

IV. The pH Scale & pH Calculations All water undergoes a very small amount of auto-ionization:

H2O(l) + H2O(l) H3O+(aq) + OH–

(aq)

At 25C, the concentrations of hydronium and hydroxide produced are relatively constant: 10–7 moles of each ion are formed per liter of water [H3O+] & [OH–] = 10–7 M

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

The product of [H3O+] x [OH–] is always the same in water and dilute aqueous solutions at 25C:

Kw = [H3O+][OH–] = 10–7 M x 10–7 M = 10–14 M2

so: Kw = 10–14 M2 = [H3O+][OH–]

The value of Kw is 10–14 M2 only at 25C! The value of Kw varies with temperature.

You have prepared a 0.0001 M solution of HNO3 for lab. Calculate the [H3O+] of this solution. Calculate the [OH–] of this solution.

Calculate the [H3O+] & [OH–] of a solution of 2.87 x 10–6 M NaOH.

pH OF AQUEOUS SOLUTIONS (STRONG ACIDS & BASES)

pH ranges from 0-14:

7 – 14 = basic

7 = neutral

0 – 7 = acidic

One pH unit represents a tenfold change in acidity or basicity.

A solution with a pH of 5 is ten times more acidic than a solution with a pH of 6 (it has ten times the [H3O+] of a solution with a pH of 6).

A solution with a pH of 3 is 100 times more acidic than a solution with a pH of 5.

A solution with a pH of 11 is 1000 times more basic than a solution with a pH of 8 (it has 1000 times the [OH–], or 1/1000th the [H3O+], of a solution with a pH of 6).

pH = –log [H3O+]

Find the pH of a 10–3 M HCl solution.

Find the pH of a 10–4 M NaOH solution.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

pOH = –log [OH–]

What is the pOH of a 1 x 10–8 M HCl solution?

pH + pOH = 14

What is the pOH of a solution whose pH is 9.45?

[H3O+] = 10–pH

Determine the [H3O+] of an aqueous solution with a pH of 6.

Determine the [OH–] of an aqueous solution with a pH of 10.49.

[OH–] = 10–pOH

What is the [OH–] of an aqueous solution with a pH of 11.35?

V. Acid-Base Titrations

An acid-base titration is a process used to find the concentration of a solution (either the acid or base, in this case, although there are other types of titrations). In a typical titration, you fill up two burettes, one with the acid, one with the base, and clamp both

to a ring stand. You must already know the concentration of one of these solutions in order to find the concentration of the other.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

You measure out a certain volume of the solution whose concentration you do not know from the burette into an Erlenmeyer flask. Also into this flask you usually add a little distilled water, and a few drops of an indicator. Indicators are weak acid or base dyes, whose colors are sensitive to pH. In other words,

indicators will change color over a certain pH range, as indicated in the figure on the next page. Examples of indicators include phenolphthalein and litmus.

So now in your Erlenmeyer flask, you have some volume of your solution of unknown concentration, distilled water and a few drops of your chosen indicator. You place your flask beneath the other burette, filled with the solution whose concentration you do know. This solution of known concentration is called the standard solution or titrant. You begin adding a few drops of titrant at a time from the burette into your flask until you get a color change in the solution from your indicator. When you get a color change that doesn’t go away after a few seconds (e.g. – a pale, persistent pink color in phenolphthalein), then you are done with your titration. You just have to record the volume of titrant added, and now you’re ready to calculate the molarity of your other solution.

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Here are the apparatus used when performing a titration

Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Indicators change color gradually over a pH range. For example, bromphenol blue is yellow at pH values below 4, then it gradually transitions from yellow to blue (and is essentially

green at the endpoint, at a pH slightly less than 5). Bromthymol blue is blue at pHs above about 5.5.

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IB ChemistryTeacher Notes

You can also measure the progress of your titration using a pH meter, if one is available, and tracking the change in pH vs. volume of titrant added. The endpoint of your titration (the point at which a rapid change in pH occurs) corresponds to the indicator color change in the previously described procedure. The endpoint specifically refers to the pH when the [H3O+] = [OH–]. The volume of titrant added

at the endpoint is called the equivalence point. In the figure below, the endpoint occurs at a pH of 7.0, and the equivalence point occurs at a volume of 20.0 mL of NaOH added. To be clear, the endpoint and equivalence point correspond to the same point on the graph – one just refers more to pH and the other to volume added.

For your particular titration, you want to select an indicator that will change color at endpoint of the titration. (That means you would need to research and find out the endpoint for the titration before you went to the lab, unless you were using a pH meter.) See the graph of pH vs. volume of titrant added, below, as an example.

Suppose an aqueous NaOH solution of unknown concentration is added dropwise to 15.3 mL of a 0.147 M aqueous solution of HCl (with an indicator – bromthymol blue). You find that the end point (= the inflection point on the graph) is reached when you have added 20.9 mL of NaOH. Find the molarity of the NaOH solution.

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IB ChemistryTeacher Notes

Be sure you can interpret titration curves for titrations involving strong acids & strong bases.

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