Acids and Bases

Agenda • “the acid test” “acid drop” “acid rain”• “put the acid on” “do acid” “acid head”

• Day 71 – Strong and Weak Acids and Bases Intro• Lesson: • Handouts: 1. Acid/Base Handout • Text: 1. P.462-466, 470-474- Dissociation vs Ionization • Arrhenius and Bronsted-Lowry Definitions of Acids and

Bases • HW: 1. page 475 # 1,2-12•  

Properties of acids and bases• Get 8 test tubes. Rinse all tubes well with water. Add acid

to four tubes, base to the other four.• Touch a drop of base to your finger. Record the feel in the

chart (on the next slide). Wash your hands with water. Repeat for acid.

• Use a stirring rod, add base to the litmus and pH papers (for pH paper use a colour key to find a number). Record results. Repeat for acid.

• Into the four base tubes add: a) two drops of phenolphthalein, b) 2 drops of bromothymol, c) a piece of Mg, d) a small scoop of baking soda. Record results. Repeat for acid.

• Clean up (wash tubes, pH/litmus paper in trash).

BubblesNRBaking sodaBubblesNRMagnesium*Yellow*BlueBromothymol

*Cloudy/ white*PinkPhenolphthalein

RedBlueLitmus (blue or red)114pH (# from the key)

Not slipperySlipperyFeel (choose slippery or not slippery)

SourBitterTasteHCl(aq)NaOH(aq)

Observations*Usually, but not always

6

1. Describe the solution in each of the following as: 1) acid 2) base or 3)neutral.

A. ___soda

B. ___soap

C. ___coffee

D. ___ wine

E. ___ water

F. ___ grapefruit

7

Describe each solution as: 1) acid 2) base or 3) neutral.A. _1_ soda

B. _2_ soapC. _1_ coffeeD. _1_ wineE. _3_ water

F. _1_ grapefruit

8

Identify each as characteristic of an A) acid or B) base

____ 1. Sour taste

____ 2. Produces OH- in aqueous solutions

____ 3. Chalky taste

____ 4. Is an electrolyte

____ 5. Produces H+ in aqueous solutions

9

Identify each as a characteristic of an A) acid or B) base

_A_ 1. Sour taste

_B_ 2. Produces OH- in aqueous solutions

_B_ 3. Chalky taste

A, B 4. Is an electrolyte

_A_ 5. Produces H+ in aqueous solutions

Properties of Acids• They taste sour (don’t try this at home).• They can conduct electricity.

– Can be strong or weak electrolytes in aqueous solution

• React with metals to form H2 gas.

• Change the color of indicators (for example: blue litmus turns to red).

• React with bases (metallic hydroxides) to form water and a salt.

Properties of Acids• They have a pH of less than 7 (more on this concept

of pH in a later lesson)

• They react with carbonates and bicarbonates to produce a salt, water, and carbon dioxide gas

• How do you know if a chemical is an acid?– It usually starts with Hydrogen.– HCl, H2SO4, HNO3, etc. (but not water!)

Acids Affect Indicators, by changing their color

Blue litmus paper turns red in contact with an acid (and red paper stays red).

Acids React with Active Metals

Acids react with active metals to form salts and hydrogen gas:

HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

This is a single-replacement reaction

Acids React with Carbonates and Bicarbonates

HCl + NaHCO3

NaCl + H2O + CO2

Hydrochloric acid + sodium bicarbonate

salt + water + carbon dioxide

An old-time home remedy for relieving an upset stomach

Effects of Acid Rain on Marble(marble is calcium carbonate)

George Washington:BEFORE acid rain

George Washington:AFTER acid rain

Acids Neutralize Bases

HCl + NaOH  → NaCl + H2O

-Neutralization reactions ALWAYS produce a salt (which is an ionic compound) and water.

-Of course, it takes the right proportion of acid and base to produce a neutral salt

Sulfuric Acid = H2SO4 Highest volume production of

any chemical in the U.S. (approximately 60 billion pounds/year)

Used in the production of paper

Used in production of fertilizers

Used in petroleum refining; auto batteries

Nitric Acid = HNO3 Used in the production of

fertilizers Used in the production of

explosives Nitric acid is a volatile acid –

its reactive components evaporate easily

Stains proteins yellow (including skin!)

Hydrochloric Acid = HCl Used in the “pickling” of

steel Used to purify magnesium

from sea water Part of gastric juice, it aids in

the digestion of proteins Sold commercially as

Muriatic acid

Phosphoric Acid = H3PO4 A flavoring agent in

sodas (adds “tart”) Used in the

manufacture of detergents

Used in the manufacture of fertilizers

Not a common laboratory reagent

Acetic Acid = HC2H3O2 (also 

called Ethanoic Acid, CH3COOH)

Used in the manufacture of plastics

Used in making pharmaceuticals

Acetic acid is the acid that is present in household vinegar

Properties of Bases (metallic hydroxides)

• React with acids to form water and a salt.• Taste bitter.• Feel slippery (don’t try this either).• Can be strong or weak electrolytes in aqueous

solution• Change the color of indicators (red litmus turns

blue).

Examples of Bases(metallic hydroxides)

Sodium hydroxide, NaOH (lye for drain cleaner; soap)

Potassium hydroxide, KOH (alkaline batteries)

Magnesium hydroxide, Mg(OH)2 (Milk of Magnesia)

Calcium hydroxide, Ca(OH)2 (lime; masonry)

Bases Affect Indicators

Red litmus paper turns blue in contact with a base (and blue paper stays blue).

Phenolphthalein turns purple in a base.

Bases have a pH greater than 7

Bases Neutralize Acids

Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl.

2 HCl + Mg(OH)2

MgCl2 + 2 H2O

Magnesium salts can cause diarrhea (thus they are used as a laxative) and may also cause kidney stones.

Acid-Base TheoriesOBJECTIVES:Compare and contrast acids and bases as defined by the theories of:

a) Arrhenius, b) Brønsted-Lowry, and

c) Lewis.

Svante Arrhenius• He was a Swedish chemist (1859-1927), and a

Nobel prize winner in chemistry (1903)• One of the first chemists to explain the chemical

theory of the behavior of acids and bases

Svante Arrhenius (1859-1927)

1. Arrhenius Definition - 1887

• Acids produce hydrogen ions (H1+) in aqueous solution (HCl → H1+ + Cl1-)

• Bases produce hydroxide ions (OH1-) when dissolved in water.

(NaOH → Na1+ + OH1-)• Limited to aqueous solutions.• Only one kind of base (hydroxides)• NH3 (ammonia) could not be an Arrhenius base: no

OH1- produced.

Polyprotic Acids?

• Some compounds have more than one ionizable hydrogen to release

• HNO3 nitric acid - monoprotic

• H2SO4 sulfuric acid - diprotic - 2 H+

• H3PO4 phosphoric acid - triprotic - 3 H+

• Having more than one ionizable hydrogen does not mean stronger!

Acids• Not all compounds that have hydrogen are acids.

Water?

• Also, not all the hydrogen in an acid may be released as ions– only those that have very polar bonds are

ionizable - this is when the hydrogen is joined to a very electronegative element

Arrhenius examples...

• Consider HCl = it is an acid!• What about CH4 (methane)?• O (e.g. H2SO4) was originally thought to cause acidic

properties. Later, H was implicated, but it was still not clear why CH4 was neutral.

• CH3COOH (ethanoic acid, also called acetic acid) - it has 4 hydrogens just like methane does…?

Arrhenius’ theory Limitation

Using Arrhenius’ theory the following would be incorrectly classified as neutral1. Compounds of hydrogen polyatomic ions (NaHCO3(aq))

2.Oxides of metals and non metals (CaO(aq) and CO2(g))

3.Bases other than hydroxides (NH3(aq) and Na2CO3(aq))

4.Acids that do not contain hydrogen (Al(NO3)3(aq))

Revised Arrhenius theory

Ionization

+Cl HH

HO

+H

HH O Cl+

Arrhenius made the revolutionary suggestion that some solutions contain ions & that acids produce H3O+ (hydronium) ions in solution.

The revised Arrhenius theory involves two key ideas not considered by Arrhenius1. Collisions with water molecules2. The nature of hydrogen ions

Agenda

• Day 72 – Conjugate Acids and Bases• Lesson: PPT• Handouts: 1. Acid/Base Handout. 2 Conjugate

Acid& Base Worksheet• Text: 1. [page 386-388 old text

photocopy!]• HW: 1. [P.389 # 18, 19 page 392 # 8, 9, 11 old

text photocopy!]

2. Brønsted-Lowry - 1923• A broader definition than Arrhenius• Acid is hydrogen-ion donor (H+ or proton);

base is hydrogen-ion acceptor.• Acids and bases always come in pairs.• HCl is an acid.

– When it dissolves in water, it gives it’s proton to water.

HCl(g) + H2O(l) ↔ H3O+(aq) + Cl-(aq)

• Water is a base; makes hydronium ion.

Johannes Brønsted Thomas Lowry (1879-1947) (1874-1936) Denmark England

Brønsted-Lowry Theory of Acids & BasesConjugate Acid-Base Pairs

General Equation

Why Ammonia is a Base

Ammonia can be explained as a base by using Brønsted-Lowry:

NH3(aq) + H2O(l) ↔ NH41+(aq) + OH1-

(aq)

Ammonia is the hydrogen ion acceptor (base), andwater is the hydrogen ion donor (acid).

This causes the OH1- concentration to be greater thanin pure water, and the ammonia solution is basic

Acids and bases come in pairs• A “conjugate base” is the remainder of the

original acid, after it donates it’s hydrogen ion

• A “conjugate acid” is the particle formed when the original base gains a hydrogen ion.

• Thus, a conjugate acid-base pair is related by the loss or gain of a single hydrogen ion.

• Chemical Indicators? They are weak acids or bases that have a different color from their original acid and base

Acids and bases come in pairs• General equation is:

HA(aq) + H2O(l) ↔ H3O+(aq) + A-

(aq)

• Acid + Base ↔ Conjugate acid + Conjugate base

• NH3 + H2O ↔ NH41+ + OH1-

base acid c.a. c.b.• HCl + H2O ↔ H3O1+ + Cl1-

acid base c.a. c.b.• Amphoteric – a substance that can act as

both an acid and base- as water shows

When life goes either wayamphoteric (amphiprotic) substances

HCO3-

H2CO3 CO3-2

+ H+ - H+

Acting like a base

Acting like an acid

accepts H+ donates H+

Brønsted-Lowry Theory of Acids & Bases

Brønsted-Lowry Theory of Acids & BasesNotice that water is both an acid & a base = amphoteric

Reversible reaction

Organic Acids (those with carbon)Organic acids all contain the carboxyl group, (-COOH), sometimes several of them. CH3COOH – of the 4 hydrogen, only 1 ionizable

The carboxyl group is a poor proton donor, so ALL organic acids are weak acids.

(due to being bonded to the highly electronegative Oxygen)

Conjugate Acid-Base Pairs

Conjugate Acid- Base Pairs

In other words: When a proton is gained by a Bronsted-Lowry base, the product formed is referred to as the base’s conjugate acid

Conjugate Acid Conjugate BaseH2O (l)

H2O (l)

NH4+

(aq)

HCO3-

OH-(aq)

H3O+(aq)

NH3 (aq)

H2CO3

CO3-2

HCO3-

Practice problemsIdentify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs:

acid base conjugate acidconjugate baseHC2H3O2(aq) + H2O(l) C2H3O2

–(aq) + H3O+(aq)

conjugate acid-base pairs

acidbase conjugate acidconjugate baseOH

–(aq) + HCO3–(aq) CO3

2–(aq) + H2O(l)

conjugate acid-base pairs

acid base conjugate acidconjugate baseHF(aq) + SO3

2–(aq) F–(aq) + HSO3–(aq)

conjugate acid-base pairs

acidbase conjugate acidconjugate baseCO3

2–(aq) + HC2H3O2(aq) C2H3O2–(aq) + HCO3

–(aq)

conjugate acid-base pairs

acid base conjugate acidconjugate baseH3PO4(aq) + OCl

–(aq) H2PO4–(aq) + HOCl(aq)

conjugate acid-base pairs

Answers: question 18(a)

(b)

(c)

acid base conjugate baseconjugate acidHCO3

–(aq) + S2–(aq) HS–(aq) + CO32–(aq)

conjugate acid-base pairs

baseacid conjugate acidconjugate baseH2CO3(aq) + OH

–(aq) HCO3–(aq) + H2O(l)

conjugate acid-base pairs

acid base conjugate acidconjugate baseH3O+(aq) + HSO3

–(aq) H2O(l) + H2SO3(aq)

conjugate acid-base pairs

8a)

8b)

11a)

base acid conjugate baseconjugate acidOH

–(aq) + HSO3–(aq) H2O(l) + SO3

2–(aq)

conjugate acid-base pairs

11b)

What is the conjugate base of the following substances?

a. H2O ________________ b. NH4

+________________ c. HNO2_______________ d. HC2H3O2_________________ What is the conjugate acid of the following 

substances?a. HCO3

-__________________ b. H2O____________ c. HPO4

2-____________ d. NH3___________

Strengths of Acids and Bases

• OBJECTIVES:– Define strong acids and weak acids.

StrengthOBJECTIVES: Define strong acids and weak acids.

•Acids and Bases are classified according to the degree to which they ionize in water:

– Strong are completely ionized in aqueous solution; this means they ionize 100 %

– Weak ionize only slightly in aqueous solution

Strength is very different from Concentration

Strength• Strong – means it forms many ions when dissolved

(100 % ionization)

• Mg(OH)2 is a strong base- it falls completely apart (nearly 100% when dissolved). – But, not much dissolves- so it is not

concentrated

HA

Let’s examine the behavior of an acid, HA, in aqueous solution.

What happens to the HA molecules in solution?

HA

H+

A-

Strong Acid

100% dissociation of HA

Would the solution be conductive?

Strong Acid Dissociation (makes 100 % ions)

HA

H+

A-

Weak Acid

Partial dissociation of HA

Would the solution be conductive?

HA

H+

A-

Weak Acid

HA H+ + A-

At any one time, only a fraction of

the molecules are dissociated.

Weak Acid Dissociation(only partially ionizes)

1.  Binary or hydrohalic acids – HCl, HBr, and HI “hydro____ic acid” are strong acids.  Other binary acids are weak acids (HF and H2S). Although the H-F bond is very polar, the bond is so strong (due to the small F atom) that the acid does not completely ionize.

Strength of ACIDS

2.  Oxyacids – contain a polyatomic ion a. strong acids (contain 2 or more oxygen per

hydrogen)HNO3 – nitric from nitrate

H2SO4 - sulfuric from sulfate

HClO4 - perchloric from perchlorate

b. weak acids (acids with l less oxygen than the “ic” ending

HNO2 – nitrous from nitrite

H3PO3 - phosphorous   from phosphite

H2SO3 - sulfurous from sulfite

HClO2 - chlorous from chlorite

c. weaker acids (acids with “hypo ous” have less oxygen than the “ous” ending

HNO - hyponitrous H3PO2 - hypophosphorus

HClO - hypochorous

d. Organic acids – have carboxyl group   -COOH - usually weak acids

HC2H3O2 - acetic acid

C7H5COOH - benzoic acid

Strength of Bases 

Strong Bases:  metal hydroxides of Group I and II metals (except Be) that are soluble in water and dissociate (separates into ions) completely in dilute aqueous solutions

Weak Bases: a molecular substance that ionizes only slightly in water to produce an alkaline (basic) solution (ex. NH3)

What is a strong Base?

A base that is completely dissociated in water (highly soluble).

NaOH(s) Na+ + OH-

Strong Bases:

Group 1A metal hydroxides(LiOH, NaOH, KOH,RbOH, CsOH)

Heavy Group 2A metal hydroxides[Ca(OH)2, Sr(OH)2, andBa(OH)2]

For the following identify the acid and the base as strong or weak .

a. Al(OH)3 + HCl

b. Ba(OH)2 + HC2H3O2

c. KOH + H2SO4

d. NH3 + H2O

Weak base Strong acid

Weak acid

Strong acid

Strong base

Strong base

Weak base Weak acid

Strength vs. Concentration• The words concentrated and dilute tell how much

of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume

• The words strong and weak refer to the extent of ionization of an acid or base

• Is a concentrated, weak acid possible?

3. Lewis Acids and Bases• Gilbert Lewis focused on the donation or

acceptance of a pair of electrons during a reaction• Lewis Acid - electron pair acceptor• Lewis Base - electron pair donor• Most general of all 3 definitions; acids don’t even

need hydrogen!

Gilbert Lewis (1875-1946)

Gilbert Lewis (1875-1946)

Summary: Definitions

• Acids – produce H+

• Bases - produce OH-

• Acids – donate H+

• Bases – accept H+

• Acids – accept e- pair• Bases – donate e- pair

Arrehenius

Bronsted-Lowry

Lewis

only in water

any solvent

used in organic chemistry,wider range of substances

Acids BasesArrhenius Acid: donates (or

produces) hydronium ions (H3O+) in water or hydrogen ions (H+) in water

Bronsted-Lowry Acid: donates a proton (H+) in water, H3O+ has an extra H+, if it donated it to another molecule it would be H2O (page 467)

HNO3 + H2O H+ + NO3-

HNO3 + H2O H3O+ + NO3-

HCl + H2O H+ + Cl-

HCl + H2O H3O+ + Cl-

Arrhenius Base: donates (or produces) hydroxide ions (OH-) in water

Bronsted – Lowry Base: accepts a proton in water, OH- needs an extra H+ if it accepts one from another molecule it would be H2O (page 468)

KOH + H2O K+ + OH-

NH3 + H2O NH4+ + OH-

Hydrogen Ions and AcidityOBJECTIVES:•Describe how [H1+] and [OH1-] are related in an aqueous solution.•Classify a solution as neutral, acidic, or basic given the hydrogen-ion or hydroxide-ion concentration. •Convert hydrogen-ion concentrations into pH values and hydroxide-ion concentrations into pOH values.• Describe the purpose of an acid-base indicator.

Agenda

• Day 73 – pH Calculations • Lesson: PPT• Handouts: 1. pH Handout, 2. pH Calculations

Worksheet• Text: 1. [page 368-374 old text

photocopy!]• HW: 1. [page # 371 # 2, 3, 4, 6 old text

photocopy!]

Hydrogen Ions from Water• Water ionizes, and forms ions:

H2O + H2O↔ H3O1+ + OH1-

• Called the “self ionization” of water• Occurs to a very small extent:

[H3O1+ ] = [OH1-] = 1 x 10-7 M

• Since they are equal, a neutral solution results from water

Kw = [H3O1+ ] x [OH1-] = 1 x 10-14 M2

• Kw is called the “ion product constant” for water

Water Equilibrium

H2O + H2O H3O+ + OH-

Does pure water conduct electrical current?

[H3O+][OH-] = 10-14

For pure water: [H3O+] = [OH-] = 10-7M

This is neutrality and at 25oC is a pH = 7.

Water is a very, very, very weak electrolyte.

How are (H3O+) and (OH-) related?

water

Lone Hydrogen ions do not exist by themselves in solution. H+ is always bound to a water molecule to form a hydronium ion

Ion Product Constant• H2O ↔ H3O1+ + OH1-

• Kw is constant in every aqueous solution:

[H3O+] x [OH-] = 1 x 10-14 M2

• If [H3O+] > 10-7 then [OH-] < 10-7

• If [H3O+] < 10-7 then [OH-] > 10-7

• If we know one, other can be determined• If [H3O+] > 10-7 , it is acidic and [OH-] < 10-7

• If [H3O+] < 10-7 , it is basic and [OH-] > 10-7

– Basic solutions also called “alkaline”

The pH concept – from 0 to 14

• pH = pouvoir hydrogene (Fr.) “hydrogen power”• definition: pH = -log[H3O+]

• in neutral pH = -log(1 x 10-7) = 7• in acidic solution [H3O+] > 10-7

• pH < -log(10-7)– pH < 7 (from 0 to 7 is the acid range)– in base, pH > 7 (7 to 14 is base range)

pH Scale [ ] brackets mean concentration or Molarity

The pH scale indicates the hydronium ionconcentration, [H3O+] or molarity, of a solution. (Inother words how many H3O+ ions are in a solution. Ifthere are a lot we assume it is an acid, if there are veryfew it is a base.)

pH

2 3 4 5 6 7 8 9 10 11 12

neutral @ 25oC(H+) = (OH-)

distilled water

acidic(H+) > (OH-)

basic or alkaline(H+) < (OH-)

natural waters pH = 6.5 - 8.5

normal rain (CO2)pH = 5.3 – 5.7

acid rain (NOx, SOx)pH of 4.2 - 4.4 in

0-14 scale for the chemists

fish populationsdrop off pH < 6 and to zero pH < 5

pH Scale

• A change of 1 pH unit represents a tenfold change in the acidity of the solution.

• For example, if one solution has a pH of 1 and a second solution has a pH of 2, the first solution is not twice as acidic as the second—it is ten times more acidic.

Calculating pOH

• pOH = -log [OH-] • [H+] x [OH-] = 1 x 10-14 M2

• pH + pOH = 14• Thus, a solution with a pOH less than 7 is basic; with

a pOH greater than 7 is an acid• Not greatly used like pH is.

pH and Significant Figures

• For pH calculations, the hydrogen ion concentration is usually expressed in scientific notation

• [H1+] = 0.0010 M = 1.0 x 10-3 M, and 0.0010 has 2 significant figures

• the pH = 3.00, with the two numbers to the right of the decimal corresponding to the two significant figures

Example Problems:1. What is the pH of a 0.001M NaOH solution?

1st step: Write a dissociation equation for NaOHNaOH Na + + OH- 0.001mol 0.001molHydroxide will be produced and the [OH-] = 0.001M

2nd step: pOH = -log [0.001] pOH = 3.0 pH = 14.0-3.0 = 11.0

PRACTICE PROBLEM1. What is the molar concentration of hydronium ion in

a solution of pH 8.25?2. What is the pH of a solution that has a molar

concentration of hydronium ion of 9.15 x 10-5M?3. What is the pOH of a solution that has a molar

concentration of hydronium ion of 8.55 x 10-10 M?4. What is the molar concentration of hydronium ion in

a solution of pH 2.45?5. What is the pH of a solution that has a molar

concentration of hydronium ion of 3.75 x 10-9 M?6. What is the pOH of a solution that has a molar

concentration of hydronium ion of 4.99 x 10-4 M?

5.623 x 10-9 M

pH = 4.0

pOH = 4.9

2. What is the pH of a 3.4X10-5M H2SO4 solution?

3. What is the pH of a solution if the pOH = 5?

4. What is the pH of a 10 liter KOH solution if 5.611 grams of KOH were used to prepare the solution?

5. What is the pOH of a 1.1X10-5M HNO3 solution?

6. If the pH of a KOH solution is 10.75, what is the molar concentration of the solution? What is the pOH? What is the [H+]?

The pH of a strong acid cannot be greater than 7. If the acid concentration [H3O+] is less than 1.0X10-7, the water becomes the important source of [H3O+] or [H+] and the pH is 7.00. Just remember to check if you answer is reasonable!

7. What is the pH of a 2.5X10-10M HCl solution? 8. What is the pH of a 1.0X10-11M HNO3 solution?

What is acid rain?

CO2 (g) + H2O H2CO3 H+ + HCO3-

Dissolved carbon dioxide lowers the pH

Atmospheric pollutants from combustion

NO, NO2 + H2O … HNO3

SO2, SO3 + H2O … H2SO4

bothstrong acids

pH < 5.3

105

Db107

Bh

Behavior of oxides in water– Group Abasic amphoteric acidic

3A 4A 5A 6A 7A

1A

2A

8A

Group B

basic: Na2O + H2O 2NaOH (O-2 + H2O 2OH-)

acidic: CO2 + H2O H2CO3

Measuring pH• Why measure pH?

Everyday solutions we use - everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc.

• Sometimes we can use indicators, other times we might need a pH meter

pH in the Digestive System

• Mouth-pH around 7. Saliva contains amylase, an enzyme which begins to break carbohydrates into sugars.

• Stomach- pH around 2. Proteins are broken down into amino acids by the enzyme pepsin.

• Small intestine-pH around 8. Most digestion ends. Small molecules move to bloodstream toward cells that use them

mouth

esophagus

stomach

small intestine

large intestine

Digestive system

pH

1 2 3 4 5 6 8 9 10 11

The biological view in the human body

gastric juice

urinesalivacerebrospinal fluid

bloodpancreatic juice

bileacidic basic/alkaline

7

How to measure pH with wide-range paper

1. Moisten the pH indicator paper strip with a few drops of solution, by using a stirring rod.

2.Compare the color to the chart on the vial – then read the pH value.

Some of the many pH

Indicators and theirpH range

Acid-Base Indicators• Although useful, there are limitations to indicators:

– usually given for a certain temperature (25 oC), thus may change at different temperatures

– what if the solution already has a color, like paint?

– the ability of the human eye to distinguish colors is limited

Acid-Base Indicators• A pH meter may give more definitive results

– some are large, others portable– works by measuring the voltage between two

electrodes; typically accurate to within 0.01 pH unit of the true pH

– Instruments need to be calibrated

Neutralization Reactions

• OBJECTIVES:– Define the products of an acid-base reaction.– Explain how acid-base titration is used to

calculate the concentration of an acid or a base.

– Explain the concept of equivalence in neutralization reactions.

Agenda

• Day 74 – Acid & Base Titration - Stoichiometry/pH Calculations

• Lesson: PPT• Handouts: 1. Titration Handout 2. Titration

Problems Worksheet • Text: 1. P. 476- 484 -Titration• HW: 1. P. 485 # 1-13

Acid-Base Reactions• Acid + Base → Water + Salt• Properties related to every day:

– antacids depend on neutralization– farmers adjust the soil pH– formation of cave stalactites– human body kidney stones from insoluble salts

Acid – Base reactions

• Each salt listed in this table can be formed by the reaction between an acid and a base.

Acid-Base Reactions• Neutralization Reaction - a reaction in which an

acid and a base react in an aqueous solution to produce a salt and water:

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

H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2 H2O(l)

According to the Bronsted-Lowry theory, in a neutralization reaction a proton is transferred form the strongest acid to the strongest base

Acid – Base Reactions

• A reaction between an acid and a base is called neutralization. An acid-base mixture is not as acidic or basic as the individual starting solutions.

Titration- Stoichiometry• Titration is the process of adding a known amount

of solution of known concentration to determine the concentration of another solution

• Remember? - a balanced equation is a mole ratio

• The equivalence point is when the moles of hydronium ions is equal to the moles of hydroxide ions (= neutralized!)

Titration• The concentration of acid (or base) in solution can

be determined by performing a neutralization reaction– An indicator is used to show when

neutralization has occurred– Often we use phenolphthalein- because it is

colorless in neutral and acid; turns pink in base

Steps - Neutralization reaction#1. A measured volume of acid of unknown

concentration is added to a flask#2. Several drops of indicator added#3. A base of known concentration is slowly added,

until the indicator changes color; measure the volume.

Neutralization• The solution of known concentration is called the

standard solution– added by using a burette– Figure 1, page 476

• Continue adding until the indicator changes color– called the “end point” of the titration– Go over Sample Problem1 and 2 , page 482

Question: Write the chemical reaction when lithium hydroxide is mixed with carbonic acid.

Step 1: write out the reactantsLiOH(aq) + H2CO3(aq)

Step 2: determine products … H2O and Li1(CO3)2

LiOH(aq) + H2CO3(aq) Li2CO3(aq) + H2O(l) Step 3: balance the equation

2LiOH(aq) + H2CO3(aq) Li2CO3(aq) + 2H2O(l) lithium hydroxide + carbonic acid lithium carbonate + water

Writing neutralization equationsWriting neutralization equationsWhen acids and bases are mixed, a salt forms

NaOH + HCl H2O + NaCl base + acid water + saltCa(OH)2 + H2SO4 2H2O + CaSO4

a) iron(II) hydroxide + phosphoric acidb) Ba(OH)2(aq) + HCl(aq)c) calcium hydroxide + nitric acidd) Al(OH)3(aq) + H2SO4(aq)e) ammonium hydroxide + hydrosulfuric acidf) KOH(aq) + HClO2(aq)

Assignment

Write balanced chemical equations for these neutralization reactions. Under each compound give the correct IUPAC name.

a) 3Fe(OH)2(aq) + 2H3PO4(aq) Fe3(PO4)2(aq) + 6H2O(l) iron(II) hydroxide + phosphoric acid iron (II) phosphate

b) Ba(OH)2(aq) + 2HCl(aq) BaCl2 (aq) + 2H2O(l) barium hydroxide + hydrochloric acid barium chloride

c) Ca(OH)2(aq) + 2HNO3(aq) Ca(NO3)2(aq) + 2H2O(l) calcium hydroxide + nitric acid calcium nitrate

d) 2Al(OH)3(aq) + 3H2SO4(aq) Al2(SO4)3(aq) + 6H2O(l) aluminum hydroxide + sulfuric acid aluminum sulfate

e) 2NH4OH(aq) + H2S(aq) (NH4)2S(aq) + 2H2O(l) ammonium hydroxide+ hydrosulfuric acid ammonium sulfide

f) KOH(aq) + HClO2(aq) KClO2(aq) + H2O(l) potassium hydroxide + chlorous acid potassium chlorite

TITRATION- MAVA = MBVB

Sample Problem: Suppose 75.00 mL of hydrochloric acid was required toneutralize 22.50 mL of 0.52 M NaOH. What is themolarity ( concentration) of the acid?

Mb = 0.773 mol/L

HCl + NaOH H2O + NaClMa Va = Mb Vb rearranges to Ma = Mb Vb / Va

so Ma = (0.52 M) (22.50 mL) / (75.00 mL) = 0.16 M

Now you try:2.If 37.12 mL of 0.843 M HNO3 neutralized 40.50 mL of KOH, what is the molarity of the base?

TITRATION- Sample Problem: If 37.12 mL of 0.543 M LiOH neutralized40.50 mL of H2SO4, what is the molarity of the acid?

2 LiOH + H2SO4 Li2SO4 + 2 H2O

1. First calculate the moles of base:0.03712 L LiOH (0.543 mol/1 L) = 0.0202 mol LiOH2. Next calculate the moles of acid:0.0202 mol LiOH (1 mol H2SO4 / 2 mol LiOH)= 0.0101 mol

H2SO4

3. Last calculate the Molarity: Ma = n/V = 0.010 mol H2SO4 / 0.4050 L = 0.248 M

Now you try it: If 20.42 mL of Ba(OH)2 solution was used to titrate29.26 mL of 0.430 M HCl, what is the molarity of the bariumhydroxide solution?

Mb = 0.308 mol/L

Titration problems1. What volume of 0.10 mol/L NaOH is needed

to neutralize 25.0 mL of 0.15 mol/L H3PO4?

2. 25.0 mL of HCl(aq) was neutralized by 40.0 mL of 0.10 mol/L Ca(OH)2 solution. What was the concentration of HCl?

3. A truck carrying sulfuric acid is in an accident. A laboratory analyzes a sample of the spilled acid and finds that 20 mL of acid is neutral-ized by 60 mL of 4.0 mol/L NaOH solution. What is the concentration of the acid?

4. What volume of 1.50 mol/L H2S will neutral-ize a solution containing 32.0 g NaOH?

Titration problems1. (3)(0.15 M)(0.0250 L) = (1)(0.10 M)(VB)

VB= (3)(0.15 M)(0.0250 L) / (1)(0.10 M) = 0.11 L

2. (1)(MA)(0.0250 L) = (2)(0.10 M)(0.040 L)

MA= (2)(0.10 M)(0.040 L) / (1)(0.0250 L) = 0.32 M

3. Sulfuric acid = H2SO4

(2)(MA)(0.020 L) = (1)(4.0 mol/L)(0.060 L)

MA = (1)(4.0 M)(0.060 L) / (2)(0.020 L) = 6.0 M

4. mol NaOH = 32.0 g x 1 mol/40.00 g = 0.800 (2)(1.50 mol/L)(VA) = (1)(0.800 mol)

VA= (1)(0.800 mol) / (2)(1.50 mol/L) = 0.267 L

Molarity and Titration• A student finds that 23.54 mL of a 0.122 M

NaOH solution is required to titrate a 30.00-mL sample of hydr acid solution. What is the molarity of the acid?

• A student finds that 37.80 mL of a 0.4052 M NaHCO3 solution is required to titrate a 20.00-mL sample of sulfuric acid solution. What is the molarity of the acid?

• The reaction equation is:H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2

1. How many milliliters of 1.25 M LiOH must be added to neutralize 34.7 mL of 0.389 M HNO3?

2. What mass of Sr(OH)2 will be required to neutralize 19.54 mL of 0.00850 M HBr solution?

3.How many mL of 0.998 M H2SO4 must be added to neutralize 47.9 mL of 1.233 M KOH?

4. How many milliliters of 1.25 M LiOH must be added to neutralize 34.7 mL of 0.389 M HNO3?

5. What mass of Sr(OH)2 will be required to neutralize 19.54 mL of 0.00850 M HBr solution?

6. How many mL of 0.998 M H2SO4 must be added to neutralize 47.9 mL of 1.233 M KOH?

7. How many milliliters of 0.75 M KOH must be added to neutralize 50.0 mL of 2.50 M HCl

10.8 mL

0.0101 g

29.6 mL

10.8 mL

0.0101 g

29.6 mL