Acids and Bases

Agenda • “the acid test” • “put the acid on”

• • • •

“acid drop” “do acid”

“acid rain” “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).

Observations *Usually, but not always

Taste Feel (choose slippery or not slippery) pH (# from the key) Litmus (blue or red)

NaOH(aq) Bitter

HCl(aq) Sour

Slippery

Not slippery

14 Blue

1 Red *Cloudy/ white *Yellow Bubbles Bubbles

Phenolphthalein

*Pink

Bromothymol Magnesium Baking soda

*Blue NR NR

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

6

Describe each solution as: 1) acid 2) base or 3) neutral. A. _1_ soda B. _2_ soap C. _1_ coffee D. _1_ wine E. _3_ water F. _1_ grapefruit

7

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

8

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

9

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 Hydrochloric acid + sodium bicarbonate

NaCl + H2O + CO2

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: George Washington: BEFORE acid rain 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 4 Highest volume production of any chemical in the U.S. (approximately 60 billion pounds/year) 4 Used in the production of paper 4 Used in production of fertilizers 4 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 Theories OBJECTIVES:

Compare and contrast acids and bases as defined by the theories of: a) b)

Arrhenius, 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 (18591927)

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 neutral 1. 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 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 Arrhenius 1. Collisions with water molecules 2. The nature of hydrogen ions

Ionization

Cl H

H

+

O H

+ H HO + H

Cl

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

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 (1879-1947) Denmark

Thomas Lowry (1874-1936) England

Brønsted-Lowry Theory of Acids & Bases Conjugate 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), and water is the hydrogen ion donor (acid). This causes the OH1- concentration to be greater than in 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+ + OH1base acid c.a. c.b. • HCl + H2O ↔ H3O1+ + Cl1acid base c.a. c.b. • Amphoteric – a substance that can act as both an acid and base- as water shows

When life goes either way amphoteric (amphiprotic) substances Acting like a base + H+ H2CO3 accepts H+

HCO3-

Acting like an acid - H+ CO3-2 donates H+

Brønsted-Lowry Theory of Acids & Bases

Brønsted-Lowry Theory of Acids & Bases Notice 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

(due to being bonded to the highly electronegative Oxygen)

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

Conjugate Acid-Base Pairs

Conjugate Acid- Base Pairs In other words: When a proton is gained by a BronstedLowry base, the product formed is referred to as the base’s conjugate acid Conjugate Acid Conjugate Base H2O (l) OH-(aq) H2O (l) H3O+(aq) NH3 (aq) NH4+(aq) HCO3CO3-2 H2CO3

HCO3-

Practice problems

Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs: HC2H3O2(aq) + H2O(l)  C2H3O2–(aq) + H3O+(aq) conjugate base conjugate acid acid base conjugate acid-base pairs

OH –(aq) + HCO3–(aq)  CO32–(aq) + H2O(l) base acid conjugate baseconjugate acid conjugate acid-base pairs

Answers: question 18 (a)

(b)

HF(aq) + SO32–(aq)  F–(aq) + HSO3–(aq) conjugate baseconjugate acid acid base conjugate acid-base pairs

CO32–(aq) + HC2H3O2(aq)  C2H3O2–(aq) + HCO3–(aq) base acid conjugate base conjugate acid (c)

conjugate acid-base pairs

H3PO4(aq) + OCl –(aq)  H2PO4–(aq) + HOCl(aq) conjugate base conjugate acid acid base conjugate acid-base pairs

8a)

HCO3–(aq) + S2–(aq)  HS–(aq) + CO32–(aq) acid base conjugate acidconjugate base conjugate acid-base pairs

8b)

H2CO3(aq) + OH –(aq)  HCO3–(aq) + H2O(l) acid base conjugate base conjugate acid conjugate acid-base pairs

11a)

H3O+(aq) + HSO3–(aq)  H2O(l) + H2SO3(aq) conjugate baseconjugate acid acid base conjugate acid-base pairs

11b) OH –(aq) + HSO3–(aq)  H2O(l) + SO32–(aq) conjugate acidconjugate base base acid conjugate acid-base pairs

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. HPO42-____________ d. NH3___________

Strengths of Acids and Bases • OBJECTIVES: – Define strong acids and weak acids.

Strength OBJECTIVES: 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

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

HA

What happens to the HA molecules in solution?

100% dissociation of HA

HA H+

Strong Acid

A-

Would the solution be conductive?

Strong Acid Dissociation (makes 100 % ions)

Partial dissociation of HA

HA H+

Weak Acid

A-

Would the solution be conductive?

HA  H+ + A-

HA H+ A-

Weak Acid At any one time, only a fraction of the molecules are dissociated.

Weak Acid Dissociation (only partially ionizes)

Strength of ACIDS 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.

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+ + OHStrong Bases: Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) Heavy Group 2A metal hydroxides [Ca(OH)2, Sr(OH)2, and Ba(OH)2]

For the following identify the acid and the base as strong or weak . a. Al(OH)3 + HCl  Weak base

Strong acid

b. Ba(OH)2 + HC2H3O2  Strong base Weak acid c. KOH + H2SO4  Strong base Strong acid

d. NH3 + H2O  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 Arrehenius only in water

• Acids – produce H+ • Bases - produce OH-

• Acids – donate H+ Bronsted-Lowry • Bases – accept H+ any solvent

Lewis used in organic chemistry, wider range of substances

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

Acids Arrhenius 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+ + NO3HNO3 + H2O  H3O+ + NO3HCl + H2O  H+ + ClHCl + H2O  H3O+ + Cl-

Bases 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+ + OHNH3 + H2O  NH4+ + OH-

Hydrogen Ions and Acidity OBJECTIVES: • 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

Does pure water conduct electrical current? Water is a very, very, very weak electrolyte. H2O + H2O  H3O+ + OHHow are (H3O+) and (OH-) related? [H3O+][OH-] = 10-14

For pure water: [H3O+] = [OH-] = 10-7M This is neutrality and at 25oC is a pH = 7. 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 ion concentration, [H3O+] or molarity, of a solution. (In other words how many H3O+ ions are in a solution. If there are a lot we assume it is an acid, if there are very few it is a base.)

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

pH

0-14 scale for the chemists 2

3

4

5

acidic (H+) > (OH-) normal rain (CO2) pH = 5.3 – 5.7

6

7

8

9

10

11

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

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

fish populations drop off pH < 6 and to zero pH < 5

natural waters pH = 6.5 - 8.5

12

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 NaOH NaOH  Na + + OH0.001mol 0.001mol Hydroxide will be produced and the [OH-] = 0.001M 2nd step: pOH = -log [0.001] pOH = 3.0 pH = 14.0-3.0 = 11.0

1. 2. 3. 4.

5. 6.

PRACTICE PROBLEM What is the molar concentration of hydronium ion in a solution of pH 8.25? 5.623 x 10-9 M What is the pH of a solution that has a molar pH = 4.0 concentration of hydronium ion of 9.15 x 10-5M? What is the pOH of a solution that has a molar pOH = 4.9 concentration of hydronium ion of 8.55 x 10-10 M? What is the molar concentration of hydronium ion in a solution of pH 2.45? What is the pH of a solution that has a molar concentration of hydronium ion of 3.75 x 10-9 M? What is the pOH of a solution that has a molar concentration of hydronium ion of 4.99 x 10-4 M?

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? Dissolved carbon dioxide lowers the pH CO2 (g) + H2O  H2CO3  H+ + HCO3Atmospheric pollutants from combustion

NO, NO2 + H2O …  HNO3 SO2, SO3 + H2O …  H2SO4 pH < 5.3

both strong acids

Behavior of oxides in water– Group A basic

1A

amphoteric

acidic 3A 4A 5A 6A 7A

2A

Group B

105

107

Db

Bh

basic: Na2O + H2O  2NaOH (O-2 + H2O  2OH-) acidic: CO2 + H2O  H2CO3

8A

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

Digestive system mouth esophagus stomach

small intestine large intestine

pH The biological view in the human body acidic 1

2

3

4

5

6

basic/alkaline 7

8

9

10

11

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 their pH 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

Writing neutralization equations When acids and bases are mixed, a salt forms NaOH + HCl  H2O + NaCl base + acid  water + salt Ca(OH)2 + H2SO4  2H2O + CaSO4 Question: Write the chemical reaction when lithium hydroxide is mixed with carbonic acid. Step 1: write out the reactants LiOH(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 +

Assignment Write balanced chemical equations for these neutralization reactions. Under each compound give the correct IUPAC name. a) b) c) d) e) f)

iron(II) hydroxide + phosphoric acid Ba(OH)2(aq) + HCl(aq) calcium hydroxide + nitric acid Al(OH)3(aq) + H2SO4(aq) ammonium hydroxide + hydrosulfuric acid KOH(aq) + HClO2(aq)

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 to neutralize 22.50 mL of 0.52 M NaOH. What is the molarity ( concentration) of the acid? HCl + NaOH  H2O + NaCl Ma 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?

Mb = 0.773 mol/L

TITRATIONSample Problem: If 37.12 mL of 0.543 M LiOH neutralized 40.50 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 LiOH 2. 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 barium hydroxide solution?

Mb = 0.308 mol/L

Titration problems 1. 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 neutralized 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 neutralize a solution containing 32.0 g NaOH?

Titration problems 1. (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 20.00mL sample of sulfuric acid solution. What is the molarity of the acid? • The reaction equation is: H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2

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 required10.8 to mL neutralize 19.54 mL of 0.00850 M HBr solution? 3. How many mL of 0.998 M H2SO4 must 0.0101 be addedg to neutralize 47.9 mL of 1.233 M KOH? mL to 4. How many milliliters of 1.25 M LiOH must29.6 be added neutralize 34.7 mL of 0.389 M HNO3? 10.8 mL 5. What mass of Sr(OH)2 will be required to neutralize 19.54 mL of 0.00850 M HBr solution? 0.0101 g 6. How many mL of 0.998 M H2SO4 must be added to neutralize 47.9 mL of 1.233 M KOH? 29.6 mL 7. How many milliliters of 0.75 M KOH must be added to neutralize 50.0 mL of 2.50 M HCl 1.

Microtitration Lab: Create an Observation Table for this Experiment Procedure 1: 1. Measure 1-2mL of 0.1M HCl into a 10mL graduated cylinder. Record volume used 2. Add one drop of phenolphthalein 3. Use 0.1M NaOH to determine the volume required to produce a colour change. 4. Record you final volume accurately 5. Repeat the same procedure 2 more times Procedure 2: 1. Titrate an unknown concentration of HCl with the 0.1M NaOH as in procedure 1.