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Acids and Bases• What’s an Acid?

– Acids called “Sauerstoff” in German– Biting or tart taste (lemons, Limes, vinegar)– Surplus of hydrogen (hydronium) ions

• What’s a Base?– Bases tend to be bitter (ammonia, carbonate)– Anti-Acid, hydroxide is opposite of hydrogen

• “Tums” the drug store antacid

– Hydrogen ion deficiency, surplus of hydroxide

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Acids

• Acidic Behavior– Sour taste (lemons, vinegar, carbonated drinks)

• HCl required for digestion

– Acids dissolve most metals (Iron, Aluminum, Zinc)• But not gold, silver, platinum

– Dissolves some rocks (limestone, other carbonates)• Often evolves gasses during reaction

– Hydrogen from metals, Carbon dioxide from carbonate.– Sodium Bicarbonate used in lab to neutralize acids

» NaHCO3 + H+ CO2 (gas) + Na+ + H2O

– Hydrogen Ion is responsible for acidity• Conceptually treated as a free proton

– Free nuclear particles don’t remain unattached for long

– Acidic ion structure is H3O+, called “Hydronium” ion

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Bases

• Base Behavior– Bitter or bad taste (ammonia)– Most properties the direct opposite of acids– However, some metals also dissolved by base

• Don’t make soap in an aluminum pan!– Soap is made from Lye and vegetable oil (or animal fat)– But there is a competing reaction in presence of Aluminum,

which happened in my Junior High School science class

– 2Al + 6NaOH 3 H2 (gas) + 2Na3AlO3 (aq)…and pan went away , leaving a nasty mess and embarrassed instructor.

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3 Definitions of Acids and Bases• Simplest concept = Arrhenius model

– Acidic when [H3O+] > [OH-]– Basic when [OH-] > [H+]– Requires Water to form ions in solution

• More common = BrØnsted model– Acids donate protons [H+] to transfer– Bases [OH-] accept protons (NH3 + H+ NH4

+)– Does NOT require water, OK for organics

• Most General = Lewis model– Acids are electron pair ACCEPTORS– Bases are electron pair DONORS– Applicable anywhere

Arrhenius model• First proposed that salts dissolving in

water create ions which can carry an electric current

• Proposed his theory in 1884 that acids provide hydrogen ions in solution, bases provide hydroxide ions

• First in 1896 to relate carbon dioxide levels in atmosphere to surface temperature of earth, he predicted global warming due to burning of fossil fuels.

• Created concept of activation energy in 1889, that most reactions require overcoming an energy barrier.

• Created over 50 concepts we use today, although radical for his time.

Copyright © 2010 Pearson Education, Inc.

Chapter Ten 6

Acids and Bases in Water• An acid produces hydrogen ions, H+, when

dissolved in water. (Arrhenius Definition)• Hydronium ion: The H3O+ ion formed when an

acid reacts with water.

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H3O is Trigonal PyramidTetrahedral …but missing one chemical bond

(EPG = Electron Pair Geometry)

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Hydrated HydroniumHydronium picks up more water via Hydrogen Bonding

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Arrhenius Limitations• Only works with water involved

– Limited to hydrogen ions and hydroxide ions• Relies on concentrations of ions in solutions

– Does not handle non-aqueous situations• Cannot explain HCl(g) + NH3(g) NH4Cl(s)

– No Hydrogen or hydroxide ions in these gases– Same result in aqueous or gas phase reactions

• Organic reactions also not addressed

– We needed a better model !

Bronsted & Lowry

Bronsted-Lowry

• Both came up with idea of proton transfer at the same time as acid-base definition.

• Bronsted “protonic theory” in 1923– Hydrogen atom in acid loses an electron and

becomes a “proton donor”– Hydroxide ion called “proton receiver” or base– Hydrogen & hydroxide combine forming water

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Brønsted-Lowry Acids and Bases• Brønsted–Lowry acids donate H+ ions (protons)• Monoprotic acids can donate 1 H+ ion,

– diprotic acids can donate 2 H+ ions, and – triprotic acids can donate 3 H+ ions.

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• A Brønsted–Lowry base accepts H+ ions.• An acid–base reaction is where a proton is

transferred. Reaction need not occur in water.

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More on Brønsted-Lowry

• Although not the most general theory, Brønsted-Lowry definition is the most widely used model.

• The strength of an acid by this definition is the stability of hydronium and the solvated conjugate base upon dissociation.

• Increasing or decreasing stability of the conjugate base will increase or decrease the acidity of a compound.

• This concept of acidity is used frequently for organic acids such as carboxylic acid.

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BrØnsted Limitations

• Relies on proton transfer– Works well in most common situations– More general than Arrhenius– Also limited in scope

• Does not handle neutralization intuitively– Products are conjugates, not salts and water

• Not all reactions require hydrogen ion– BF3 + NH3 BF3NH3 … no water, no protons

• Need a better model ! … (again)

Gilbert Norton Lewis

G. N. Lewis contributions• Developed electron dot model in 1902,using

cubes surrounding atoms with 8 corners for electrons, later published in 1916

• First described the covalent bond in 1916• Formulated the electron pair theory of acid-base

reactions in 1923– Lewis acid is electron pair acceptor– Lewis base is electron pair donor

• Named the “photon” as smallest unit of radiant energy in 1926

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Lewis Reaction, a 3-D view

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Soda Pop … which model works?

• CO2(g) + H2O(l) H2CO3(aq) – How do we explain this reaction?

• Ahrennius does NOT work, carbon dioxide is a gas … no hydogen or hydroxide ions.

• Bronsted does NOT work, no proton exchange between water and a gas

• The Lewis acid-base theory can be used to explain why CO2 dissolves in water.

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Soda Water FormationThe water molecule acts as an electron-pair donor, or Lewis base. The electron-pair acceptor is the carbon atom in CO2.

When carbon atom bonds to a pair of water molecule electrons, it no longer needs double bonds with both of it’s oxygen atoms. The net result of the reaction between CO2 and water is carbonic acid, H2CO3, in carbonated drinks

Which model to use?

• Always use the simplest one which works

• Many or most reactions are in water– Water reactions handled by Arrenihius, – The simplest and most appropriate for us

• Bronsted handles no water reactions– Limited due to proton transfer

• Lewis the most versatile, more complex– Good for gases, mixed phase reactions.

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Summary Definitions of Acids and Bases

• Simplest concept = Arrhenius model– Acidic when [H+] > [OH-]– Basic when [OH-] > [H+]– Requires Water

• Most common = BrØnsted model– Acids have protons [H+] available to transfer– Bases [OH-] accept protons (H+ OH- HOH)– Does NOT require water, OK for organics

• Most General = Lewis model– Acids are electron pair ACCEPTORS– Bases are electron pair DONORS– Applicable anywhere

• We tend to use Arrhenius model in practice– Hydrogen and hydroxide ions in water– Neutralization yields water (not conjugates)– Reactions in solutions most common

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Common Acids

• Sulfuric acid, H2SO4 , is the largest quantity industrial chemical, and used in automobile batteries.

• Hydrochloric acid, HCl, is “stomach acid” in the digestive systems of most mammals.

• Phosphoric acid, H3PO4 , is used in fertilizers, and in soft drinks providing a tart taste.

• Nitric acid, HNO3 , is a strong oxidizing agent used for making fertilizers, explosives, many applications

• Acetic acid, CH3CO2H, acidic component of vinegar, natural product from oxidizing alcohol (turns sour)

• Formic acid, HCOOH, simplest organic acid, main constituent of bee and hornet stings.

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Common Bases• Sodium hydroxide, NaOH, or lye, is used in the

production of aluminum, glass, and soap. Drain cleaners often contain NaOH because it reacts with the fats and proteins found in grease and hair.

• Calcium hydroxide, Ca(OH)2 , or slaked lime, is made by treating lime (CaO) with water. It is used in mortars and cements. An aqueous solution is often called limewater.

• Magnesium hydroxide, Mg(OH)2 , or milk of magnesia, is an additive in foods, toothpaste, and many over-the-counter medications. Many antacids contain magnesium hydroxide.

• Ammonia, NH3 , is used primarily as a fertilizer. A dilute solution of ammonia is frequently used around the house as a glass cleaner.

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Water is BOTH acid or base

• Water can dissociate into ions– H2O H+ + OH-

– This reaction is very slight– Acid and base in water = 10-7 mole/liter

• Addition of other materials upsets balance– Acids such as HCl create big surplus of H+

• Acidic solutions are the result

– Bases such as NH4OH create surplus of OH-

• Basic solutions are the result

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Relative Strength• Some acids are recognized as “Strong”

– Sulfuric & Phosphoric Acids dissolve rust– Nitric Acid dissolves almost anything

• Others are “Weak” in everyday sense– Vinegar on salads, lemonade, soda-water

• Same idea applies to bases– Lye is strong …Drano (Lye+Aluminum)– Sodium Carbonate and ammonia are weak

• How about the “in between” material– Weak or diluted acids and bases

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Strong Acids• Highly dissociated ions in water

– Most of acid is dissociated into hydrogen ions• >50% dissociated is arbitrary definition of “strong”• “strong” is not an indicator of chemical reactivity

– Seven ordinary strong acids• Halogen Acids

– Hydrochloric, HCl swimming pools, etch steel, remove rust– Hydrobromic, HBr less aggressive than HCl, more costly– Hydroiodic, Hi

• Sulfuric, H2SO4 Auto batteries• Nitric, HNO3 manufacture of explosives• Chlorine Oxy-acids

– Chloric, HCLO3 strong Oxidizer– Perchloric, HClO4 strong Oxidizer

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Acid-Base strength

• Acid strength– “Strong” Acids highly dissociated

• Lots of Hydrogen Ion available• Same as strong electrolytes (70% dissociated)

• As acid consumed, MORE dissociates into H+

– Dissociation continues until all acid ion is made available and consumed

– “Weak” Acids only modestly dissociated• Equilibrium between dissociated and bonded• Weak electrolytes are weak acids and vice versa • Acetic acid 1.3% dissociated in 1M solution

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Strong Electrolytes

• Strong Acids/Bases, electrolytes– Highly Ionized

• HCl ≈ 80% ionized in 0.1M solution

• BIG arrow in direction of dissociation • Little arrow in direction of recombination

– Strong electrolytes are good electrical conductors• Auto batteries, dry cells, laptop batteries

– Zinc Chloride, Manganese Dioxide, “gel-cells”

• Dangerous to be between power line and ground– Electrical system uses “ground” as return path

– Electrocution potential between puddle and power line

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Mineral Acids

• All are polar, soluble in water– Simple combinations, small molecules

• Binary Inorganic halogen acids• HCl, HBr, HF, HI• HCl is “stomach acid” for digestion

• Oxyacids, oxygen containing anions• HNO3, H2SO4, H3PO4

• Used in car batteries, etc.

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Organic Acids• Many found in plants and animals

– Most are “carboxylic” acids

• Fruit & Vegitable acids– Malic Acid (apples)– Tartaric Acid (grapes)– Citric Acid (oranges, lemons, limes)– Oxalic Acid (rhubarb)

• Bodily Acids– Lactic Acid (milk)– Ascorbic Acid (vitamin C)

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Carboxylic Acid StructureWe are interested in –COOH, the acidic part of molecule“R” represents rest of molecule to which -COOH attached

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Carboxylic acids are characterized by the presence of at least one carboxyl group. The general formula is R-COOH, where R is a mono-valent functional group. A carboxyl formation consists of a carbonyl (RR'C=O) and a hydroxyl (R-O-H), which has the formula -C(=O)OH, usually written as -COOH or -CO2H

Carboxylic acids are Brønsted-Lowry acids because they donate proton (H+). They are the most common type of organic acid. Among the simplest examples are formic acid H-COOH, which occurs in ants, and acetic acid CH3-COOH, which gives vinegar its sour taste. Acids with two or more carboxyl groups are called di-carboxylic, tri-carboxylic, etc. The simplest di-carboxylic example is oxalic acid (COOH)2, which is just two connected carboxyls. Mellitic acid is an example of a hexa-carboxylic acid. Other important natural examples are citric acid (in lemons) and tartaric acid (in tamarinds).

Salts and esters of carboxylic acids are called carboxylates (-ic -ate). When a carboxyl group is de-protonated, its conjugate base, a carboxylate anion is formed. Carboxylic acids are typically weak acids, they only partially dissociate into H+ cations and RCOO– anions in aqueous solution. At room temperature, only 0.4% of all acetic acid molecules are dissociated.

Carboxylic acids often have strong odors. Most common are acetic acid (vinegar) and butanoic acid (rancid butter). On the other hand, esters of carboxylic acids tend to have pleasant odors and many are used in perfumes.

Ionization of Carboxylic AcidTerminal H of -COOH group dissociates to hydrogen ion

Balance becomes carboxylate ion (conjugate base)

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More examples of organic acids

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Chemistry in poetry(Poetry in chemistry?)

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Acids and Bases

• For our experiment, Arrhenius model– Requires water, that’s what we’re using– Treats acid and base as simple ions

• Water is somewhat unusual– Dissolves ions of other materials– Ions give solution acidic or basic properties– Water only slightly dissociates into ions itself

What is pH• How to handle wide range of values

– Earthquakes, Tsunamis– Radiation, sound– Acid and base strength

• Usual moles/liter is often inconvenient– Huge range of + and – exponents– What if we just use the exponents instead?– Definition: pH = - log[H+]

• Log of an exponent is the exponent itself• For acid 10^-5 moleH+/liter - log (10^-5) = pH 5• For neutral 10^-7 moleH+/Liter -log(10^-7) = pH 7 • For basic 10^-14 mole +//iter -log(10^-14) = pH 14

• Each pH unit changes concentration by 10X

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Use of LogarithmsEssential feature is scale compression

Huge range turned into set of small numbersSame idea for Earthquakes and Tsunamis

pH• Definition: pH = - log[H+]

– Scale compressed using logarithms– pH fully equivalent to moles/liter of H+ ion

• Example– Lemon Juice is 0.01 mole/Liter for H+

• pH = - log[0.01] = - - 2.0 = +2.0

• Applications– Stomach acid, fruit juice

• Calculations– [H+] to pH, pH to [H+]

Logarithms• Base

– Number being raised to an exponent power

• Exponent– How many times base is multiplied by itself

• Number– Result of base raised to an exponent

• Example– Base of 10 with exponent 3

• 10*10*10 = 103 = 1,000

• Logarithm is the value of the exponent– Log 1000 = 3 (by inspection)– Log 1500 = 3.17 (use a calculator or tables)

Note the scale compressionlog(1)=0, log(10)=1, log(100)=2, log(1000)=3, etc.

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pH math

• pH definition: pH = –log [H+]

• pH from known value of [H+]– If H+ = 10^-5 molar, pH = -log[10^-5] = 5– If H+ = 1.5E-5 molar, pH = -log[1.5E-5] = 4.8

• [H+] from known value of pH– Must use “anti-logs” or exponentiation– pH = 5, [H+] = 10^-5 molar (by inspection)

– pH =4.8, [H+] = 10^-pH = 10^-4.8 = 1.5E-5

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pH & pOH math

• pH definition: pH = –log [H+]• pOH definition: pOH = -log[OH-]• [H+]*[OH-] = 10^14 for aqueous solutions• If [H+]=10^-2, [OH-]=10^-12, product is 10^-14

– If one goes up, other must go down … like a see-saw

• pH + pOH=14 (exponents add when multiplying)– If pH=5, pOH=11 sum is 14, same see-saw effect

pH & pOH symmetry

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pH of acids

• pH is –log[H+] = -log[10-7] for water– As exponent gets smaller, [H+] goes UP– This is due to exponent being NEGATIVE– As acid gets stronger, pH gets smaller

• [H] = 1/1000 m/l = [10^-3], so pH = -log[10^-3] = 3.0• [H] = 1/10 mole/liter =[10^-1], so pH=-log[10^-1] =1.0

– Smaller pH = MORE acidic.

• pOH applies to hydroxide ion– As [OH-] increases, [H+] decreases– Sum of (pH + pOH) ≡ 14 by definition– Water [H+]*[OH-] = [10^-7]*[10^-7]=10^-14– Anything else = [10^3.7]*[10^10.3]– See-Saw between H+ and OH-

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Can use pH instruments

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pH to mole/liter example• This conversion either simple or complex

– Integer pH numbers (1, 2, 3 …) are simple• Mole/Liter exponent = pH value• [H+] = 10-2 mole/liter, pH=2

• Non Integer numbers require a calculator– pH meter photo shows orange juice– pH =3.44 … what H+ concentration?– [H+] = antilog of - 3.44 (exponentiation)

• 10^-pH = 10^-3.44• 10^-3.44 = 3.63*10-4 moles/Liter

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pH to mole/liter example

• pH meter photo shows orange juice– If pH =3.44, what is H+ concentration?– [H+] = antilog of 3.44 = 10^-pH = 10^-3.44– 10^-3.44 = 3.63E-4 moles/Liter

• Calculators vary, for a TI-30stat (not TI-30x)– Enter 3.44– Change sign (+/-) to -3.44– (2nd funtion, or INV) “10X” (often over the “log” key)– “=“ 3.63E-4, or 0.000363 in moles/liter

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Examples, mole/liter [H+] to pH

• Getting pH from concentration is easiest– Enter concentration, H+ ion is 0.000363 moles/liter– Use “Log” key, result is - 3.44– Definition of pH is – log [moles/liter]– So pH is – (--3.44) or pH = 3.44 (acidic side)

• What if H+ is 0.0000000363– Enter 0.0000000363 (7 zeros)– Use “Log” key, result is -7.44 (about neutral)

• H+ must be VERY low to have basic pH– Try 0.000000003 (8 zeros)– Use “Log” key, result is -8.44 (somewhat basic)

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pH to mole/liter example, using multiple calculators

• For Casio fx-280– Enter 3.44– Change sign (+/-) to -3.44– (2nd function, or shift) , then “10X” (over the “log” key)– press “=“ to get 3.63E-4, or 0.000363 in moles/liter

• For TI 30x (similar to bookstore model)– Press 2nd function (blue key)– press 10x (above log key)

• Will get 10^( • Use lower right key (-) to make value negative• Enter value of pH (e.g. 3.44)• Close parenthesis by adding “)” on the right• Push “enter” or “=“ for the answer … 3.63x10^-4

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Stomach cross-section

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Stomach has ≈ 2 moles/liter HCL

Gastric HCl• Gastric acid is a digestive fluid, formed in the stomach. It has a pH of 1.5 to

3.5 and is composed of hydrochloric acid (HCl) (around 0.5%, or 5000 parts per million), and large quantities of KCl and NaCl.

• The acid plays a key role in digestion of proteins, by activating 

digestive enzymes, and making ingested proteins unravel so that digestive enzymes can break down the long chains of amino acids.

• Gastric acid is produced by cells lining the stomach, which are coupled to systems to increase acid production when needed. Other cells in the stomach produce bicarbonate, to buffer the fluid, ensuring that it does not become too acidic. These cells also produce mucus, which forms a viscous physical barrier to prevent gastric acid from damaging the stomach.

• Cells in the beginning of the small intestine, or duodenum, further produce large amounts of bicarbonate to completely neutralize any gastric acid that passes further down into the digestive tract.

• The presence of gastric acid in the stomach and its function in digestion was first characterized by U.S. Army surgeon William Beaumont around 1830. Beaumont studied the stomach action of fur trapper Alexis St. Martin.

Stomach acid fairly strongpH of 2 similar to lemon juice acidity

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HCl in the stomachwe don’t drink HCl … so how does it get there?

2 minute U-Tube Videoproduction of HCL in stomach by Parietal Cells

http://wn.com/parietal_cell#

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http://wn.com/parietal_cell#

http://wn.com/parietal_cell#

http://wn.com/parietal_cell#

Basic pH idea is concentration

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More pH examples• [H+] pH Example• Acids • 1 mole/liter 0 1 molar HCl• 1 x 10-1 1 Stomach acid, HCl• 1 x 10-2 2 Lemon juice, Citric Acid• 1 x 10-3 3 Vinegar, Acetic Acid• 1 x 10-4 4 Soda Water, Carbonic Acid• 1 x 10-5 5 Rainwater• 1 x 10-6 6 Milk

Neutral • 1x 10-7 7 Pure water

• Bases • 1 x 10-8 8 Egg whites• 1 x 10-9 9 Baking soda, sodium bicarbonate• 1 x 10-10 10 Tums® antacid• 1 x 10-11 11 Ammonia, ammonium hydroxide• 1 x 10-12 12 Mineral lime - Ca(OH)2• 1 x 10-13 13 Drano®, NaOH + aluminum metal• 1 x 10-14 14 NaOH

Soil

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Lemon Juice … pH of 2A “tribasic” acid, 3 H+ ions per molecule

pH and the Human Body

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Acid Rain Formationnon-polar sulfur soluble in non-polar oil (like dissolves like)

Easily gets into fuel supplies, burns to form SO2

Major problem in 1800’s London due to sulfur in coal

• S + O2 SO2 from burning coal, liquid fuels

• sulfur dioxide is further oxidized in air

SO2 + 1/2 O2 SO3

• In the presence of water sulfur trioxide (SO3) is

converted rapidly to sulfuric acid:SO3(g) + H2O(aq) → H2SO4(aq)

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Acid rain cycle

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Acid rain and Marble (limestone)

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Acid Rain and the Taj MahalDamage from acid rain mars some of the world’s finest cultural monuments. Emissions reductions,

however, have helped slow the rate of damage in North America and Europe. India’s Taj Mahal

hasnot not fared as well. Scientists blame pollution local foundries and a nearby oil refinery. 

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Air pollution sources

Acid rain across USA

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pH of rain and fish habitat

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Aquatic life, sensitivity to pH

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How to measure pH?• Indicators

– Swimming pool kits for pH and chlorine/bromine– Color of dye can depend on acidity, chlorine, etc.– Natural dyes include berries, rhubarb, flowers. – Synthetic dyes designed for specific pH – Can combine dyes for rainbow of colors

• Instrumentation– Electric current through glass depends on H+

– Thin glass membrane used as sensor– Voltage translates to pH

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Natural pH indicator dyes in plantsA nice science fair project last year

Beets Geranium petals TeaBlueberries Petunia petals Curry powder Carrots Pansy petals ThymeCherries Poppy petals TumericGrapes Violet petals  OnionPurple peonies Strawberries Rose petals  Rhubarb Tulip petals  Red Cabbage

Retrieved from http://en.wikipedia.org/wiki/PH_indicator

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Litmus PaperLitmus is a mixture of 10 to 15 different dyes extracted from lichens. It is

absorbed onto filter paper which becomes a pH indicator to test for acidity. Blue litmus paper turns red under acidic conditions and red litmus paper turns blue under basic conditions. Neutral litmus paper is purple in

color.

Parmelia sulcata

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Litmus paper color changes

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Synthetic indicator dyesengineered for specific pH “switch”

“Universal Indicator” dye combinations give color rainbow

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Laboratory Determination of Acidity(a) Universal indicator colors in solutions of known pH from 1 to 12. (b) Testing pH with a paper strip. Comparing the color of the strip with

the code on the package gives the approximate pH.

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Starch+Iodineblue/black colorIodine can be used to tell if “Red Delicious” apples are ripeImmature “green” apples have excess starch blue stain “overripe” apples more sugar, little starchno iodine stain

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Color indicators not too precise

Today’s Experiment

• Making our own indicator– Red Cabbage “soup”

• Testing household products– Cleaners, degreasers,etc.

Kitchen Chemistry, pH indicators

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Now to the experiment

• Making an indicator – Our “cabbage soup” used as pH indicator

• Chop up sample of red cabbage• Boiling water used to extract dyes• Cool before using, use ice tray• Remove & discard solids, retain dye solution

– Testing household products in the lab• 5 drops indicator to test tubes + household chems• Note colors obtained, decide if acid or base

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Experiment details

• Using the indicator – Use white ceramic plate with 12 “pockets”

• Put 5-6 drops known pH solutions in 10 pockets• This your “reference” set of colors for each pH

– Test household products, 2nd ceramic plate• Try ALL of the unknowns• Put 5-6 drops unknown + cabbage indicator• Match colors to determine pH of “unknowns”• Label each with pH and note if acidic or basic

Cabbage Juice = pH Indicator

Cabbage Juice pH colors

Red Cabbage + Household Chem.

Cabbage colors in pH solutions

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Testing outside of the lab

• Testing at home … take home some indicator– Identify pH of soil, leaves, plants

• Add materials to water, let soak, measure• Dice large samples (e.g. leaves into small pieces)

– Identify pH of items at home (not highly colored ones)• Coffee, Tea, lemonade, milk, white wine, sodas• Cleaning products, ammonia, detergent, oxy-clean, etc.

– Holiday candidates• Cranberry sauce, apple juice• Soapy water• Berry pie residue (add soap to dish)

Now to the benches …

• Let’s do it!

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Hydronium Geometrytetrahedral electronic, pyramidal molecule

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Hydrogen ion solvation

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Examples, mole/liter to pH

• Getting pH from concentration is easy– Enter concentration, H+ ion is 0.000363 moles/liter– Use “Log” key, result is -3.44– Definition of pH is – log [moles/liter]– So pH is – (--3.44) or pH = 3.44 (acidic side)

• What if H+ is 0.0000000363– Enter 0.0000000363 (7 zeros)– Use “Log” key, result is -7.44 (about neutral)

• H+ must be VERY low to have basic pH– Try 0.000000003 (8 zeros)– Use “Log” key, result is -8.52 (somewhat basic)