acids, bases, and salts

Acids, Bases, and Salts

Acids and Baseshttp://www.unit5.org/chemistry/AcidBase.html1ACIDS and BASES

In nature, acids can be found in fruits: citric acid is responsible for the sharp taste of lemons. Vinegar contains acetic acid, and tannic acid from tree bark is used to tan leather. The stronger mineral acids have been prepared since the Middle Ages. One of these, aqua fortis (nitric acid), was used by assayers to separate gold from silver. Car batteries contain sulfuric acid, also strong and corrosive. A base is the opposite of an acid. Bases often feel slippery; bicarbonate of soda and soap are bases, and so is lye, a substance that can burn skin. Bases that dissolve in water are called alkalis. In water, acids produce hydroxide ions. When an acid and a base react together, the hydrogen and hydroxide ions combine and neutralize each other, forming water together and a salt. The strength of acids and bases can be measured on a pH scale.

Eyewitness Science Chemistry , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 42

Sulfate Ion Concentrations1985-2003198519861984

44Guiding QuestionsWhat is an acid?

What makes acids dangerous? Is acid rain an issue for us? What does pH balanced mean?

3What you should learn:Acids are essential for life, are commonly found in the home, and have a myriad of uses outside the home. The strength of an acid is determined by its electron structure Acids can cause chemical changes by breaking or weakening chemical bonds.

More Specifically...:Properties and Theory List general properties of acids and bases Compare Arrhenius and Brnsted-Lowry theories of acids Classify acids and bases as Arrhenius or Brnsted-Lowry Identify acid base conjugate pairs Convert between chemical formula and name using acid naming rules Relate the acidic and basic properties of aqueous solutions to the dissociation of water Distinguish between strong and weak acids in terms of a. degree of ionization/dissociation b. conductivity c. Ka d. pH Write and balance neutralization reaction an acid and base react to form a salt and water Sketch titration curves and be able to suggest a suitable indicator for a particular titration using a table of indicators Distinguish between 'equivalence point' and 'end point' List causes and dangers of acid rain

MathCalculate pH of acids and bases from hydrogen and hydroxide ion concentration (pH = -log[H3O+]) Convert between pH and pOH Calculate unknown concentration using titration data

Acids, Bases, and SaltsYou should be able to

Understand the acid-base theories of Arrhenius, Brnsted-Lowry,and Lewis.

Identify strong acids and bases and calculate their pHs.

Calculate the pH of a weak acid or base.

Calculate the concentration of a strong or weak acid or base from its pH.

Calculate the pH and ion concentration in a polyprotic acid.

Predict the pH of a salt from its formula and then calculate the pH of the salt.

Be familiar with titration curves and selection of an acid-base indicator.

5Fast Track to a 5 (page 61)pH scale

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14ACIDBASE NEUTRALEach step on pH scale represents a factor of 10.pH 5 vs. pH 6 (10X more acidic)pH 3 vs. pH 5 (100X different)pH 8 vs. pH 13 (100,000X different): measures acidity/basicity10x10x10x100x

Sren Sorensen(1868 - 1939)6pH is defined as the negative base-10 logarithm of the hydrogen ion concentration pH = log [H+] or [H+] = 10-pH

Hydrogen ion concentration in pure water is 1 x 10-7 M at 25C; the pH of pure water is log [1.0 x 10-7] = 7.00.

pH decreases with increasing [H+] adding an acid to pure water increases the hydrogen ion concentration and decreases the hydroxide ion concentration.

Adding a base to pure water increases the hydroxide ion concentration and decreases the hydrogen ion concentrationpH increases with decreasing [H+].

pH scale runs from pH = 0 (corresponding to 1 M H+) to pH 14 (corresponding to 1 M OH).

Relationships between acidity, basicity, and pH: If pH = 7.0, the solution is neutral. If pH < 7.0, the solution is acidic. If pH > 7.0, the solution is basic.

A change of 1,0 in the pH of a solution corresponds to a tenfold change in the hydrogen ion concentration because the pH scale is logarithmic.

65. 7Richter Scale - Earthquakes 3 4 1 27Area Radius Richter Scale DiameterpH[H+]100000056471128 mm10.1 M 100000 178.46 356.820.01 M 10000 56.45 112.830.001 M 1000 17.84 35.640.0001 M 100 5.643 11.2850.00001 M 10 1.7842 1.12860.000001 M 1 0.564 1 ------ 7 0.0000001 MCalibration CurveVinegarAmmonia 1 mL 3 mL 5 mL 10 mL 15 mLvinegarammonia

Using 3 mL vinegar titrate with 0.130 M NaOH solution.Calculate molarity (M) of acetic acid. M1V1 = M2V2M1 V1 = M2 V2(Macetic acid)(3.0 mL) = (0.130 MNaOH )(19.6 mL)Macetic acid = 0.8493 molarIt required 19.6 mL of NaOH to reach the endpoint.AcidBasevinegar ammonia248TEACHER NOTE: I mix 26 g of NaOH in 5 L of distilled water to yield a 0.130 M NaOH solution.When using household vinegar it requires ~20 mL of 0.130 M NaOH / 3.0 mL of vinegar.You may choose to add 52 g of base to reduce volume of base to reach endpoint.

I use phenolpthalein indicator its cheap. It is not the proper indicator for use with a weak acid (vinegar) and weak base (ammonia). A more accurate result would be achieved using the indicator bromthymol blue.radius = 112.8 cmradius = 35.7 cmradius = 11.3 cmradius = 3.6 cmradius = 1.1 cmradius = 0.1128 cmpH = 2pH = 3..pH = 1pH = 4 5 69Area Radius Richter Scale DiameterpH100000056471128 mm1 100000 178.46 356.82 10000 56.45 112.83 1000 17.84 35.64 100 5.643 11.285 10 1.7842 1.1286 1 0.564 1 ------ 7 Acids and Bases

pH < 7pH > 7taste sourtaste bitterreact w/basesreact w/acidsproton (H1+) donorproton (H1+) acceptorturn litmus redturn litmus bluelots of H1+/H3O1+lots of OH1react w/metalsdont react w/metals Both are electrolytes.10Acid vs. Base AcidpH > 7bitter tastedoes notreact withmetalspH < 7sour tastereact withmetalsAlikeDifferentRelated toH+ (proton)concentrationpH + pOH = 14Affects pHand litmus paperBaseDifferentTopicTopic11Acids A substance with at least one hydrogen atom that can dissociate to form an anion and an H+ ion (a proton) in aqueous solution, thereby forming an acidic solution

Bases Compounds that produce hydroxide ions (OH) and a cation when dissolved in water, thus forming a basic solution

Neutral Solutions that are neither basic nor acidic

PropertiesACIDSBASESelectrolytesturn litmus redsour tastereact with metals to form H2 gasslippery feelturn litmus bluebitter tasteChemASAPvinegar, milk, soda, apples, citrus fruitsammonia, lye, antacid, baking sodaelectrolytes12Copyright 2007 Pearson Benjamin Cummings. All rights reserved.AcidSour tasteTurns blue litmus redReacts with some metals to produce H2Dissolves carbonate salts, releasing CO2

BaseBitter tasteTurns red litmus blueSlippery to the touch13Properties of acids and bases

Acids Sour taste Turns blue litmus paper red Reacts with some metals to produce H2 Dissolves carbonate salts, releasing CO2

Bases Bitter taste Turns red litmus paper blue Slippery to the touch

Copyright 2007 Pearson Benjamin Cummings. All rights reserved.

14Properties of acids and bases

Acids Sour taste Turns blue litmus paper red Reacts with some metals to produce H2 Dissolves carbonate salts, releasing CO2

Bases Bitter taste Turns red litmus paper blue Slippery to the touch

Common Acids and BasesStrong Acids (strong electrolytes)HClhydrochloric acidHNO3nitric acidHClO4perchloric acidH2SO4sulfuric acidWeak Acids (weak electrolytes)CH3COOHacetic acidH2CO3carbonicStrong Bases (strong electrolytes)NaOH sodium hydroxideKOH potassium hydroxideCa(OH)2 calcium hydroxideWeak Base (weak electrolyte)NH3 ammoniaKotz, Purcell, Chemistry & Chemical Reactivity 1991, page 145Weak Base (weak electrolyte)

NH4OH ammoniaNH3 + H2O NH4OH15Phosphoric acid, H3PO4, is another acid commonly found in the laboratory. It is on the borderline between a strong and weak acid.Acid + Base Salt + WaterOrange juice + milk bad tasteEvergreen shrub + concrete dead bushUnder a pine tree + fertilizer white powder

HCl + NaOH NaCl + HOH salt water16Acid-Base Neutralization1+1-++Hydronium ionHydroxide ionH3O+OH-WaterH2OWaterH2OWaterH2OWaterH2ODorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 58417Acid-Base Neutralization1+1-++Hydronium ionHydroxide ionWaterH3O+OH-H2OWaterH2ODorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 58418Acid Rain19Acid Precipitation

Click HereClick Herehttp://nadp.sws.uiuc.edu/amaps2/20Click on the SO2 and NOx letters to view a map of the United States and the levels of these pollutants.

Acid rain is rainfall whose pH is less than 5.6 due to dissolved carbon dioxide, which reacts with water to give the weak acid carbonic acid.

Source of the increased acidity in rain due to the presence of large quantities of sulfate (SO42-) and nitrate (NO3-) ions, which come from nitrogen oxides and sulfur dioxide produced both by natural processes and by the combustion of fossil fuels

These oxides react with oxygen and water to give nitric acid and sulfuric acid.

Damage caused by acid rain 1. Dissolves marble and limestone surfaces due to a classic acid-base reaction 2. Accelerates the corrosion of metal objects 3. Decreases the pH of natural waters 4. Biological effects

Formation of Sulfuric Acid

Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 302SO2(g) + H2O(l) H2SO3(aq)2SO2(g) + O2(g) 2SO3(g)SO3(g) + H2O(l) H2SO4(aq)SO2(g) + H2O2(l) H2SO4(aq)

Catalyzed by atmospheric dustSulfuric acid+

+

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Figure courtesy of the National Atmospheric Deposition Program, Champaign, Ill.The progressively darker red areas on the map indicate the lowest pH levels and areas most prone to problems from acid rain. Hydrogen ion concentration as pH from measurementsmade at the field laboratories during 2003National Atmospheric Deposition Program/National Trends Networkhttp://nadp.sws.uiuc.edu

made at the Central Analytical Laboratory, 199922http://www.professionalroofing.net/images/1004_26_1.jpg

Acid Rain

Smoke stacks pollute SO2into the atmosphere. Thiscombines with water to formacid rain.Estimated sulfate ion deposition, 1999

23Marble is made of limestone. SO2 combines with water to form H2SO3 acid. The acid reacts with the CaO in the limestone to produce CO2 g and slowly dissolves.

SO4 Levelshttp://nadp.sws.uiuc.edu24

Nitrate Ion Concentrations1985-2003

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Coal Burning Power PlantCopyright 2007 Pearson Benjamin Cummings. All rights reserved.63Coal-fired units produce electricity by burning coal in a boiler to heat water to produce steam. The steam, at tremendous pressure, flows into a turbine, which spins a generator to produce electricity. The steam is cooled, condensed back into water, and returned to the boiler to start the process over.

http://www.ohiocitizen.org/campaigns/coal/smokestack_tall.jpg

Statewide Coal-Fired Power PlantsLegend

Existing Power Plant

Proposed Power PlantCity

64http://www.environmentillinois.org/uploads/Ps/Of/PsOfNh7eUYva5HjekgP3xQ/powerplantsmap1.gif

http://www.astecindustries.com/images/photos/Coal_Hands.jpgCoal Burning Power Plant

65http://images.google.com/imgres?imgurl=http://www.tva.gov/power/images/coalart.gif&imgrefurl=http://www.tva.gov/power/coalart.htm&h=312&w=504&sz=27&hl=en&start=2&tbnid=GbeRwGGykEt4jM:&tbnh=80&tbnw=130&prev=/images%3Fq%3Dcoal%2Bburning%2Bpower%2Bplant%26gbv%3D2%26hl%3Den%26sa%3DX

Copyright 2007 Pearson Benjamin Cummings. All rights reserved.66CO2 (g)H2O (l)H2CO3 (aq)

Carbon dioxideCarbonic acidWater Weak acidCopyright 2007 Pearson Benjamin Cummings. All rights reserved.67

Carbon dioxideWater

Carbonic Acid68

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Carbon dioxideWater70

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Carbonic AcidCopyright 2007 Pearson Benjamin Cummings. All rights reserved.79Common Acids80Common AcidsSulfuric AcidH2SO4

Nitric AcidHNO3

Phosphoric AcidH3PO4

Hydrochloric AcidHCl

Acetic Acid CH3COOH

Carbonic Acid H2CO3Battery acidUsed to make fertilizersand explosivesFood flavoringStomach acidVinegarCarbonated water

81Common AcidsFormulaName of AcidName of Negative Ion of Salt

HFhydrofluoricfluorideHBrhydrobromicbromideHIhydroiodiciodideHClhydrochloricchlorideHClOhypochloroushypochloriteHClO2chlorouschloriteHClO3chloricchlorateHClO4perchloricperchlorateH2ShydrosulfuricsulfideH2SO3sulfuroussulfiteH2SO4sulfuricsulfateHNO2nitrousnitriteHNO3nitricnitrateH2CO3carboniccarbonateH3PO3phosphorousphosphiteH3PO4phosphoricphosphate

82 Formation of Hydronium Ions1+hydronium ionH3O++hydrogen ion H+ waterH2O1+(a proton) 1+83

Sulfuric Acid, H2SO4Sulfuric acid is the most commonly produced industrial chemical in the world.

Uses: petroleum refining, metallurgy, manufacture of fertilizer, many industrial processes: metals, paper, paint, dyes, detergentsSulfuric acid is used in automobile batteries.H2SO4

oil of vitriol

84Compound produced in the largest quantity in the industrial world and is one of the oldest chemical compounds known

Production starts with elemental sulfur obtained through a technique called the Frasch process, in which very hot water forces liquid sulfur out of the ground in nearly pure form.

Sulfuric acid is produced by the reaction of sulfur dioxide with oxygen in the presence of vanadium(V) oxide (the contact process), followed by the absorption of the sulfur trioxide in concentrated sulfuric acid to produce oleum.

Uses to make fertilizers and potash, one of the major ingredients in fertilizers

http://en.wikipedia.org/wiki/Image:Sulfuric-acid-3D-vdW.png

Nitric Acid, HNO3Nitric acid stains proteins yellow (like your skin).

Uses: make explosives, fertilizers, rubber, plastics, dyes, and pharmaceuticals.

HNO3aqua fortis

OOONH85For making fertilizers. About 75% of the nitric acid produced in the United States is used in the manufacture of fertilizers. Ammonium nitrate is the most important nitrate so used, and is readily manufactured in plants using the combined Haber-Ostwald processes. Sodium nitrate and potassium nitrate are also used as fertilizer ingredients.

For making explosives. Many modern explosives are made directly or indirectly from nitric acid. The acid itself is not an explosive, but many of the compounds derived from it form the most violent explosives known. Among these are nitroglycerine, smokeless powder, and TNT.

For making dyes. Nitric acid reacts with several products obtained from coal tar, forming nitro compounds. One of these coal tar products, benzene, reacts with nitric acid to form nitrobenzene, C6H5NO2. Aniline, C6H5NH2, a compound used in making different dyes, is made by reducing nitrobenzene with hydrogen.

For making plastics. Cotton which consists mainly of cellulose, (C6H10)5)n, is treated with a mixture of nitric acid and sulfuric acid to make nitrocellulose plastics. A variety of products is formed, depending on the amount of nitric acid used, the temperature, and the length of time the acid is allowed to act on the cellulose. Manufacturers use sulfuric acid to absorb the water that is formed in the reaction. Celluloid, pyroxylins, photographic film, and many other products are made from such nitrocellulose plastics.

Modern Chemistry Metcalfe, Williams, and Catska (1966) pg. 491

http://en.wikipedia.org/wiki/Image:Nitric-acid-3D-balls-B.png

Hydrochloric Acid, HClThe stomach produces HCl to aid in the digestion of food.

Uses: For pickling iron and steel.Pickling is the immersion of metals in acid solution to removesurface impurities.

A dilute solution of HCl is called muriatic acid (available in many hardwarestores). Muriatic acid is commonly used to adjust pH in swimming pools and in the cleaning of masonry.HCl(g) + H2O(l) HCl(aq)hydrogen chloride water hydrochloric acid

86http://en.wikipedia.org/wiki/Image:HCl_molecule_model-VdW_surface.svgCommon Bases87Common BasesSodium hydroxideNaOHlye or caustic soda

Potassium hydroxideKOHlye or caustic potash

Magnesium hydroxideMg(OH)2milk of magnesia

Calcium hydroxideCa(OH) 2 slaked lime

Ammonia waterNH3 H2Ohousehold ammoniaName Formula Common Name.NH4OHNH41+ + OH1-ammonium hydroxide

hydroxideion OH1-

88Bases

Bases are ionic compounds that contain the hydroxide ion and a metal cation and have the general formula M(OH)n . When a base reacts with an acid, it accepts a proton (H+) and is therefore called a proton acceptor. Aqueous ammonia solution is also a common base. Replacing a hydrogen atom of NH3 with an alkyl group results in an amine (RNH2), which is also a base. Amines have pungent odors.

Common BasesSodium hydroxideNaOHlye or caustic soda

Potassium hydroxideKOHlye or caustic potash

Magnesium hydroxideMg(OH)2milk of magnesia

Calcium hydroxideCa(OH) 2 slaked lime

Ammonia waterNH3 H2Ohousehold ammoniaName Formula Common Name.NH4OHNH41+ + OH1-ammonium hydroxide

hydroxideion OH1-89Bases

Bases are ionic compounds that contain the hydroxide ion and a metal cation and have the general formula M(OH)n . When a base reacts with an acid, it accepts a proton (H+) and is therefore called a proton acceptor. Aqueous ammonia solution is also a common base. Replacing a hydrogen atom of NH3 with an alkyl group results in an amine (RNH2), which is also a base. Amines have pungent odors.

Relative Strengths of Acids and BasesperchloricHClO4hydrogen chlorideHClnitricHNO3sulfuricH2SO4hydronium ionH3O+hydrogen sulfate ionHSO4-phosphoricH3PO4aceticHC2H3O2carbonicH2CO3hydrogen sulfideH2Sammonium ionNH4+hydrogen carbonate ionHCO3-waterH2OammoniaNH3hydrogenH2Decreasing Acid Strengthperchlorate ionClO4-chloride ionCl-nitrate ionNO3-hydrogen sulfate ionHSO4-waterH2Osulfate ionSO42-dihydrogen phosphate ionH2PO4-acetate ionC2H3O2-hydrogen carbonate ionHCO3-hydro sulfide ionHS-ammoniaNH3carbonate ionCO32-hydroxide ionOH-amide ionNH2-hydride ionH-Decreasing Base StrengthAcid Formula Conjugate baseFormulaMetcalfe, Williams, Catska, Modern Chemistry 1966, page 229acid conjugate base + H+90Binary Hydrogen CompoundsOxysalts + H2O Oxyacids91Binary Hydrogen Compoundsof Nonmetals When Dissolved in Water(These compounds are commonly called acids.)The prefix hydro- is used to represent hydrogen, followed by the nameof the nonmetal with its ending replaced by the suffix ic and the wordacid added.Examples:

*HCl

HBr*The name of this compound would be hydrogen chloride if it was NOT dissolved in water.Hydrochloric acidHydrobromic acid92Naming Simple Chemical CompoundsIonic (metal and nonmetal)Covalent (2 nonmetals)MetalFormsonly onepositive ionFormsmore thanone positiveionNonmetalUse the name of element

Use elementname followedby a Romannumeral toshow the chargeFirstnonmetalSecondnonmetalBeforeelement name use a prefixto matchsubscriptUse a prefixbeforeelement name and end with ideSingleNegative IonPolyatomic Ion

Use the nameof theelement, butend with ide

Use thename ofpolyatomicion (ate orIte)93Naming Ternary Compounds from OxyacidsThe following table lists the most common families of oxy acids.one moreoxygen atommostcommonone lessoxygentwo lessoxygenHClO4perchloric acidHClO3chloric acidHClO2chlorous acidHClOhypochlorous acidH2SO4sulfuric acidH2SO3sulfurous acidH3PO4phosphoric acidH3PO3phosphorous acidH3PO2hypophosphorous acidHNO3nitric acidHNO2nitrous acid(HNO)2hyponitrous acid94An acid with aname ending inA salt with aname ending in-ic-ous-ate-iteformsformsHill, Petrucci, General Chemistry An Integrated Approach 1999, page 6095Oxyacids Oxysalts If you replace hydrogen with a metal, you have formed an oxysalt.A salt is a compound consisting of a metal and a non-metal. If thesalt consists of a metal, a nonmetal, and oxygen it is called anoxysalt. NaClO4, sodium perchlorate, is an oxysalt.HClO4perchloric acidHClO3chloric acidHClO2chlorous acidHClOhypochlorous acidNaClO4sodium perchlorateNaClO3sodium chlorateNaClO2sodium chloriteNaClOsodium hypochloriteOXYACID OXYSALT96ACID SALT per stem ic changes to per stem ate stem ic changes to stem ate stem ous changes to stem itehyper stem ous changes to hypo stem iteHClO3 + Na1+ NaClO3 + H1+ acid cation salt97Definitions98Arrhenius Acids and BasesAcids release hydrogen ions in water.Bases release hydroxide ions in water.

An acid is a substance that produces hydronium ions, H3O+, when dissolved in water.

Lewis DefinitionsA Lewis acid is a substance than can accept (and share) an electron pair.A Lewis base is a substance than can donate (and share) an electron pair.

Lewis Acid

Brnsted-Lowry DefinitionsA Brnsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.A Brnsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.Brnsted-Lowry

ArrheniusacidsAcid Definitions99SVANTE AUGUST ARRHENIUS

Svante Arrhenius was born in Sweden. He learned to read at the age of three and became interested in mathematics and physics at an early age. He proposed in his doctoral thesis that electrolytes split into ions in water. For his efforts he was awarded the barest of passes. Fortunately, William Ostwald and Jacobus vant Hoff promoted his work on electrolytic theory. He was awarded the 1903 Nobel prize for Chemistry for roughly the same thesis that had been nearly rejected nineteen years previously. He had universal interests in science and proposed the greenhouse effect.

Although this paper contains the standard chronological biography of Dr. Arrhenius, our goal is broader. First, we hope to show a more personal view of Dr. Arrhenius as related to us by Dr. Hubert Alyea in an interview. Second, we would like to present an overview of Dr. Arrhenius' Nobel Prize winning work and the difficulty he had in gaining acceptance in the scientific community. Finally we will offer a lab that simulates Dr. Arrhenius's work in hopes that some teachers will let their students experience a little piece of chemical history. Reminiscing about Svante A. ArrheniusDr. Hubert Alyea worked in Sweden under Dr. S. A. Arrhenius during 1925 and 1926. He was Arrhenius' last graduate student and has very fond memories of the great scientist. The path that led Dr. Alyea to Arrhenius' lab began in 1920 when he entered Princeton at the age of 15. His work and studies were delayed when he contracted polio at the age of 19; however, he feels that the year he spent in bed as a result of his illness was a time of great inner reflection and that he emerged with a strong commitment to accomplish something with his life that would contribute to the good of humanity. After Dr. Alyea graduated from Princeton he was awarded a grant to study with Arrhenius, and he left for Sweden. At the time of Dr. Alyea's arrival, Svante Arrhenius was already 66 years old, and the bulk of his research had been completed. What had been a bustling lab in earlier years now supported only a few graduate students. Dr. Alyea felt this atmosphere was perfect for him; he received the mentoring that he needed to thrive. Although Dr. Alyea worked directly under Beckstrum, Arrhenius' assistant, he recalls how Svante Arrhenius came in to the lab at least twice every day to ask how the work was going and what Alyea planned to do that day. Dr. Alyea was working on the idea of a free atom, a concept not readily accepted at that time. Svante Arrhenius seemed very interested in the idea, possibly because he may have thought of it as an extension of his own work on ions. Often Arrhenius would offer suggestions, but he did not insist on a specific research plan. Dr Alyea characterized Arrhenius as a man who "spread joy in the lab". Dr. Arrhenius related the following story of his youth to Hubert Alyea one day when he visited the lab. He said that he was working in a lab at the time, and one of the chemists gave him a vial of liquid to dispose of. The vial contained mercaptans, an extremely smelly substance, but at the time the young Svante Arrhenius was not aware of this. Having forgotten to dispose of the liquid properly while still at the lab, Arrhenius tossed the vial beside the road as he was riding his bicycle home. The cap of the vial came loose, and the mercaptans diffused rapidly causing a wide area of town to smell. A committee was appointed to investigate the problem, and after several weeks they concluded that unusual meteorological conditions had caused the mercaptans to form. They also concluded it was highly unlikely that this would ever happen again. Arrhenius was very sure that it would never happen again! Although there were only a few graduate students working under Dr. Arrhenius at this time, there was a constant flow of visitors from around the world, and Dr. Arrhenius seemed to enjoy the attention as well as their company. Often large groups were invited to dinner. One memorable dinner party was interrupted by Arrhenius when the aurora borealis began that evening. Arrhenius required everyone to accompany him outside as he explained to them the cause of this beautiful phenomenon. By the time he allowed his guests to return to dinner, all of the food was cold. Dr. Alyea said that the aurora borealis was not the only thing about the sky that held Arrhenius' attention. He believed that life on earth was brought here by spores carried through space from other planets by radiation. He believed that likewise these spores could have carried life to many planets resulting in life throughout the universe. Dr. Svante August Arrhenius won the Nobel Prize for his work in 1903, and Dr. Alyea's favorite story about Dr. Arrhenius as director of the Nobel Institute was about the 1926 Nobel Awards Dinner. Dr. Arrhenius took the podium to honor the guests, and as he said a few words about each of the famous men present he offered a toast with the alcoholic beverage of that man's choice. (Because this was during prohibition in the United States, Dr. Arrhenius toasted the American with water!) By the end of this series of toasts, Dr. Arrhenius appeared to be the only one of the distinguished guests still sober. In closing this inspirational interview, Dr. Hubert Alyea offered the following classification of teachers. A good teacher is one who explains a concept; a better teacher is one who asks questions about the concept; and the best teacher is one who demonstrates the concept then solicits the questions from the students. Dr. Alyea has been a wonderful role model for all teachers, and he has certainly been among the best of teachers. Biographical Information Arrhenius was born in Wijk, Sweden on February 19, 1859. He was an infant prodigy, teaching himself to read at three and graduating from high school as the youngest and brightest in his class. While attending the University of Uppsala studying under E. Edluld he worked on the properties of electricity passing through solutions. As a doctoral candidate he knew of Faraday's "ions", and it seemed to him that the puzzles involved in his conductivity studies could best be explained if one assumed that molecules could break up into electrically charged fragments called "ions". For his PhD thesis in 1884 he presented his "ionic theory", but it turned out to be a bit too revolutionary for his examiners' taste. He just barely passed with a fourth class rank, "not without merit". (See the next section of this paper, Arrhenius' Dissertation, for a more complete discussion of the Thesis, its rejection, and final acceptance.). Arrhenius obtained a travel grant and worked with Ostwald and Van't Hoff during which time his reputation increased as he clarified the ionic theory to his fellow chemists. In the late 1890's when electrically charged subatomic particles were discovered, Arrhenius' ionic theory suddenly made sense, and in 1903 he received the Nobel Prize in chemistry for it. In an extension of his ionic theory Arrhenius proposed definitions for acids and bases. He believed that acids were substances which produce hydrogen ions in solution and that bases were substances which produce hydroxide ions in solution. It is interesting that neither Bronsted nor Lewis received the Nobel Prize for continuing the work on the theory of acids and bases and for expanding the definition of these substances. It is noted that Arrhenius never did accept the Bronsted or Lewis definitions of acids and bases. Arrhenius studied reaction rates as a function of temperature, and in 1889 he introduced the concept of activation energy as the critical energy that chemicals need to react. He also pointed out the existence of a "greenhouse effect" in which small changes in the concentration of carbon dioxide in the atmosphere could considerably alter the average temperature of a planet. In 1908 Arrhenius advanced a new concept that had less success but has remained interesting to scientists and science fiction readers alike. He suggested that life might have orginated through spores blown up through the atmosphere of life-bearing worlds and then driven by radiation pressure across the gulfs of space to fall on hospitable but as yet sterile worlds. He died in Stockholm on October 2, 1927.

Arrhenius' DissertationHindsight shows that this young chemist had both great data and a revolutionary explanation, but neither the logic nor the data could change the mind set of the established chemists. Yet in the end his persistence and his model prevailed. He has been recognized by the professional societies and the Nobel Prize committee. This is the story that needs to be told to high school students. At the age of 24, Arrhenius had determined the conductivity of many electrolytes and planned his dissertation proposal. His data may have taken the following format: Conductivity (ohm-1cm2mol-1)Electrolyte0.001M0.0050.010.050.10.5Acetic acid4120146.54.62.0Hydrochloric acid377373370360351327Sodium acetate757270646149

From this data he noted: The resistance of an electrolyte is increased when the dilution is doubled. In very dilute solutions the conductivity is nearly proportional to the concentration. The conductivity of a solution is equal to the sum of conductivities of the salt and the solvent. If these laws are not observed, it must be due to a chemical reaction between the substances including the solvent. The electrical resistance rises with increasing viscosity, complexity of the ion, and the molecular mass of the solvent. (incorrect)

Arrhenius concluded from the above statements that the "molecule" breaks apart into a positive fragment and negative fragment, called ions, by its interaction with the solvent. This was a great leap in thinking. The concept of dissociation did not come at first but developed as time allowed Arrhenius to talk with other chemists. The review committee at University of Uppsala was very reluctant to award the doctorate degree to Arrhenius but finally gave a fourth rank (barely passing) for the dissertation. The disgrace prevented Arrhenius from aspiring to any professorship. Arrhenius accepted a chemistry position at Technical High School of Stockholm. From this position he sent copies of his research to chemists in many of the laboratories of Europe. The older chemists formally and quietly rejected his thesis; they were firmly convinced that molecules could not break up and could not carry an electric charge. They, like the dissertation committee, could not imagine sodium, a metal that violently reacts with water, and chlorine, a gas with toxic properties, existing as independent tragments after sodium chloride dissolved in water. By 1887 Arrhenius had worked out the language of his model with statements like "In all probability all electrolytes are completely dissociated at extreme dilutions". He could explain weak and strong acids by the concentration of the ions, known as percent dissociation today. Fortunately a few of the younger men in the new field of physical chemistry were intrigued by the notion of ions and by Arrhenius' reasoning. His model gained acceptance as the network of young chemists started to explain their results in terms of ions and dissociation. Van't Hoff is one example; he explained the higher osmotic pressure than predicted for electrolytes. Data continued to support the concept that ionic and polar covalently bonded substances dissociate in water. Some substances dissociate to a greater extent. These statements explain the colligitave properties and differences in pH of similar acid concentrations. Bronsted, Lowry, Lewis and others have developed a more general model of acids and bases for non-water solvent systems. We also must be prepared to accept the next model for the dissociation of substances. Biographical ResourcesAsimov, Issac, Asimov of Chemistry, Anchor Books, Garden City, NY, p103 - This source has a fine portrait and biography. Snelders H. A. M., "Arrhenius, Svante August", Charles C. Gillispie, editor, Dictionary of Scientific Biography, Volume 1, Charles Scribner's Sons, New York, New York, pp296 - 302 - Great seven page chronology of Arrhenius and a description of findings in his dissertation on ionic theory. Jaffe, Bernard, Crucibles: The Story of Chemistry, Simon and Schuster, Inc., New York,1930, pp 219 - 241. - This is the best 22 page resource for the human side to Arrhenius' personality and for the personal conflicts of accepting a revolutionary idea. Kendall, James, Great Discoveries by Young Chemists, Thomas Crowell Co, New York, 1935, pp 124 -138.- Chapter V is the chemistry of solutions which shows the supportive relationchip of Van't Hoff, Ostwald and others. Shakhashiri, Bassam, Demonstrations in Chemistry, Volume 3,University of Wisconsin, 1989, pp 3 - 26. - Within the framework of acid/base concept development, Arrhenius' data on conductivity is simulated with more modern measurements. The table of data, page 8, lists the conductivities of 12 compounds at dilutions of 0.05 M to O.001 M. Demonstration 8.21 shows how to build a light bulb conductivity tester with two sets of electrodes for serial comparison.

http://www.woodrow.org/teachers/chemistry/institutes/1992/Arrhenius.html

GILBERT NEWTON LEWIS:AMERICAN CHEMIST (1875-1946)At the end of a manuscript in the Bancroft Library at UC Berkeley, we find the words: "I have attempted to give you a glimpse...of what there may be of soul in chemistry. But it may have been in vain. Perchance the chemist is already damned and the guardian the blackest. But if the chemist has lost his soul, he will not have lost his courage and as he descends into the inferno, sees the rows of glowing furnaces and sniffs the homey fumes of brimstone, he will call out-: 'Asmodeus, hand me a test-tube.'"(1)These are fitting words from a man who was the most eminent figure in a great revolution that brought America to the forefront in chemistry. Gilbert Newton Lewis was probably the greatest and most influential of American chemists. Through the nineteenth century, Europe dominated science, but the first half of the twentieth century brought a tidal wave of scientific research that thrust America to the forefront. Lewis influenced this revolution by both his teaching and his research. During his career he published over 150 papers. Gilbert Newton Lewis was born at Weymouth, Massachusetts, on October 23, 1875. He was educated at home by his parents in the style of the English tutoring system. His only public schooling occurred between the ages of 9 to 14 years in Lincoln, Nebraska. At age fourteen, Lewis entered the University of Nebraska but transferred to Harvard College after three years. In 1899 he was awarded his PhD at age 24 under the supervision of T.W. Richards. Richards trained Lewis in experimental techniques and careful measurements and fostered his interest in thermodynamics. Richards' idea that "Believing in Faraday's methods ... fact is more important than theory..." (2) influenced Lewis' approach to research throughout his career. Conflicts with Richards over bonding in atomic and molecular structures caused Lewis to leave Harvard. This ended a two-year period in which he published nothing, the only non-productive time in his career. Lewis spent one year in the Phillipines as the Superintendent of the Bureau of Weights and Measures before joining the faculty at MIT where he found a group of young, talented physical chemists interested in doing research. This group was brought together by A.A. Noyes who, like Richards, had received his doctorate under Ostwald at the University of Leipzig. This research center provided an energizing atmosphere where Lewis spent seven productive years during which he undertook the systematic determination of the electrode potentials of the elements. Lewis left MIT when he was appointed the Chairman of the Department of Chemistry and the Dean of the College of Chemistry at UC Berkeley in 1912, positions he held until his sudden death in his laboratory on March 23, 1946, a remarkable 34-year tenure. Lewis believed that a chemistry department should simultaneously teach science and advance it, always remembering that the most important emphasis must be placed on fundamental principles rather than its technical applications. To achieve his goals, Lewis nurtured the Chemistry Department at Berkeley with his own vision. Dr. Lewis focused on young brilliant minds in chemistry, choosing to work with the exceptional rather than the average students. Course content was aimed at giving superior understanding of chemistry fundamentals rather that lots of facts. Students were encouraged to think for themselves by free discussions between students and teacher. Following his early training, Lewis had graduate students teach lower division courses. The present use of problem sets which challenge even the best young minds in organic chemistry can be traced back to Lewis' penchant for innovative study sheets. Upper division honor students were required to do research. Lewis strove to give graduate students freedom in selecting professors for research advisors with the option to change when desired. Students were given the run of the lab facilities and storerooms. Everyone cooperated with each other and no one was too busy to help with another's research. This open policy built initiative and morale among all participants in UC Berkeley's Chemistry Department. There were weekly research seminars at which Lewis sought to educate the graduate students, the staff and himself. At these meetings, research papers were presented by staff, students and eminent visitors on current events, project proposals and journal articles. Throughout his career, Lewis remained open-minded to criticism and was well known for his wit and insight. At UC Berkeley Lewis built a remarkably strong department which trained a large number of chemists including Nobel Prize Laureates, members of the National Academy of Science, and numerous individuals who became department and division chairs. Besides his contributions to education, his research in four areas of science give Lewis his remarkable status. Most of his research is focused on thermodynamics and its relationship to chemical equilibrium, the electron-pair bonding theory of atoms and molecules, isotopes (mainly deuterium), and the interaction of light with matter. During his tenure at MIT, Lewis researched standard electrode potentials, conductivity, free energy and other thermodynamic constants for the elements. These tables are still being used. His ability to organize and apply the scattered laws of thermodynamics brought about the evolution of physical chemistry into the science as it is known today. Lewis once defined physical chemistry as encompassing "everything that is interesting".(3) In his theory of the shared electron pair, Lewis did not believe only that an electron completely transfers from one atom to another, as in the positive-negative theory. He describes the partial transfer of two electrons, one from each of the two bonding atoms, so that there is a shared pair of electrons between them. This eliminates the need for the formation of oppositely charged atoms when there was no indication of individually charged atoms (ions) in a compound. This was the first description of covalent bonding. Lewis theorized that electrons in an atom pair up around the nucleus, usually forming a tetrahedral arrangement. Although he never actually used the term "octet" for four pairs of electrons, the octet rule is often associated with Lewis. His main concern was with individual bonds between atoms rather than with all the electron pairs around each nucleus. Lewis' book, Valence and the Structure of Atoms and Molecules, is a classic, one of the greatest contributions to modern bonding theory. Lewis' research on isotopes is an example of his wide-ranging and prolific interests. He published twenty-six papers on heavy hydrogen and heavy water, isotopes of lithium, and neutron physics. He predicted the existence of naturally occuring heavy water before he isolated it. Lewis continued with his research until his death on March 23, 1946, in his laboratory surrounded by his beakers and test tubes. There is no scientist in American history who has contributed more extensively to all fields in chemistry than Gilbert Newton Lewis. His thinking was far ahead of his time and his theories have had profound influence on modern chemistry. Although Lewis was never to receive the Nobel Prize, it is commonly felt that his work in thermodynamics and valence theory more than merited this award. GILBERT NEWTON LEWIS: AMERICAN CHEMIST (1875-1946)DATEEVENT1875Born on October 23 in Weymouth, Mass.1889Entered Univ. of Nebraska at age 14.1892Transferred to Harvard College.1899PhD at age 24.1900Instructor at Harvard College. Began research on electrochemistry and chemical equilibrium under Richards.1904Superintendent of Weights and Measures in the Phillipines.1905Faculty position at MIT. Began work on thermodynamics and free energies for elements.1912Married Mary Sheldon, daughter of Harvard Professor.1912Appointed Chairman of Dept. of Chemistry and Dean of the College of Chemistry, Berkeley.1918World War I, France. Appointed Chief of the Defense Division of Chemical Warfare Service. Received the Distinguished Service Medal (USA) and the Cross of the Legion of Honor (France).1923Authored Thermodynamics and the Free Energy of Chemical Substances with M. Randall. Authored Valence and the Structure of Atoms and Molecules. Clarified electron-pair bonding in covalent substances. Began work on a more inclusive acid-base theory.1930'sWorked on deuterium. First to prepare pure deuterium and its compounds. Published 26 papers on deuterium and other isotopes.1938Franklin Institute lectures on acids and bases. 1940's Worked on photochemical processes.1946Unexpected death in laboratory on March 23.

Annotated ReferencesJournal of Chem. Ed., "Gilbert Newton Lewis: Report of the Symposium", 1984, 61, p. 2-21, p. 93-108, p. 185-204. --The G.N. Lewis Symposium, presented at the Las Vegas ACS meeting in March of 1982, provides the best overall reference material about Lewis. Chairman Derek A. Davenport of Purdue University organized the symposium honoring Dr. Lewis while many of his friends and colleagues were still alive and able to give valuable insights into his life and far-reaching contributions to modern chemistry. The proceedings of the Lewis symposium were printed in the January 1984 Journal of Chemical Education and its subsequent two issues. Lewis, Gilbert Newton, Valence and the Structure of Atoms and Molecules, Dover Publications, Inc., New York, 1966. --Lewis' book Valence and the Structure of Atoms and Molecules is one of the few science classics that can be largely understood by teachers at the secondary level. Much of the information is currently being taught in chemical education although some of the material in his book is now known to be inaccurate. Ihde, Aaron J., The Development of Modern Chemistry, "Gilbert Newton Lewis", Dover Books, NY, 1984, p. 505, p. 534, p.536-41, p. 542. --This is a reference for general information about Lewis' research and publications. Servos, John W., Physical Chemistry from Ostwald to Pauling, Princeton University Press, Princeton, New Jersey, 1990. --This final reference contains materials for the individual who is more interested in specific details and facts on the work and life of G.N. Lewis and other American physical chemists. It is not as useful as the symposium articles. Bibliography1. Davenport, Derek A., "Gilbert Newton Lewis: 1875 - 1946", J. of Chem Ed., 1984, 61, p.2. 2. Servos, John W., "G.N.Lewis: The Disciplinary Setting", J. of Chem Ed., 1984, 61, p. 7. 3. Seaborg, Glenn T., "The Research Style of Gilbert N. Lewis", J. of Chem Ed., 1984, 61, p. 93.

http://www.woodrow.org/teachers/chemistry/institutes/1992/Lewis.htmlAcid DefinitionsLewis acids

Brnsted-Lowry

ArrheniusacidsThe Arrhenius model of acidsand bases was broadened bythe Brnsted-Lowry model.

The Lewis acid-base model isthe most general in scope.

The Lewis definition of an acidincludes any substance thatis an electron pair acceptor;a Lewis base is any substancethat can act as an electron pair donor.Ralph A. Burns, Fundamentals of Chemistry 1999, page 483100Arrhenius Acids and BasesAcids release hydrogen ions in water.Bases release hydroxide ions in water.


Brnsted-Lowry DefinitionsA Brnsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.A Brnsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.

Lewis DefinitionsA Lewis acid is a substance than can accept (and share) an electron pair.A Lewis base is a substance than can donate (andshare) an electron pair.

Acids and bases can be defined in different ways: 1. Arrhenius definition: An acid is a substance that dissociates in water to produce H+ ions (protons), and a base is a substance that dissociates in water to produce OH ions (hydroxide); an acid-base reaction involves the reaction of a proton with the hydroxide ion to form water

2. BrnstedLowry definition: An acid is any substance that can donate a proton, and a base is any substance that can accept a proton; acid-base reactions involve two conjugate acid-base pairs and the transfer of a proton from one substance (the acid) to another (the base)

3. Lewis definition: A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor

Lewis acids

Brnsted-Lowry

ArrheniusacidsThe Arrhenius model of acidsand bases was broadened bythe Brnsted-Lowry model.

The Lewis acid-base model isthe most general in scope.

The Lewis definition of an acidincludes any substance thatis an electron pair acceptor;a Lewis base is any substancethat can act as an electron pair donor.Ralph A. Burns, Fundamentals of Chemistry 1999, page 483Acid Definitions101Arrhenius Acids and BasesAcids release hydrogen ions in water.Bases release hydroxide ions in water.


Brnsted-Lowry DefinitionsA Brnsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.A Brnsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.

Lewis DefinitionsA Lewis acid is a substance than can accept (and share) an electron pair.A Lewis base is a substance than can donate (and share) an electron pair.

Acid Base SystemsTypeAcidBase ArrheniusH+ or H3O + producerOH - producer Brnsted-LowryProton (H +) donorProton (H +) acceptor LewisElectron-pair acceptorElectron-pair donor

102Svante Arrhenius (1859-1927), a Swedish scientist, won the 1903 Nobel Prize for his work on ionization. He introduced the idea of compounds dissociating, or splitting, into their constituent ions in solution. He explained how the strength of acids in aquos (water) solution depend on its concentration of hydrogen ions.

- Eyewitness Science Chemistry , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 43

Svante A. Arrhenius, 1859 - 1927Johannes N. Brnsted, 1879 - 1947Thomas M. Lowry, 1874 - 1936Gilbert N. Lewis, 1875 - 1946

Acids and bases can be defined in different ways: 1. Arrhenius definition: An acid is a substance that dissociates in water to produce H+ ions (protons), and a base is a substance that dissociates in water to produce OH ions (hydroxide); an acid-base reaction involves the reaction of a proton with the hydroxide ion to form water

2. BrnstedLowry definition: An acid is any substance that can donate a proton, and a base is any substance that can accept a proton; acid-base reactions involve two conjugate acid-base pairs and the transfer of a proton from one substance (the acid) to another (the base)

3. Lewis definition: A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor

Arrhenius Acid1+++hydronium ionH3O+1-chloride ionCl-waterH2Ohydrogen chlorideHCl(an Arrhenius acid)Any substance that releases H+ ions as the only positive ion in the aqueous solution.103Arrhenius definition of acids and bases Acids are substances that dissolve in water to produce H+ ions. Bases are substances that dissolve in water to produce hydroxide (OH) ions. Two limitations 1. Definition applied only to substances in aqueous solution 2. Definition predicted that only substances that dissolve in water to produce H+ and OH ions should exhibit the properties of acids and bases

DefinitionsArrhenius - In aqueous solutionHCl + H2O H3O+ + Cl Acids form hydronium ions (H3O+)

HHHHHHClClOO+acidCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

104DefinitionsArrhenius - In aqueous solutionBases form hydroxide ions (OH-)NH3 + H2O NH4+ + OH-

HHHHHHNNOO+HHHHbaseCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

105Arrhenius Bases and Their PropertiesAccording to the definition of Arrhenius a:Base - "a substance whose water solution yields... Are NaOH and NH3 considered to be Arrhenius bases?1) Bases are electrolytesDissociation equation for NH3NH3(g) + H2O(l) NH41+(aq) + OH1-(aq)Dissociation equation for NaOHNaOH(s) Na1+(aq) + OH1-(aq)2) Bases cause indicators to turn a characteristic color3) Bases neutralize acidsNaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)YES4) Water solutions of bases tasted bitter and feel slippery.hydroxide ions (OH-) as the only negative ions."106NeutralizationNeutralization is a chemical reaction between an acid and a base to produce a salt (an ionic compound) and water.NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)baseacidsaltwaterSome neutralization reactions:H2SO4(aq) + NaOH(aq)Na2SO4 +HOHsulfuric acidsodium hydroxidesodium sulfatewaterHC2H3O2(aq) + Ca(OH)2(aq)Ca(C2H3O2)2 +HOHacetic acidcalcium hydroxidecalcium acetatewater2222107NeutralizationACID + BASE SALT + WATERHCl + NaOH NaCl + H2OHC2H3O2 + NaOH NaC2H3O2 + H2OSalts can be neutral, acidic, or basic.Neutralization does not mean pH = 7.weakstrongstrongstrongneutralbasicCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

108ACID + BASE SALT + WATERHCl + NaOH NaCl + H2OHC2H3O2 + NaOH NaC2H3O2 + H2OSalts can be neutral, acidic, or basic.Neutralization does not mean pH = 7.weakstrongstrongstrongneutralbasicCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

109SaltsSalts - Ionic compounds containing a positive ion other than the hydrogen ion and a negative ion other than the hydroxide ion.i.e., a metal and a non-metalUnder what conditions do salts conduct current?NaCl(s) + H2O(l) Na1+(aq) + Cl1-(aq)Formulas and names of common saltsSALTFORMULACommon Namesodium chlorideNaCl(table) saltsodium nitrateNaNO3Chile saltpetersodium bicarbonateNaHCO3baking sodapotassium carbonateK2CO3potashammonium chlorideNH4Clsal ammoniac

NaCl110Salt FormationNaOHHClstrongbasestrongacidsalt of a strong base and a strong acidNaClNaOHHC2H3O2strongbaseweakacidsalt of a strong base and a weak acidNaC2H3O2Note: that in each case H-OH (water) is formedNaOH + HCl NaCl + H2ONaOH + HC2H3O2 NaC2H3O2 + H2O111Salt FormationNH3H2SO4weakbasestrongacidsalt of a weak base and a strong acid(NH4) 2SO4NH3HC2H3O2weakbaseweakacidsalt of a weak base and a weak acidNH4C2H3O2Note: that in each case H-OH (water) is also formedNH4OHH2SO4NH4OH + H2SO4 (NH4)2SO4 + H2ONH4OH + HC2H3O2 NH4C2H3O2 + H2ONH4OH112NH3H2SO4weakbasestrongacidsalt of a weak base and a strong acid(NH4) 2SO4NH4OHH2SO4ammonium ionNH4+hydroxide ionOH-1+1-NH4+OH-1+1-sulfuric acid(NH4)2SO4HOH1+HOH1+sulfate ion2 NH4OH + H2SO4 (NH4)2SO4 + 2 HOHwaterammonium sulfate2-H2SO42 NH4OH + H2SO4 (NH4)2SO4 + 2 H2O113phosphoric acidammonium hydroxideammonium phosphateReactions that produce salt acid + base salt +waterH3PO4NH4OH(NH4)3PO4H2Onitric acidmagnesium hydroxidemagnesium nitrateHNO3Mg(OH)2Mg(NO3)2H2Ocarbonic acidpotassium hydroxidepotassium carbonateH2CO3KOHK2CO3H2Oacetic acidaluminum hydroxidealuminum acetateHC2H3O2Al(OH)3Al(C2H3O2)3H2Operchloric acidbarium hydroxidebarium perchlorateHClO4Ba(OH)2Ba(ClO4)2H2O++andyieldsandwater114Brnsted-Lowry Acids and Bases1++hydronium ionH3O+1-chloride ionCl-(base)H2O(acid)HCld+d-Acid = any substance that donates a proton.Base = any substance that accepts a proton.115Brnsted Lowry definition of acids and bases

A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutions

BrnstedLowry acids and bases1. Base any species that can accept a proton2. Acid any substance that can donate a protonBrnsted-Lowry Acids and Bases1++hydronium ionH3O+1-chloride ionCl-(base)H2O(acid)HCld+d-Acid = any substance that donates a proton.Base = any substance that accepts a proton.116Brnsted Lowry definition of acids and bases

A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutionsBrnsted-Lowry Acids and Bases(acid)H2O(base)NH3d+d-1++ammonium ionNH4+1-hydroxide ionOH-117Brnsted Lowry definition of acids and bases

A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutionsBrnsted-Lowry Acids and Bases(acid)H2O(base)NH3d+d-1++ammonium ionNH4+1-hydroxide ionOH-118Brnsted Lowry definition of acids and bases

A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutionsBrnsted-Lowry Acids and Bases1++ammonium ionNH4+1-hydroxide ionOH-(acid)H2O(base)NH3d+d-119Brnsted Lowry definition of acids and bases

A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutionsBrnsted-Lowry Acids and Bases1++ammonium ionNH4+1-hydroxide ionOH-(acid)H2O(base)NH3d+d-120Brnsted Lowry definition of acids and bases

A more general definition of acids and bases An acid is any substance that can donate a proton. A base is any substance that can accept a proton. Not restricted to aqueous solutionsDefinitionsBrnsted-LowryHCl + H2O Cl + H3O+Acids are proton (H+) donors. Bases are proton (H+) acceptors.conjugate acidconjugate basebaseacidCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

121DefinitionsH2O + HNO3 H3O+ + NO3 CBCAABCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

HHOHOOONBaseAcid122DefinitionsAmphoteric - can be an acid or a base.NH3 + H2O NH4+ + OH-CACBBACourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

HHOHNBaseAcidHH123DefinitionsF -H2PO4-H2OHFH3PO4H3O+Give the conjugate base for each of the following:Polyprotic - an acid with more than one H+Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

124DefinitionsBr -HSO4-CO32-HBrH2SO4HCO3-Give the conjugate acid for each of the following:Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

125DefinitionsLewisAcids are electron pair acceptors. Bases are electron pair donors.Lewis base

Lewis acidCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

126pH Scale

Sren Sorensen(1868 - 1939)127The pH scale was invented by the Danish chemist Sren Sorensen for a brewery to measure the acidity of beer.pH Scale

AcidBase0714Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 515[H+] pH

10-14 14

10-13 13

10-12 12

10-11 11

10-10 10

10-9 9

10-8 8

10-7 7

10-6 6

10-5 5

10-4 4

10-3 3

10-2 2

10-1 1

100 0

1 M NaOHAmmonia(householdcleaner)BloodPure waterMilkVinegarLemon juiceStomach acid1 M HClAcidic Neutral Basic128S.P.L. Sorensen (1868-1939) introduced the pH scale to measure the concentration of hydrogen ions in solution. The more hydrogen ions, the stronger the acid. The amount of hydrogen ions in solution can affect the color of certain dyes found in nature. These dyes can be used as indicators to test for acids and alkalis. An indicator such as litmus (obtained from lichen) is red in acid. If base is slowly added, the litmus will turn blue when the acid has been neutralized, at about 6-7 on the pH scale. Other indicators will change color at different pHs. A combination of indicators is used to make a universal indicator.

- Eyewitness Science Chemistry , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 42

pH of Common SubstancesTimberlake, Chemistry 7th Edition, page 3351.0 MHCl0gastricjuice1.6vinegar2.8carbonated beverage3.0orange3.5apple juice3.8tomato4.2lemonjuice2.2coffee5.0bread5.5soil5.5potato5.8urine6.0milk6.4water (pure)7.0drinking water7.2blood7.4detergents8.0 - 9.0bile8.0seawater8.5milk of magnesia10.5ammonia11.0bleach12.01.0 MNaOH(lye)14.089101112141334562170acidicneutralbasic[H+] = [OH-]129pH of Common Substance 14 1 x 10-14 1 x 10-0 0 13 1 x 10-13 1 x 10-1 1 12 1 x 10-12 1 x 10-2 2 11 1 x 10-11 1 x 10-3 3 10 1 x 10-10 1 x 10-4 4 9 1 x 10-9 1 x 10-5 5 8 1 x 10-8 1 x 10-6 6

6 1 x 10-6 1 x 10-8 8 5 1 x 10-5 1 x 10-9 9 4 1 x 10-4 1 x 10-10 10 3 1 x 10-3 1 x 10-11 11 2 1 x 10-2 1 x 10-12 12 1 1 x 10-1 1 x 10-13 13 0 1 x 100 1 x 10-14 14NaOH, 0.1 MHousehold bleachHousehold ammonia

Lime waterMilk of magnesia

Borax

Baking sodaEgg white, seawaterHuman blood, tearsMilkSalivaRain

Black coffeeBananaTomatoesWineCola, vinegarLemon juice

Gastric juiceMore basicMore acidicpH [H1+] [OH1-] pOH 7 1 x 10-7 1 x 10-7 7130Acid Base ConcentrationspH = 3pH = 7pH = 11OH-H3O+OH-OH-H3O+H3O+[H3O+] = [OH-] [H3O+] > [OH-] [H3O+] < [OH-] acidicsolutionneutralsolutionbasicsolutionconcentration (moles/L)10-1410-710-1Timberlake, Chemistry 7th Edition, page 332131pH

pH = -log [H1+]Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 285132pH CalculationspHpOH[H3O+][OH-]pH + pOH = 14pH = -log[H3O+][H3O+] = 10-pHpOH = -log[OH-][OH-] = 10-pOH[H3O+] [OH-] = 1 x10-14133The pH scale is a concise way of describing the H3O+ concentration and the acidity or basicity of a solution

pH and H+ concentration are related as follows: pH = log10[H+] or [H+] = 10pH

pH of a neutral solution ([H3O+] = 1.00 x 107 M) is 7.00

pH of an acidic solution is < 7, corresponding to [H3O+] > 1.00 x 107

pH of a basic solution is > 7, corresponding to [H3O+] < 1.00 x 107

The pH scale is logarithmic, so a pH difference of 1 between two solutions corresponds to a difference of a factor of 10 in their hydronium ion concentrationsThere is an analogous pOH scale to describe the hydroxide ion concentration of a solution; pOH and [OH] are related as follows: pH = log10[OH] or [OH] = 10pOH

A neutral solution has [OH] = 1.00 x 107, so the pOH of a neutral solution is 7.00

The sum of the pH and the pOH for a neutral solution at 25C is 7.00 + 7.00 = 14.00 pKw = log Kw = log([H3O+] [OH]) = (log[H3O+]) + (log[OH]) = pH + pOH

At any temperature, pH + pOH = pKw, and at 25C, where Kw = 1.01 x 1014, pH + pOH = 14.00; pH of any neutral solution is just half the value of pKw at that temperature

pH = - log [H+]pH = 4.6pH = - log [H+]4.6 = - log [H+]- 4.6 = log [H+]- 4.6 = log [H+]Given:2nd log 10x antilogmultiply both sides by -1substitute pH value in equationtake antilog of both sidesdetermine the [hydronium ion]choose proper equation[H+] = 2.51x10-5 MYou can check your answer by working backwards.pH = - log [H+]pH = - log [2.51x10-5 M]pH = 4.6Recall, [H+] = [H3O+]134pH and pOH CalculationsKeys

pH and pOH Calculations

pH and pOH Calculations

http://www.unit5.org/chemistry/AcidBase.html135Acid DissociationmonoproticdiproticpolyproticHA(aq)H1+(aq) + A1-(aq) 0.03 M0.03 M0.03 MpH = - log [H+]pH = - log [0.03M]pH = 1.52e.g. HCl, HNO3H2A(aq)2 H1+(aq) + A2-(aq) 0.3 M0.6 M0.3 MpH = - log [H+]pH = - log [0.6M]pH = 0.22e.g. H2SO4Given: pH = 2.1find [H3PO4]assume 100% dissociatione.g. H3PO4H3PO4(aq)3 H1+(aq) + PO43-(aq) ? Mx MpH = ?136Acids differ in the number of hydrogen ions they can donate. Monoprotic acids are compounds capable of donating a single proton per molecule. Polyprotic acids can donate more than one hydrogen ion per molecule.

Polyprotic acids contain more than one ionizable proton, and the protons are lost in a stepwise manner.

The fully protonated species is always the strongest acid because it is easier to remove a proton from a neutral molecule than from a negatively charged ion; the fully deprotonated species is the strongest base. Acid strength decreases with the loss of subsequent protons, and the pKa increases. The strengths of the conjugate acids and bases are related by pKa + pKb = pKw, and equilibrium favors formation of the weaker acid-base pair.

When a strong base is added to a solution of a polyprotic acid, the neutralization reaction occurs in stages.

1. The most acidic group is titrated first, followed by the next most acidic, and so forth 2. If the pKa values are separated by at least three pKa units, then the overall titration curve shows well-resolved steps corresponding to the titration of each proton

Given: pH = 2.1find [H3PO4]assume 100% dissociationH3PO4(aq)3 H1+(aq) + PO43-(aq) X M0.00794 MStep 1) Write the dissociation of phosphoric acidStep 2) Calculate the [H+] concentrationpH = - log [H+]2.1 = - log [H+]- 2.1 = log [H+]2ndlog- 2.1 = log [H+]2ndlog[H+] = 10-pH[H+] = 10-2.1[H+] = 0.00794 M[H+] = 7.94 x10-3 M7.94 x10-3 MStep 3) Calculate [H3PO4] concentrationNote: coefficients (1:3) for (H3PO4 : H+)7.94 x10-3 M3=0.00265 M H3PO4 137How many grams of magnesium hydroxide are needed to add to 500 mL of H2O to yield a pH of 10.0?Step 1) Write out the dissociation of magnesium hydroxideMg2+OH1-Mg(OH)2Mg(OH)2(aq)Mg2+(aq)2 OH1-(aq)+Step 2) Calculate the pOHpH + pOH = 1410.0 + pOH = 14pOH = 4.0Step 3) Calculate the [OH1-]pOH = - log [OH1-][OH1-] = 10-OH[OH1-] = 1 x10-4 M1 x10-4 M0.5 x10-4 M5 x10-5 MStep 4) Solve for moles of Mg(OH)2

x = 2.5 x 10-5 mol Mg(OH)2Step 5) Solve for grams of Mg(OH)2x g Mg(OH)2 = 2.5 x 10-5 mol Mg(OH)21 mol Mg(OH)2= 0.00145 g Mg(OH)258 g Mg(OH)2138Practice Problems - KeyKeys

Practice Problems - Answer Key Practice Problems - Answer Key http://www.unit5.org/chemistry/AcidBase.html139LeChateliers Principle

Henry Le Chatelier (1850-1936)

an influential French chemist of the 19th century 140What you should learn:All systems strive for equilibrium, a state in which there is no net change. Reactions occur when molecules collide with enough energy and in the right orientation so the more likely a collision, the more likely a reaction will take place. More Specifically...:Reaction Rates Describe the steps involved in a chemical reaction: reaction mechanism Explain the concept of activation energy and activated complex State that catalysts are used to lower the activation energy of a reaction Use Collision Theory to describe how temperature, concentration, surface area, degree of randomness and catalysts affect the rate of reaction Interpret tabular and graphical data relating to rates

Equilibrium Classify reactions as reversible or irreversible Give an example of dynamic equilibrium State that equilibria take a finite time to be achieved Use LeChtelier's Principle to predict the effect of a change in the number of moles, volume, or temperature upon the position of an equilibrium Use Le Chatelier's principle to recommend methods to increase the yield of a reaction State that equilibrium is reached when the rate of the forward reaction equals the rate of the reverse reaction Write equilibrium constants (Keq) for a chemical reaction

EquilibriumLeChateliers Principle

CO2 + CaO CaCO3chicken breathfoodegg shell I WISH I HADSWEAT GLANDS.As temperature increases, chickens pant more.This stresses the system and shifts the equilibrium to the LEFTThis makes the egg shells THIN and fragile.

141Le Chateliers Principle

Henri Louis Le Chatelier (1850-1936) was a French professor of chemistry. He studied the gases that caused explosions in mines and the reactions that occur in a blast furnace. These studies led him to put forward the principle that bears his name. The principle explains that if conditions such as temperature or pressure are changed in a system containing reactants and products, then any chemical reaction occurring will tend to counteract the change and return the conditions to equilibrium.

Eyewitness Science Chemistry , Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 39

[ CaO ] , shift[ CO2 ] , shift-- shift ; eggshells are thinnerIn a chickenCaO + CO2 CaCO3 (eggshells)In summer, [ CO2 ] in a chickens blood due to panting.How could we increase eggshell thickness in summer?-- give chickens carbonated water-- put CaO additives in chicken feed-- air condition the chicken houseTOO much $$$-- pump CO2 gas into the chicken housewould kill all the chickens!

I wish I had sweat glands.

142LeChateliers PrincipleN2 + 3 H2 2 NH3 + heat Raising the temperaturefavors the endothermic reaction (the reverse reaction) in which the rise in temperature is counteracted by the absorption of heat.Increasing the pressurefavors the forward reaction in which 4 mol of gas molecules is converted to 2 mol.Decreasing the concentrationof NH3favors the forward reaction in order to replace the NH3 that has been removed.Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532Animation by Raymond ChangAll rights reserved143HENRI LOUIS LE CHATELIERA Man of PrincipleIn the introduction to Linus Pauling's chapter on chemical equilibrium, he gives the student the following advice: The student (or the scientist) would be wise to refrain from using the mathematical equation unless he understands the theory that it represents, and can make a statement about the theory that does not consist just in reading the equation. It is fortunate that there is a general qualitative principle, called Le Chatelier's principle, that relates to all the applications of the principles of chemical equilibrium. When you have obtained a grasp of Le Chatelier's principle, you will be able to think about any problem of chemical equilibrium that arises, and, by use of a simple argument, to make a qualitative statement about it....Some years after you have finished your college work, you may (unless you become a chemist or work in some closely related field) have forgotten all the mathematical equations relating to chemical equilibrium. I hope, however, that you will not have forgotten Le Chatelier's principle.(1) What is this principle which is so highly recommended, and who is its author? The nature of the principle itself is as easy to grasp as it is difficult to state succinctly, and the man himself is almost totally eclipsed by the popularity of his most famous observation. Any system in stable chemical equilibrium, subjected to the influence of an external cause which tends to change either its temperature or its condensation (pressure, concentration, number of molecules in unit volume), either as a whole or in some of its parts, can only undergo such internal modifications as would, if produced alone, bring about a change of temperature or of condensation of opposite sign to that resulting from the external cause.(2) This rather cumbersome expression of the Principle, in a note first published in 1884, seems to show that Le Chatelier himself was the first to experience the frustration of trying to express a generalized form of the concept in a clear and succinct manner. Four years later, in a long article in the Annales des Mines, Le Chatelier restated the Principle in a simpler and more comprehensive form: Every change of one of the factors of an equilibrium occasions a rearrangement of the system in such a direction that the factor in question experiences a change in a sense opposite to the original change.(3) Subsequent authors, particularly those writing textbooks, have opted more for brevity than rigor, usually emphasizing the qualitative cause and effect nature of the idea: If the conditions of a system, initially at equilibrium, are changed, the equilibrium will shift in such a direction as to tend to restore the original conditions.(4) or, even more briefly, If a stress is applied to a system at equilibrium, then the system readjusts, if possible, to reduce the stress.(5) or, Le Chatelier's principle is easier to illustrate than to state. Nevertheless, let us begin with this statement:If a "stress" is applied to a system at equilibrium, the equilibrium condition is upset; a net reaction occurs in that direction which tends to relieve the "stress," and a new equilibrium is obtained.(6) This last example voices the realization all teachers of chemistry have discovered for themselves, that the generalization is easier to teach with examples than words. In fact, the long 1888 paper contains a large number of applications to equilibria including ones involving electromotive force as well as pressure and temperature. The extension of this argument to include applications to biology and social sciences have been made by later writers and were not represented in Le Chatelier's original explanation. One of the most promising areas of research at this time was the possibility of synthesizing ammonia from atmospheric nitrogen and hydrogen. It also presented an excellent study, then and now, in equilibrium. N2 (g) + 3H2 (g) 2NH3 (g) + heat From the equation, we can infer that,.according to Le Chatelier's generalization, increasing the pressure should have the effect of favoring the production of ammonia, because a contraction from four volumes to two volumes occurs during the reaction. As in any exothermic reaction, the temperature would have to be limited to prevent the reverse (endothermic) reaction from being favored, but not kept so low as to seriously inhibit the rate of production. Considering these problems, and informed by the notes of Thenard who had shown complete dissociation of ammonia at 600oC in the presence of metallic iron, Le Chatelier in 1901 attempted the direct combination of the two gases at a pressure of 200 atmospheres using Thenard's temperature. The mixture of gases was forced by an air compressor into a steel Berthelot bomb, where they and the reduced iron catalyst were heated by a platinum spiral. An accidental contamination of the reaction chamber with air resulted in an explosion which blew fragments of the steel container through the floor and ceiling. Le Chatelier abandoned the project, and less than five years later, Haber and Claude were successful in producing ammonia on a commercial scale, acknowledging that the account of Le Chatelier's failed attempt had accelerated their research. Near the end of his life, Le Chatelier wrote, "I let the discovery of the ammonia synthesis slip through my hands. It was the greatest blunder of my scientific career." The great American chemist Josiah Willard Gibbs had anticipated in a mathematical way Le Chatelier's main results, but the fact that he had published in a journal with extremely limited readership outside the United States, combined with the abstract and difficult presentation of his ideas, limited its value to the audience Le Chatelier served, the practicing chemist. Le Chatelier himself, however, was very interested in thermodynamics and was an early champion of Gibbs' work in France, being responsible for the first translation of his papers into the French language. Le Chatelier's interest in and use of the mathematics of thermodynamics may be surprising to teachers who, informed by the standard textbook treatment of his Principle, use Le Chatelier as an example of a qualitative treatment of equilibrium. However, a more thorough look at his life and career reveals a man who consistently integrated theory with practice and whose most successful research was directed toward the problems of industry. Born in Paris on October 8,1850 to a family of architects and engineers, Henry Louis Le Chatelier received his early training in mathematics and chemistry from his father, Louis Le Chatelier, an accomplished engineer. He assisted his father while the latter helped to create the aluminum industry in France, and thus gained much first-hand information about metallurgy. His mother was a rigid disciplinarian and a devout Catholic whose family background and love of poetry fostered in her son that appreciation of art and letters which was evident throughout his life. Le Chatelier's career was marked by a disdain for unsubstantiated speculation, coupled with an almost uncanny sense of the interrelationship between theoria and praxis in science. After finishing his formal education at the Ecole Polytechnique, the Ecole des Mines, and the College de France, his intention was to devote himself to government service as a mining engineer. After two years in the Corps des Mines at Besancon, however, he was offered, much to his surprise, a position as professor of chemistry at the Ecole des Mines. He spent the rest of his life in Paris, where he lectured at the Ecole polytechnique in 1882, was made professor at the College de France in 1883 and became professor at the Sorbonne in 1887. In the same year, he returned to the Ecole des Mines as Professor of Industrial Chemistry and Metallurgy. In 1889 he returned to the College de France where he remained until 1908 as Professor of Inorganic Chemistry. His choice of research projects clearly reflects his appreciation of the interplay between theoretical science and its practical applications. Indeed, he consistently chose research problems of wide interest which seemed to give promise of industrial applications. His very first area of investigation, for example, brought him into the largely uncharted area of the chemistry of cements. He began by repeating the experiments of Lavoisier and Payen on the preparation of plaster of Paris. He discovered that good plaster of Paris consists of the hemihydrate of calcium sulfate, which he identified, and not of the anhydride as was previously believed. These investigations led him to a theory applicable to the setting of many kinds of cements. By this theory, when a cement comes into contact with water, a supersaturated solution is formed which deposits a less soluble hydrated material. This process of solution and solidification results, Le Chatelier explained, in the production of an interlaced, coherent mass of minute crystals. His studies with cements also give evidence for the extraordinary practical skill with which Le Chatelier selected, modified, or in some cases actually invented the instruments needed to carry on his researches. In his work with cements, for example, he chose a method of thermal analysis devised by Regnault but little known at the time. The method would eventually prove of even greater importance in the study of steel and alloys. Methods known at the time for determining high temperatures, such as the use of gas thermometers and of thermocouples, soon proved inaccurate for Le Chatelier's work with cements, as Regnault himself had predicted. It was left to Le Chatelier to isolate the source of the erratic behavior of the instruments. This he did, and the determination of high temperatures soon became a routine procedure following the development of the platinum-platinum-rhodium couple. As an outgrowth of his researches on cements, Le Chatelier significantly widened the range of applications thermodynamics to chemistry. He reasoned that thermodynamics should yield valuable information about those chemical phenomena which were of most concern to him in his study of cements, such as the solubility of salts and their reaction with water. The result of these investigations are seen in his well-known principles of equilibria and the displacement of equilibria to which his name is attached and for which he is certainly best known among students of chemistry. Given his keen sense for the applications of chemical theory, it is perhaps not surprising that he should have accomplished the synthesis of ammonia in 1901, anticipating Fritz Haber as was discussed above. Without his equilibrium principle, it is possible that the practical applications of the phase rule and phase law diagrams would have remained hidden for quite some time. Another major area of research for Le Chatelier was in the combustion and explosion of gaseous mixtures. Once again, his research was closely intertwined with a practical concern for a particular problem. The specific problem that directed his attention into this area was a series of coal mine disasters in France. He attacked the problem by initiating a scientific study of the combustion of methane. Together with Mallard, a professor at the Ecole des Mines, he determined the temperature of ignition, the explosive ratio of air and methane, the speed of propagation in explosions, the explosive pressure, etc. As was often the case with Le Chatelier's scientific investigations, his pioneering work in a new field required the creative development of the appropriate apparatus. Having done this, Le Chatelier and Mallard were able to extend their techniques to the study of hydrogen, carbon monoxide, acetylene, and cyanogen. During these studies, Le Chatelier also developed devices for detecting and determining small quantities of marsh gas. The use of these instruments contributed greatly to the safety of mines, as did the development by Le Chatelier and Mallard of safer, more suitable explosives for use in the mines. Working with Mallard on the allotropic transformations of crystalline materials, such as quartz, Le Chatelier was once again in the position of needing an instrument capable of following these transformations at temperatures above the 200 to 300 degree maximum allowed by the technique then available, polarizing microscopy. He met the challenge by contriving the so-called differential dilatometer which allowed him to follow the transformations by observing the expansion rates of a given specimen relative to the refractory material on which it rested. These investigations led to his study of the thermal expansion of glasses, to the complex reactions that take place in the production of ceramics, and then to an extended study of alloys: their thermal expansion, electrical conductivity, thermoelectric potentials, cooling and heating characteristics, tempering, and annealing. While working in the field of metallurgy he greatly improved upon the techniques then available for microscopic metallography. His improved microscope revealed the formation of compounds between iron and carbon in steel and proved the value of heat treatment in steel. This empirical data, coupled with his theoretical application of the phase rule to the allotropic transformations of a wide variety of alloys, proved to be of enormous value to the world of industry. Still another contribution to the scientific progress of metallurgy can be seen in his creation of the Revue de Metallurgie, which was published beginning in 1904 with Le Chatelier as editor for the next ten years. Besides his many scientific papers, he wrote books on metal alloys, steel, clays and ceramics and acted as an industrial consultant in the manufacture of steel, cement and synthetic ammonia. Altogether Le Chatelier published over five hundred journal articles and books. They included not only works on chemistry and ceramics, but also numerous biographies and, toward the end of his life, articles on social welfare, the scientific management of industries ("Taylorism"), the interrelationship between pure and applied science, and the relation of science to economics.(7) Throughout his life, Le Chatelier was a leader in progressive movements, giving much of his time and effort to causes which he believed worthwhile. In particular, he wrote and spoke extensively on educational reform, taking the lead through the example of his own highly respected career as a teacher. At the Sorbonne, where he was made professor in 1887, he directed the work of over one hundred graduate students during the period from 1908-1922. Ralph Oesper aptly describes his contribution to educational theory and practice in these words: Le Chatelier's teaching opened a new era in chemical education. Enumeration of compounds, their properties, methods of preparation--such was the unappetizing content of the lectures usually offered to the students. He refused to follow the fashion whose effect was the alienation of the student's interest. His courses were built around general laws and principles and he presented facts only as applications of these. He stressed the value of these laws in predicting new facts and emphasized the necessity of precise measurements since these alone can lead to valid general conclusions. Never did he usurp the function of a dictionary or an encyclopedia, and his sole function was to develop the reasoning power rather than the memory of those to whom he was unfolding the beauties of chemistry.(8) A modest man himself, Le Chatelier advised his students to be content with adding a bit to the structure of science, keeping their eyes open for the unusual in nature, but avoiding any deliberate grasping for "sensational discoveries, which do not come by merely wishing for them."(9) He expressed this and other of his lifetime ideals on the occasion of the fiftieth anniversary of his graduation from the Ecole Polytechnique, celebrated by the French Academy of Sciences in 1922. Some of what he said on that occasion has been summarized by Alexander Silverman in a paper presented before the Division of the History of Chemistry of the American Chemical Society on September 6, 1937: He stressed the importance of discipline which his parents had imposed upon him and which was practiced at the Ecole Polytechnique while he was a student. He deplored the decreasing seriousness of study and the increasing tendency toward pleasure and even license in modern colleges and universities. He likened the irresponsible student to a bold individual who dodges vehicles in crossing a street and risks being crushed, to say nothing of the fact that he seriously ties up traffic. Of himself Le Chatelier said, '. . . throughout my scientific career I strove without any desire of the sensational, contenting myself each day with the conscientious pursuit of the task of the day. In the end I was amply rewarded.'(10) Bibliography1. L. Pauling, College Chemistry, 3rd ed., Freeman, San Francisco, CA, 1964, pp. 437-438. 2. H. L. Le Chatelier, Comptes rendus, 99, 786 (1884). 3. H. L. Le Chatelier, Annales des Mines, 13 (2), 157, (1888). 4. L. Pauling, p. 44. 5. M. J. Sienko and R. A. Plane, Chemical Principles and Properties, 2nd ed., McGraw-Hill, New York, NY, 1974, p. 216. 6. R. H. Petrucci, General Chemistry: Principles and Modern Applications, 1st ed., Macmillan, New York, NY, 1972, p. 275. 7. "Henri Le Chatelier: His Publications," Ceram. Abs., 16, (Oct., 1937). 8. Ralph E. Oesper, "The Scientific Career of Henry Louis Le Chatelier," J. Chem. Ed., 1931, 8, p. 451. 9. R. Oesper, op . cit., p. 444. 10. Alexander Silverman, "Henri Le Chatelier: 1850 to 1936" in Aaron J. Ihde and William F. Kieffer, ed., Selected Readings in the History of Chemistry, Division of Chemical Education, A

acids, bases, and salts

Science

tannic acid

sulfuric acid

citric acid

acetic acid

polyprotic acid

strong acids

ph scale

acidbase indicator