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Chapter 15-16 pH Acids and Bases

Burettes are used in titrations, page 498. Chapter 15-16 pH Acids and Bases. Pages 452-507. Acids and Bases are all around us. Our very existence relies upon the interactions of acidic and basic substances.

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Chapter 15-16 pH Acids and Bases

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  1. Burettes are used in titrations, page 498 Chapter 15-16pHAcids and Bases Pages 452-507 Acids and Bases are all around us. Our very existence relies upon the interactions of acidic and basic substances. In neutral water, there is a balance of equal concentrations of H3O+ and OH-. (1x10-7M of each)

  2. Properties of Acids • Acids… • Taste sour (but don't taste them in lab!!) • Form Electrolytes • React with bases to form salts and water. • Turn Blue Litmus Paper to Red • React with metals, dissolving the metal atoms into ions, and releasing H2 gas. • Hydronium ion (H3O+) is the “acid ion”

  3. Properties of Bases • Bases… • Feel slippery • Taste bitter (but don't taste them in lab!!) • Form electrolytes • React with acids to form salts and water • Turn Red Litmus Paper to Blue • Greater threat to biological material due to their hydrophilic nature. • Hydroxide (OH-) “base ion”

  4. Some examples of acids and bases • HNO3 – can give up one H+ to make H3O+ • acid • NaOH – will deposit the OH- ion • base • HCl – can give up one H+ to form a Cl- and H3O+ ions. • acid • On your own, predict if each of the following is an acid or a base: • HIO4, HClO4, Ba(OH)2, H3PO4,Na3PO4.

  5. Strong Acids • Strong acids have a very electronegative end • opposite from the H+ end. • The more polar an acid, the stronger it will be. • Strong acids • dissociate completely • Diprotic, Triprotic, Polyprotic Acids • can give up more than one H+ • H2SO4 is a strong acid. • Why? • HClO4 is one of the strongest acids we can use • Why?

  6. Strong Bases • Strong bases • dissociate completely(dumping OH-). • Weak bases • take H+ to leave OH- behind. • Ca(OH)2 is a strong base. • Why? • K2CO3is a weak base. • Why? • Ammonia, NH3 is a weak base. • How?

  7. Three ways to describe acids/bases • Arrhenius Model of Acids and Bases: • Acids shift a water solution toward H+ imbalance • Bases remove H+ ions from solution, leaving an OH- imbalance. • Bronsted-Lowry Model of Acids and Bases: • Acids are proton donors. • Bases are proton acceptors. • Lewis acid-base model: • Acid = e- pair acceptor • Base = e- pair donor • Acid-Base reactions produce pairs of molecules, known as conjugate pairs.

  8. Conjugate Acid - Base Pairs • Acidsconjugate bases. How? • Basesconjugate acids. How? Acid Base Conjugate Acid Conjugate Base

  9. Identify Conjugate pairs: Base Acid Conjugate Base Conjugate Acid Acid Base Conjugate Acid Conjugate Base

  10. Strong vs Weak • Strong Acid / Base: dissociatecompletely. • Dissociation constant, K: • Numerical constant represents acid and base strength. • Large K = stronger acid/base • Small K = weaker acid/base • Acid strength: Ka • Base strength: Kb • To calculate KaorKb, we use the concentrations of products divided by reactants at equilibrium.

  11. The pH scale • The “p” = “negative log” (-log) • The “H” = hydronium ion [H3O+] concentration. • pH = negative log of the hydronium ion concentration-log [H3O+] • Low pH high H3O+concentration. “Acids” • High pHlow H3O+concentration. “Bases” • When a solution increases in pH, does it become more acidic, or more basic?

  12. The pH scale • Low pHhigh H3O+ concentration. “Acids” • High pHlow H3O+ concentration. “Bases”

  13. pH calculations: “p” = -logarithm • Try These: • What is the pH of a solution whose [H3O+] is 1 x 10-5 M? • What is the [H3O+] concentration of a solution with a pH of 9? • What is the pH of a solution with a [H3O+] concentration of 1.0 x 10-12 M? • What is the pH of a solution whose [H3O+] concentration is 3.0 x 10-3 M? • What is the [H3O+] concentration of a solution whose pH is 6.5? 5 1x10-9 M 12 2.5 3.2x10-7 M

  14. H3O+ and OH- • In pure water: • Perfect balance of H3O+ and OH-. • [H3O+]= 10-7 M and [OH-]= 10-7 M • Acids and Bases: • Shift balance towards [H3O+]or [OH-] • Kw: • Autoionization of water. • = [H3O+]x [OH-]=1x10-14 • pH + pOH = 14 Amphoteric: Substances that can act as an acid or base

  15. Sample H3O+ and OH- Problems • Determine hydronium [H3O+] and hydroxide [OH-] concentrations of the following solutions: • 1.0 x 10-4 M HCl • 1.0 x 10-3 M HNO3 • 3.0 x 10-2 M NaOH • 1.0 x 10-4 M Ca(OH)2

  16. Neutralization • Strong acid + strong basewater + a salt Acid Base Ions Ions Water Salt Water = HOH HOH = H+, OH- Salt = cation + anion

  17. Indicators • Indicators • Are one chemical in an acid, another in a base. • Appear to change color when placed in different concentrations of acids and bases. • Transition interval • the range of pH that they are effective. • Phenolphthalein • colorless in the presence of an acid and magenta in the presence of a base. • Universal Indicator • displays a spectrum of colors, depending upon acid and base concentration.

  18. Titrations • Titrations • analytically determine an acid’s or base’s concentration (M). • far more precise than indicators. • Titrations are performed by… • Adding a known amount of a known-concentration acid to an unknown base or… • Adding a known amount of a known-concentration base to an unknown acid. • If the acid is diprotic or the base is (OH)2… M=molarity or V=volume

  19. Titration Curves • The equivalence point is the point at which: • mole acid = mole base

  20. Acid / Base Buffers, Blood Buffers • Buffers resist changes in pH. • pairs or sets of compounds, similar to the acid or the base that is being buffered. • When an acid is added: • the buffer absorbs the extra H+ ion. • When a base is added: • the buffer releases a H+ ion to react with the OH- of the base and produce H2O. • We have special buffers in our blood to stabilize the pH of our blood. • protect delicate organ systems from pH imbalances from increased or decreased levels of dissolved gases and minerals in our blood. • Our normal blood pH is 7.4

  21. Acid Rain and Stream Water pH • Acid rain • SOx and NOx from air pollution react with water to produce acids. • The worst reported case of acid rain was in West Virginia, measured at a pH of 1.5, slightly weaker than stomach acid.

  22. Effects of Acid Rain on structures • The high acidity dissolves the calcium compounds use in bridges, building mortar, pavement, foundations, weakening the material and putting the structural integrity of the buildings at risk. • Marble and granite used in very old statues is at a greater risk due to the absence of modern compounds used to reduce mineral dissolution. • The thin layers of automotive paint suffer substantial damage from acid rain with only small amount of exposure.

  23. Effects of Acid Rain on Humans • Acidic water vapor in the air we breath causes inflammation in the lungs and deterioration of lung tissue. In those already at risk for respiratory disruption, the effect can be fatal. • The CDC estimates that strict reductions in acid rain and treatment under the Public Health Acid Rain Programwill save the US $50 billion by the year 2010. • Another effect of acid rain is reduced visibility -acidic particles in the air react to produce photochemical smog – a haze that reflects or blocks light like a strong mist on the horizon.

  24. Effects on Forest Ecosystems • The reduced pH of soils • dramatically reduces the ability of trees and underbrush to grow. • As pH decreases in aqueous systems • metals and other positive ions are dissolved more readily and bound up in complex molecules. • As the soils grow more acidic • valuable minerals are dissolved and bound up, • made unavailable to the plants that normally draw the minerals from the soils through their roots.

  25. Effects of Acid Rain on Stream water • Acid rain is most devastating to aquatic life. • Some metal ions will not dissolve into ions above a certain pH. For Aluminum, that pH is about 5.0. • When the pH of a lake drops below this pH, Al lying at the bottom of the lake goes into solution as ions. • When the ion-containing acidic water enters the higher-pH gills, the metal ions come out of solution, forming a film on the gills that suffocates the fish.

  26. Reduction of Acid Rain • Due to multinational cooperation, acid rain is on the decline. • Coal-burning power plants and industrial manufacturing plants must have installed elaborate (and expensive) “scrubbers” to absorb SOxbefore it can be released. • Due to costs of pollutant-reducing equipment, developing nations still release large amounts of SOx pollutants. • Further, basic compounds and acid-base buffers are being added to main lake waters to protect aquatic wildlife. For more info: Page 493, or EPA’s Clean Air Program: http://www.epa.gov/airmarkets/acidrain/

  27. Scrubbers! • We can add limestone, CaCO3 to coal-burning power plants to react with sulfur compounds before they are released. End of chapter 15-16

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