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Acids & Bases

Acids & Bases

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Acids & Bases

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  1. Acids & Bases PROPERTIES & STUFF

  2. Properties of Acids & Bases • There are two types of specialized solutions (acidic and basic) • Acids have common properties: • taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases • Bases have common properties: • feel slippery, change litmus blue, and become less basic when mixed with acids

  3. Names & Formulas of Acids • An acid is a solute that ionizes when aqueous, producing H+ ions • Therefore the chem formulas of acids are of the general form HX • Where X is a monatomic or polyatomic ion • When the compndHCl (g) dissolves in water to form HCl (aq), it is named as an acid.

  4. Names & Formulas of Acids

  5. Names & Formulas of Bases • A base is a solute that when aqueous produces OH-1 in water • We use the ionic compound naming rules to name a base • The name of the cation followed by the name of the anion • NaOH  sodium hydroxide

  6. Properties of Acids/Bases

  7. Hydrogen Ions and Acidity • We’ve learned that water is a collection of polar molecules in constant motion connected by hydrogen bonds • Collision theory indicates that: • Occasionally, the collisions between water molecules are energetic enough to transfer a hydrogen from one water molecule to another • A water molecule that loses a hydrogen becomes a OH- ion

  8. Hydrogen Ions and Acidity • A water molecule that gains a hydrogen becomes a positively charged ion, AKA hydronium, H3O+ • This reaction is called the self-ionization of water • Establishes an equilibrium

  9. Hydrogen Ions and Acidity • The Hydrogen ions in aqueous solution have several aliases. • Protons, Hydrogen ions, Hydronium ions, & Solvated protons • And can symbolized by • H3O+ and/or H+ • In pure water, this self-ionization occurs to a very small extent • [H3O+]=1.0x10-7 • [OH-]=1.0x10-7

  10. Hydrogen Ions and Acidity • Notice the concs of the two components are equal • This described as a neutral soln • For aqueous solns, the product of [H+] & [OH-] equals 1.0x10-14 • [H3O+][OH-1] =1.0x10-14 • The product of the conc of the ions will always equal 1.0x10-14 • called the ion-product constant for water (Kw) Kw=[H3O+][OH-1]=1.0x10-14

  11. Hydrogen Ions and Acidity • Therefore the ions are interdependent • when [H3O+]increases then [OH-] decreases • If additional ions of either component are added the equilibrium shifts to compensate

  12. Hydrogen Ions and Acidity • Of course not all solns are neutral • When a substance dissolves in water, which contributes H+, the [H+] increases, so it produces an acidic soln([H+] > [OH-]) • When a substance dissolves in water and contributes into OH-, the [OH-] increases, so it produces a basic soln([OH-] > [H+])

  13. Hydrogen Ions and Acidity

  14. Hydrogen Ions and Acidity

  15. The pH Concept • Using conc to express the hydrogen ion content is difficult • A more widely used system is the pH (Potential Hydrogen) scale • A logarithmic scale (log base 10) • Ranges from 0 to 14 (can be <0 and > 14 • [H+]=[OH-] corresponds to 7 on the pH scale (neutral) • pH of 0 is considered highly acidic • pH of 14 is considered highly basic

  16. The pH Scale

  17. The pH Concept

  18. The pH Concept • Calculating the pH of a solution is straightforward • pH = -log[H3O+] • You can also calculate pOH, but it isn’t used as often • pOH = -log[OH-1] • pOH would be a scale to decide how basic a substance is • If you know the pH you can calculate pOH automatically • pH + pOH = 14

  19. The pH Concept • Remembering the equilibrium of H3O+ and OH- and that the sum of the pH and pOH always equals 14 • Using some simple math, we can bounce back and forth between pH, pOH, [H+], & [OH-]

  20. =14 - pH pH pOH =14 - pOH =10 -pH = -log[OH-] =10 -pOH = -log[H+] = 1x10-14/[H+] [H+] [OH-] = 1x10-14/[OH-]

  21. A pH Example: Given the following info fill in the missing pieces. =14-pH 2 =14-12 1x10-12 1x10-2 =10-pOH =10-2 B =10-pH =10-12 3.98x10-12 A =10-11.4 =10-pOH 11.4 2.6 =14-2.6 =14-pH -log(.0025) -log[H+]

  22. pH Practice… • A student is making a solution of calcium hydroxide. She mixes 7.55 grams of calcium hydroxide into 500 ml of water. • What is the concentration of the solution? • What is the [OH-]? • What is the pOH of the solution? • What is the pH of the solution? • What is the [H3O+]? • If she then dilutes the solution to ¼ of its original concentration what is the [OH-]? • What is the new pOH & pH of the solution?

  23. The pH Concept • People need to be able to measure the pH of the solns they use • maintaining the correct acid-base balance in a pool • Creating soil conditions ideal for plant growth • Making medical diagnoses • Indicators or pH meters are often used in measuring pHs • An indicator is an acid or base that changes color in a known pH range

  24. Acid-Base Indicators

  25. Acid-Base Indicators • Knowing the range over which the color change occurs gives a rough estimate of pH • Although indicators are useful tools, they are limited • Some are dependent on temp • If the soln being tested isn’t colorless, the indicator may not show up well • Dissolved salts in a soln may affect the dissociation of the indicator

  26. Acid-Base Indicators

  27. Acid-Base Indicators • A pH meter usually gives a more accurate, more precise measurement of pH • The color and cloudiness of the unknown solution aren’t an issue • Meters are used in hospitals, sewage plants, industry, etc. • When concs of .01M matter

  28. Using Indicators

  29. Theories: Arrhenius Acids & Bases • Chemists had recognized the properties of acids & bases • but they were not able to explain the chemical theory of this behavior • Svante Arrhenius proposed a new way of thinking about acids & bases • Acids produce H+ in soln • Bases produce OH- in soln

  30. Theories: Arrhenius Acids & Bases • The table on the next slide lists some common acids • An acid that contains one ionizable hydrogen is called a monoprotic acid • An acid that contains two ionizablehydrogens is called a diprotic acids • Three ionizable H+ are called triprotic acids

  31. Theories: Arrhenius Acids & Bases

  32. - + + + ⇌ HCl+ H2O ⇌ H3O++Cl- + + ⇌ CH4+ H2O ⇌ CH4+ H2O Theories: Arrhenius Acids & Bases • Not every hydrogen is created equal • Only those Hs attached to highlyelectronegative atoms are acidic

  33. Theories: Arrhenius Acids & Bases • Arrhenius bases are solublehydroxides • The most common in sodium hydroxide • Sodium reacts with water to produce sodium hydroxide • Extremely caustic, commonly known as lye, a major component of products used to clean clogged drains

  34. Arrhenius Acids/Bases

  35. Theories: Bronsted-Lowry A & B • The Arrhenius definition of acids and bases is not a very comprehensive one • It’s too narrow a definition, doesn’t include all substances with acidic or basic properties • For example solutions of NH3 and NaCO3 are basic, but neither contain the hydroxide ion

  36. Theories: Bronsted-Lowry A & B • 2 chemists in 1923, independently proposed an alternative theory (Bronsted & Lowry Theory) • An acid is a hydrogen ion donor • A base is a hydrogen ion acceptor • This new theory allows for ammonia’s basic character, and other discrepencies • When ammonia is dissolved in water it accepts a hydrogen ion from the water

  37. Bronsted-Lowry Acids/Bases

  38. NH3+ H2O ⇌ NH4++ OH- + - + + ⇌ Theories: Bronsted-Lowry A & B • The acceptor (NH3) is labeled a Bronsted-Lowry base • The donor (H2O) is labeled a Bronsted-Lowry acid • H+ are transferred from H2O to NH3 • Causes the OH-conc to be greater than it is in pure H2O • therefore, ammonia solns are basic makes soln basic acceptor donor

  39. Theories: Bronsted-Lowry A & B • The ammonia interaction with water is in equilibrium • The NH4+ will donate its acquired H+ to the OH- to give NH3 & H2O • In the reversedirection NH4+ acts as a B-L acid; & the OH- acts as a B-L base • The reverse direction’s compon-ents become conjugates of the parent acids and bases • A conjugate acid is the particle that results from a base accepting a H+

  40. Theories: Bronsted-Lowry A & B • A conjugate base is the particle that results from an acid donating its H+ • 2 components related by the loss or gain of a single H+ are called conjugate acid/base pairs • The NH3 molecule and the NH4+ ion are a conjugate acid/base pair • The H2O molecule and OH- ion are also a conjugate acid/base pair

  41. Conjugate acid accepts a H+, so it’s a base donates a H+,so it’s an acid Conjugate base Theories: Bronsted-Lowry A & B Acid - Base Pairs NH3 + H2O ⇌NH4+ + OH-

  42. Conjugate base donates a H+, so it’s an acid accepts a H+,so it’s a base Conjugate acid Theories: Bronsted-Lowry A & B Acid - Base Pairs HCl + H2O ⇌Cl- + H3O+

  43. Conjugate Acid-Base Pairs

  44. Theories: Bronsted-Lowry A & B • Did you notice that when water was interacting with NH3 it was acting as an acid, and that when it interacted with HCl it acted as a base? • A substance that can act as both an acid and a base is amphoteric • if an acid is present H2O acts as a base/if a base is present H2O acts as an acid

  45. •• •• O H H •• - H+ + O H •• Theories: Lewis acid/base • In the 3rd def Gilbert Lewis defined an acid as an e- acceptor, while a base is an e- donor • A Lewis acid has too few e-s • A Lewis base has too many e-s Lewis acid Lewis base ⇌

  46. Theories: Acid-Base Definitions

  47. Definitions Practice… Identify (by formula) the conjugate bases of the following acids: HPO4-2 HCl HSO4-1 H2O NH4+1 Identify (by formula) the conjugate acids of the following bases: F-1 NO2-1 CO3-2 S-2 H2O Identify the acids in the following rxn. Next Identify the bases.: HNO2 + H2O ⇌ H3O+ + NO2- Identify the Lewis acid and base in the following rxn: NO+ + NO3−⇌ N2O4

  48. Strong and Weak Acids & Bases • Acids can be classified as weak or strong depending on the degree to which they ionize in water • Strong acids completely ionize in aqueous solns • Weak acids only partially ionize in aqueous solns

  49. Strong and Weak Acids & Bases

  50. Strong and Weak Acids & Bases