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Acids, Bases, and Salts

Acids, Bases, and Salts. Chapter 14, 15. Some Properties of Acids. 1. The word acid comes from the Latin word acere , which means "sour." All acids taste sour. 2. In 1663, Robert Boyle wrote that acids would make a blue vegetable dye called "litmus" turn red.

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Acids, Bases, and Salts

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  1. Acids, Bases, and Salts Chapter 14, 15

  2. Some Properties of Acids 1. The word acid comes from the Latin word acere, which means "sour." All acids taste sour. 2. In 1663, Robert Boyle wrote that acids would make a blue vegetable dye called "litmus" turn red. 3. Acids react with bases (they destroy the chemical properties of bases). 4. Acids conduct an electric current. 5. Upon chemically reacting with an active metal, acids will evolve hydrogen gas (H2).

  3. Some Properties of Bases 1. The word "base" has a more complex history and its name is not related to taste. All bases taste bitter. 2. Bases are substances which will restore the original blue color of litmus after having been reddened by an acid. 3. Bases destroy the chemical properties of acids (will react with acids) 4. Bases will conduct an electric current. 5. Bases feel “slippery” (soap, bleach) on your skin. This is because they dissolve the fatty acids and oils from your skin and this cuts down on the friction between your fingers as you rub them together.

  4. Some Properties of Salts 1. A salt is the combination of an anion (- ion) and a cation (+ ion). 2. Salts are products of the reaction between acids and bases. 3. Solid salts are usually crystalline. 4. If a salt dissolves in water solution, it usually dissociates into the anions and cations that make up the salt (depends on Ksp)

  5. The Acid Base Theory The three main theories regarding acids and bases are: 1. Arrhenius 3. Lewis 2. Brønsted-Lowry

  6. Arrhenius Theory – late 1890s DEFINITIONS: Acid - any substance which donates hydrogen ions (H+) to water (produces hydronium ions, H3O+): HA → H+ + A¯ Base - any substance which produces hydroxide ions (OH¯) in water. XOH → X+ + OH¯ When acids and bases react, they neutralize each other, forming water and a salt: HA + XOH → H2O + XA

  7. Problems with Arrhenius Theory • The theory did not explain why ammonia (NH3) was a base. • The theory only considers water as a solvent. We know that an acid added to benzene will not dissociate. Solvents are crucial to acid definition. • The end result of mixing certain acids and bases can be a slightly acidic or basic solution. Arrhenius had no explanation for this phenomenon (degrees of acidity).

  8. Brønsted – Lowry Theory – Early 1920s • Two chemists, independent of one another, proposed a new definition of an acid and a base: • An acid is a substance from which a proton can be removed (donates protons). • A base is a substance that can remove a proton from an acid (proton acceptor). *This definition does not require acids and bases to be in aqueous solutions.

  9. Reactions Based on Bronsted - Lowry • Which are the acids and bases?: HCl + H2O → H3O+ + Cl¯ • HCl - this is an acid, because it has a proton available to be transferred (it can give a proton). • H2O - this is a base, since it gets the proton that the acid lost (it has the capacity to accept a proton).

  10. Conjugate acid-base pairs Example: HCl + H2O → H3O+ + Cl¯ • Notice that each pair (HCl and Cl¯ as well as H2O and H3O+ differ by one proton (H+). These pairs are called conjugate pairs. Example: HNO3 + H2O → H3O+ + NO3¯ • The acids are HNO3 and H3O+ and the bases are H2O and NO3¯. What are the pairs?

  11. Bases and Conjugate Acid

  12. Lewis Theory –Early 1920s • Remember drawing Lewis Dot Structures for ionic and covalent compounds? • Lewis Theory focuses on the nature of electrons rather than proton transfer. DEFINITIONS: • An acid is an electron pair acceptor and a base as an electron pair donor. • Lewis Theory is much more general and apply to reactions that do not involve hydrogen or hydrogen ions.

  13. Lewis acid-base reaction: BF3 accepts an electron pair from ammonia: A Lewis acid must have an empty orbital to accept an electron pair. A Lewis base must have a pair of unshared electrons that can be donated. Typical Lewis bases are OH-, H2O, NH3, Cl-, CN-… due to lone pair electrons.

  14. Lewis AB reactions and Formation of Coordinate Complexes The metal ion is a Lewis acid and the ligands coordinated to the ion are Lewis bases.

  15. Autoionization of Water

  16. Strong Acids and Bases • Strong acids are those that ionized completely in water. • The dissociation of a strong base also looks like the diagram at the right in that it dissociates into positive and negative ions.

  17. HNO3 - nitric acid HCl - hydrochloric acidHBr - hydrobromic acidHI - hydroiodic acid H2SO4- sulfuric acid HClO4 - perchloric acid HClO3 - chloric acid (wanna be) 7 Strong Acids • Strong acids are assumed to ionize completely (100%) • in water. They exist as H3O+ ions in water. This is known • as “the leveling effect”. Water has a greater affinity for H+ • than the conjugate bases do.

  18. ANIMATION LINKS • Acid ionization equilibrium demo

  19. Weak Acids and Bases • Some acids and bases ionize only slightly in water. • These are considered weak. • The most important weak base is ammonia.

  20. Balance of ions in solutions Acidic Neutral Solution Solution

  21. Uncommon in labs because too expensive Strong Bases GROUP 1 hydroxides LiOH - lithium hydroxideNaOH - sodium hydroxideKOH - potassium hydroxideRbOH - rubidium hydroxideCsOH - cesium hydroxideBa(OH)2 - barium hydroxideSr(OH)2 - strontium hydroxideCa(OH)2 - calcium hydroxide Some GROUP 2 hydroxides • Strong bases also ionize completely in water, except • for Sr(OH)2 and Ca(OH)2 which are only slightly soluble • (remember Mg(OH)2 is insoluble).

  22. Polyprotic Acids • Polyprotic acids dissociate in a stepwise fashion with different Ka values for each step… In the second and subsequent ionizations the acids are always weak, whether or not the original is a strong or weak acid. • For most of these acids (ex. H3PO4), the first dissociation contributes the significant amount of H+ for pH calculations, and the rest are negligible (except for H2SO4 where second ionization is significant).

  23. Naming Acids -REVIEW -ide ending (elements): “hydro____ic acid” ex. chloride (HCl): hydrochloric acid -ate ending (polyatomics): “______ic acid” ex. chlorate (HClO3): chloric acid -ite ending(polyatomics): “______ous acid” ex. chlorite (HClO2): chlorous acid

  24. Net Ionic Equations -REVIEW • For aqueous acid-base reactions reactions, it is common to write equations in the net ionic form. Standard form: NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l) Ionic form: Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) Na+(aq) + Cl-(aq) + H2O(l) Net ionic form: OH-(aq) + H+ (aq)  H2O(l) (No spectator ions are included)

  25. Things to remember when writing Net Ionic Equations • Binary Acids: HCl, HBr, and HI are strong: all other binary acids and HCN are weak. Strong acids are written in ionic form; weak acids are written in molecular form. • Ternary Acids: If the number of oxygen atoms in an inorganic acid molecule exceeds the number of hydrogen atoms by two or more, the acid is strong (complete dissociation). We will consider all organic acids as weak. • Strong: HClO3, HClO4, H2SO4, HNO3 • Weak: HClO, H3AsO4, H2CO3, H4SiO4, HNO2

  26. Polyprotic Acids: (acids that contain more than one ionizable hydrogen atom. Ex: H2SO4, H3PO4, H2CO3). • Bases: Hydroxides of Group 1 and 2 elements (except Be(OH)2 and Mg(OH)2) are strong bases. All others including ammonia, hydroxlamine, and organic bases are weak. • Salts: Salts are written in ionic form if soluble, and in undissociated form if insoluble. *Know the solubility rules. • Oxides: Oxides are always written in molecular or undissociated form (ex: MgO). • Gases: Gases are always written in molecular form (ex: SO2).

  27. Practice Net Ionic Equations 1. AgNO3 (aq) + H2SO4 (aq)  2. H4SiO4 (aq) + NaOH (aq)  3. HBr (aq) + KOH (aq)  1. Ag+ + HSO4- → AgHSO4(s) 2. H4SiO4 + OH- → H3SiO4- + H2O 3. H+ + OH- → H2O

  28. Weak acids and bases will have Ka or Kb values less than one, but greater than water dissociation, Kw

  29. Relationship between Ka and Kb Ka x Kb = Kw • For any acid and it’s conjugate base, this relationship can be used to determine Ka or Kb. • Ex: NH3 + H2O ↔ NH4+ + OH- NH4+ + H2O ↔ NH3 + H3O+ Kb(NH3)=[NH4+][OH-] Ka(NH4+)=[NH3][H3O+] [NH3] [NH4+] • Therefore, Ka x Kb = [OH-][H3O+]= Kw

  30. Strength of Acid-base pairs • Strong acids yield WEAK conjugate bases… they have a low affinity for H+ • Weak acids yield STRONG conjugate bases • Strong bases yield WEAK conjugate acids • Weak bases yield STRONG conjugate acids

  31. Soren Sorenson (1868-1939) invented the pH scale while creating a way to test the acidity of beer. Beer has a pH of about 4.5. pH Scale • The pH scale (potential hydrogen scale) is a measure of hydronium ion (H3O+) concentration. • Hydronium ion concentration indicates acidity. Each increase in pH # means a 10-fold decrease in [H+]. • The higher the [H3O+], the higher the acidity.

  32. pH scale and [H+]

  33. Calculating pH • The concentration (M or mol/L) of H3O+ is expressed in powers of 10, from 10-14 to 100. • Scientists use pH which is the negative log of [H3O+]. pH = -log[H3O+] • Note: The significant figures for logarithmic numbers are given after the decimal, and the numbers preceding the decimal give the exponent.

  34. Calculating pH of a strong acid: Ex: Given a solution of 0.50M HCl, what is the pH? Step 1: Find [H3O+] in mol/L 0.50mol/L = 5.0 x 10-1 mol/L Step 2: Place value in equation and solve. • pH = -log[5.0 x 10-1] = 0.30

  35. Practice pH Calculations • Find pH of the following solutions if [H3O+] is: • 1.00 x 10-3 • 6.59 x 10-6 • 9.47 x 10-10 • Find [H3O+] if the pH is: • 6.678 3. 10.0 • 2.533 4. 2.56

  36. pOH • You can calculate the pH of a solution if you know the concentration of hydronium ion. [OH-] • If we use the ion product constant of water we can derive this equation: [pH][pOH] = 1.00 x 10-14 • Working with this equation leads to: pH + pOH = 14

  37. Calculating pH of a strong base: Ex: Find the pH of a solution with an [NaOH] of 1.0 x 10-8. Step 1: Solve for [H3O+] in equation: [H3O+] = 10-14 [OH-] Step 2: Place values in: [H3O+] = 10-14 = 10-6 M [1.0 x 10-8]

  38. Step 3: Solve for pH by placing [H3O+] in pH = –log[H3O+] pH = -log(1.0 x 10-6) pH = 6.0

  39. Practice pH Calculations Using pOH • Find the pH of the following solutions with [OH] of: • 1.00 x 10-4 • 2.64 x 10-13 • 5.67 x 10-2 • 3.45 x 10-11

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