Acids and Bases

1. Acids and Bases

2. Acids: Compounds that contain one or more hydrogen ions (H+) when dissolved in water Bases: Compounds that contain one or more hydroxide ions (OH-) when dissolved in water

3. Properties of Acids • Acids taste sour • Acids have a pH lower than 7 • Acids effect indicators • Blue litmus turns red • Universal indicator turns red • Acids are proton (hydrogen ion, H+) donors • Acids react with active metals, produce H2 • Acids react with carbonates • Acids neutralize bases

4. Acids Effect Indicators Blue litmus paper turns red in contact with an acid.

5. Acids Have a pH less than 7

6. Acids React with Active Metals Acids react with active metals to form salts and hydrogen gas. Mg + 2HCl  MgCl2 + H2(g)

7. Acids Neutralize Bases HCl + NaOH  NaCl + H2O Neutralization reactions ALWAYS produce a salt and water.

8. Acids you must know: Strong Acids Weak Acids Sulfuric acid, H2SO4 Phosphoric acid, H3PO4 Hydrochloric acid, HCl Acetic acid, HC2H3O2 Nitric acid, HNO3

9. Sulfuric Acid • Highest volume production of any chemical in the U.S. • Used in the production of paper • Used in production of fertilizers • Used in petroleum refining

10. Nitric Acid • Used in the production of fertilizers • Used in the production of explosives • Nitric acid is a volatile acid – its reactive components evaporate easily • Stains proteins (including skin!)

11. Hydrochloric Acid • Used in the pickling of steel • Used to purify magnesium from sea water • Part of gastric juice, it aids in the digestion of protein • Sold commercially as “Muriatic acid”

12. Phosphoric Acid • A flavoring agent in sodas • Used in the manufacture of detergents • Used in the manufacture of fertilizers • Not a common laboratory reagent

13. Acetic Acid • Used in the manufacture of plastics • Used in making pharmaceuticals • Acetic acid is the acid present in vinegar

14. Organic Acids Organic acids all contain the “carboxyl” group, sometimes several of them. The carboxyl group is a poor proton donor, so ALL organic acids are weak acids.

15. Examples of Organic Acids • Citric acid in citrus fruit • Malic acid in sour apples • Deoxyribonucleic acid, DNA • Amino acids, the building blocks of protein • Lactic acid in sour milk and sore muscles • Butyric acid in rancid butter

16. Ionization of HCl and formation of hydronium ion, H3O+ H2O + HCl  H3O+ + Cl- Proton acceptor Proton donor

17. Strong Acids vs. Weak Acids Strong acids are assumed to be 100% ionized in solution (good proton donors). HCl H2SO4 HNO3 Weak acids are usually less than 5% ionized in solution (poor proton donors). H3PO4 HC2H3O2 Organic acids

18. Strong Acid Dissociation

19. Weak Acid Dissociation

20. Properties of Bases • Bases taste bitter • Bases have a pH greater than 7 • Bases effect indicators • Red litmus turns blue • Universal indicator turns blue • Phenolphthalein turns purple • Bases are proton (H+) acceptors and usually have hydroxide ions- OH- • Solutions of bases feel slippery • Bases neutralize acids

21. Examples of Bases • Sodium hydroxide (lye), NaOH • Potassium hydroxide, KOH • Magnesium hydroxide, Mg(OH)2 • Calcium hydroxide (lime), Ca(OH)2

22. Bases Effect Indicators Red litmus paper turns blue in contact with a base. Phenolphthalein turns purple in a base.

23. Bases have a pH greater than 7

24. Bases Neutralize Acids Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl. 2 HCl + Mg(OH)2 MgCl2 + 2 H2O

25. Products of Neutralization HCl + NaOH  NaCl + H2O H2SO4 + Ca(OH)2 CaSO4 + 2 H2O HNO3 + KOH  KNO3 + H2O The products of neutralization are always a ______ and _______. salt water

26. Theories of Acids and Bases • Arrhenius (AH!) Theory: • Acids produce hydrogen ions in a solution HBr + H2O  Br - + H2O + H+ • Bases produce hydroxide ions in a solution NaOH + H2O  Na++ H2O + OH-

27. Bronsted (+) Lowry Theory: • Acids are proton donors H2O + HCl H3O+ + Cl- • Bases are proton acceptors  H2O + NH3 NH4+ + OH-

28. Solutions

29. Parts of a Solution • Solute-the part that gets dissolved • Solvent-the part that does the dissolving • Solvation – the process when solute molecules become surrounded by solvent molecules

30. Practice Identify the solute and solvent in the following solutions: • 10.0g of sugar & 40.0g of water solute: ____________solvent: ___________ b) 75g of KBr & 100 g water solute: ____________solvent: ___________ water sugar water KBr

31. Solubility – a measure of how much solute can dissolve in a certain solvent • Substances that dissolve easily in a certain substance are called soluble. • Example: Kool Aid in Water • Substances that do not dissolve in a certain solvent are called insoluble. • Example: Chalk (CaCo3) and Water • Two liquids that dissolve in each other are called miscible. • Example: Ethanol and Water • Two liquids that will not dissolve in each other are called immiscible. • Example: Gasoline and Water

32. Molarity and Dilutions

33. Molarity • Molarity is a measurement of concentration. • Molarity (M) is the number of moles in a liter of solution.

34. Dilutions • Watering down a solution: • Reduces the molarity • Increases the volume • The product of the molarity and volume you start with is equal to the product of the molarity and volume you end with. • The number of moles does NOT change. That is why the equations works.

35. Solubility Curves

36. Solubility Curves • Solubility curves are graphs that show the relationship between solubility and temperature for specific substances.

37. Solution Definitions • Soluble - capable of being dissolved • 2. Solution – homogeneous mixtures • 3. Saturated – exactly the right amount of dissolved solute for a particular solvent • 4. Unsaturated –less solute than can be dissolved by a solvent to be saturated • 5. Supersaturated -more dissolved solute than saturated. Has to be specially made; unstable

38. 3. Based on the solubility chart below, decide whether each of the following is U: unsaturated(Under line), S: saturated (on line), SS: supersaturated (above line), or N: not enough information is given. a) 50 g KCl in 100 g of water at 90°C. ____ b) 50 g KCl in 100 g of water at 60°C. ____ d) 50 g KNO3 in 50 g of water at 60°C. ____ 50 g KNO3 = X X = 100 50 g water 100 g water

39. Polarity

40. Why won’t some things dissolve in water? • “Like dissolves like.” • The solute and the solvent have to be similar to each other.

41. Electronegativity is the attraction between valence electrons and the nucleus of an atom. • Polar molecules have unevenly shared electrons in their bonds. They have + and – areas that are attracted to each other. Water is an example of a polar molecule.

42. Non-polar molecules have evenly shared electrons and there are no partial charges on the molecule. They are not attracted to each other or to polar molecules. Oil is an example of a non-polar molecule. • Polar solvents, such as water will dissolve polar and ionic solutes like salt and sugar. Non-polar solvents such as oils will dissolve other nonpolar solutes like wax.

43. Factors that increase rate of Dissolving for solids in a liquid solution 1. Surface area - Increase of surface area increases rate of dissolving. Increase surface area by breaking/grinding solid 2. Agitation – Shaking/stirring a solution will increase the rate of dissolving 3. Temperature- increase in temperature usually increases the rate of dissolving

44. Factors that increase the rate of dissolving for a gas in a liquid 1. Temperature- increase in temperature usually decreases the rate of dissolving 2. Pressure- Increase in pressure increases the rate of dissolving. ***Pressure has little/no effect on the dissolving rate of a solid in a liquid***