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Acids and Bases

Learn about the properties and definitions of acids and bases, including pH scale, common acids and bases, types of acids, and the concepts of Arrhenius, Bronsted-Lowery, and Lewis acids and bases.

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Acids and Bases

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  1. Acids and Bases

  2. Acids • Sour tastes • pH: 0 – 6.9 • Reacts with carbonate to produce CO2 gas. • Turns Litmus from Blue to Red. • Reacts with metals to produce hydrogen gas.

  3. Acids • Acids Solutions conduct electricity (In ion form) • Strong (99%) HCl (aq) + H2O (l) H+(aq) + Cl-(aq) or H2SO4(aq) + H2O (l) H+(aq) + HSO4-(aq) • Weak or strong acids depends on how much the acid dissociates.

  4. Bases • Bitter Taste • pH: 7.1 – 14 • Turns Litmus from Red to Blue • Reacts with acids to form salts. • Slippery feel

  5. Bases • Weak or Strong depends on how much it dissociates. • Weak (4%) Na2CO3(aq) + H2O (l) Na+(aq) + CO3-(aq) • Strong (99%) NaOH (aq) + H20 (l) Na+(aq) + OH- (aq)

  6. What is an Acid and What is a Base? • Define acid and base in terms of water. H2O or HOH H-OH <Molecular Structure H+ = Hydrogen (Acid) OH- = Hydroxide (Base) If you have more hydrogen then hydroxide, you have an acid

  7. What is an Acid and What is a Base? • If you have more hydroxide then hydrogen you have a base. H+ = Acid OH- = Base HOH (H2O) = Neutral • We measure acid and Base on the “pH” scale.

  8. What is an Acid and What is a Base?

  9. pH Scale • The pH scale goes from 0 14

  10. pH Scale • Acid: pH 0 – 6.9 • Base: pH 7.1 – 14 • pH of 7 = Neutral pH

  11. Common Acids • Hydrochloric Acid = HCl • Sulfuric Acid = H2SO4 • Nitric Acid = HNO3

  12. Common Bases • Sodium Hydroxide = NaOH • Ammonium Hydroxide = NH4OH • Potassium Hydroxide = KOH

  13. Types of Acids • Monoprotic acids contain only one acidic hydrogen.

  14. Types of Acids • Diprotic contains 2 acidic hydrogens • Triprotic contains 3 acidic hydrogens • Polyprotic- more than 1 acidic hydrogen

  15. Types of Acid • Most acids are oxyacids, where the acidic proton is attached to an oxygen atom. • HNO3 Nitric Acid • H2SO4 Sulfuric Acid • H3PO4 Phosphoric Acid

  16. Types of Acids • Amphoteric- can act as an acid or a base. H2O + H2O H3O+ + OH- Acid (1) Base (1) Acid (2) Base (2) The above is an autoionization of H2O, and involves the transfer of a proton from one water molecule to another to produce a OH- and H3O+.

  17. Buffers and Buffer Solutions • A buffered solution is one that “resists a change is its pH”, when either a hydroxide, OH- or hydrogen, H+, are added. • Blood is a good example of a buffer. • A buffer contains 2-components: an acid to neutralize the addition of OH-, and a base to neutralize the H+ from the addition of an acid.

  18. Buffers and Buffer Solutions • When base (OH-) is added to a buffer solution, the acid in the buffer provides H+ ions, which neutralizes the base, thus, preventing a large change in pH.

  19. Buffers and Buffer Solutions • The blood’s primary buffer system is made up of carbonic acid (H2CO3) and sodium bicarbonate(NaHCO3). H+ + HCO3- H2CO3 From Acid In Buffer Carbonic Acid OH- + H2CO3 HCO3- + H2O From Base In Buffer Bicarbonate Ion

  20. Final pH of Buffer Close to original Original Buffer pH Added OH- ion replaced by acid ion or Added H+ ion replaced by base ion Buffers and Buffer Solutions

  21. Acids and Bases • Bases are ionic compounds containing metal cations and the hydroxide ion, OH-. • When a “Base” completely dissociates in water to produce OH-, it is referred to as Alkaline.

  22. Bronsted-Lowery Acids and Bases • Bronsted–Lowery Acid is a molecule or ion that is a proton (H+) donor. • Bronsted-Lowery Base is a molecule or ion that is a proton (H+) acceptor.

  23. Bronsted-Lowery • Bronsted-Lowery Acid is a molecule that is a Proton Donor. Example HCl + NH3 NH4+ + Cl- • The proton is transferred from the hydrogen chloride to ammonia.

  24. Bronsted-Lowery • Bronsted-Lowery Base is a molecule that is a proton acceptor. Example HCl + NH3 NH4+ + Cl- Proton Proton DonorAcceptor AcidBase

  25. Dissociates / Ionizes • A strong acid is one that ionizes completely in an aqueous solution. • A strong acid is a strong electrolyte. • Electrolyte – Any compound that conducts electricity when melted or dissolved in water.

  26. Dissociates / Ionizes • Acids that are weak electrolytes are known as weak acids. • They do not dissociate or ionize very much. • Dissociation- The separation of ions that occurs when an ionic compound dissolves. • Ionization-The process where ions form from a covalent compound.

  27. Dissociates / Ionizes

  28. Arrhenius Acids and Bases • An Arrhenius Acid is a chemical compound that increases the concentration of hydrogen ions, H+, in an aqueous solution. • Arrhenius Base is a substance that increases the concentration of hydroxide ions, OH-, in an aqueous solution.

  29. Conjugated Acid and Bases • When an acid gives up a proton, it can re-accept the proton and acts as a base. HF + H2O F- + H3O+ Acid Base Conjugated Conjugated Base Acid • In the above reaction the water molecule is a Bronsted-Lowery Base. The hydronium ion is now able to donate a hydrogen proton, so it is called a conjugated acid.

  30. Lewis Acids and Bases • A Lewis acid is an electron-pair acceptor. • A Lewis base is an electron-pair donor. H+ + [ O – H ]- H O H Lewis Acid Lewis Base

  31. Graphic organizer time!

  32. Neutralization Reaction • The reaction of an acid and base is called a neutralization reaction because the properties of both the acid and base are diminished of neutralized when they react.

  33. Concentration of Solution • The concentration of a solution is a measure of the amount of the solute (solid) in a given amount of solvent (Liquid). • Molarity- The number of moles of a solute in one liter of solution. Molarity, M= Moles/Liter

  34. Concentration of Solution • If 3 moles of LiCl are added to 100 L of water, what is the molarity of the solution? 3mol/100 L = 0.03 M of LiCl

  35. Concentration of Solution • If you put 20.0 g of NaOH in 1-liter of water, what is the molarity (M). 20.0g NaOH 1 mole = 0.5 mole 40.0 g NaOH 0.5 mol/1.0 L = 0.5 M of NaOH

  36. Concentration of Solution • If you put 50.0 g of HCl in 2.0 liters of water, what is the molarity (M)? 50.0g HCl 1 mole = 1.37 mole 36.46 g HCl 1.37 mol/2.0 L = 0.69 M of HCl

  37. Concentration of Solution • When making-up a solution: M1V1=M2V2 M1 = Initial or Beginning Molarity V1 = Initial or Beginning Volume M2 = Final or Ending Molarity V2 = Final or Ending Volume

  38. Concentration of Solution • If you have 300mL of a 0.5 M solution of HNO3. What volume of water needs to be added to get 0.8 M solution of HNO3? M1 = 0.5 M V1 = 300 mL M2 = 0.8 M V2 = ? M1V1=M2V2

  39. If you have 56 mL of a 3 M solution of H2SO4. What volume of water needs to be added to get 0.1 M solution of H2SO4? M1 = 3 M V1 = 56 mL M2 = 0.1 M V2 = ?

  40. If you have 56 mL of a 3M solution of H2SO4. What will the concentration of H2SO4 be if we add 1000mL (1L) of water? M1 = 3M V1 = 56 mL M2 = ? M V2 = 1000 mL

  41. If you have 5 mL of a 12 M solution of HCl. What will the concentration of HClbe if we add 500mL (1L) of water? M1 = 12 M V1 = 5 mL M2 = ? M V2 = 500 mL

  42. Acid-Base Titration • The general process of determining the molarity of an acid or a base through the use of an acid-base reaction is called an acid-base titration.

  43. Acid-Base Titration • The known reactant molarity is used to find the unknown molarity of the other solution. • Solutions of known molarity that are used in this fashion are called standard solutions. • In a titration, the molarity of one of the reactants, acid or base, is known, but the other is unknown.

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