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Acids, Bases & Salts

Acids, Bases & Salts. Chapter 19 Pages 586-629. 19.1 Acid-Base Theories. Acids - taste sour, will change the color of an acid-base indicator (red), and can be strong or weak electrolytes in aqueous solutions Common acids: vinegar, carbonated drinks, citrus foods. Common Acids & their Names.

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Acids, Bases & Salts

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  1. Acids, Bases & Salts Chapter 19 Pages 586-629

  2. 19.1 Acid-Base Theories • Acids - taste sour, will change the color of an acid-base indicator (red), and can be strong or weak electrolytes in aqueous solutions • Common acids: vinegar, carbonated drinks, citrus foods

  3. Common Acids & their Names • HCl - hydrochloric acid • HNO3 - Nitric acid • H2SO4 - sulfuric acid • H3PO4 - Phosphoric acid • CH3COOH - ethanoic acid (vinegar) • H2CO3 - Carbonic acid

  4. Bases • Defined by their bitter taste, slippery feeling; will change the color of an acid-base indicator (blue color), can be strong or weak electrolytes in aqueous solution • Soap, salt water, milk of magnesia

  5. Arrhenius Acids & Bases • Arrhenius gave a more specific definition to acids and bases aside from their feel and taste. • He said that acids are hydrogen-containing compounds that ionize to yield H+ in aqueous solutions. • He said that bases are compounds that ionize to yield OH- in aqueous solutions

  6. Classification of acids • HCl - hydrochloric acid (monoprotic) • HNO3 - Nitric acid (monoprotic) • H2SO4 - sulfuric acid (diprotic) • H3PO4 - Phosphoric acid (triprotic) • CH3COOH - ethanoic acid (vinegar) (triprotic) • H2CO3 - Carbonic acid (diprotic) • You write the definition of mono-, di- and triprotic

  7. Bronsted-Lowry Acids & Bases • Used because Arrhenius isn’t very detailed • Defines acids as H+ donors and a base as a H+ acceptor

  8. Conjugate Acids & Bases • NH3 + H2O --> NH4+ + OH- • base + acid --> conj acid + conj base

  9. Lewis Acids & Bases • Acids accept a pair of electrons during a reaction • Bases donate a pair of electrons • H+ + OH- --> H2O • Lewis acid + Lewis base --> cmpd

  10. 19.2 Hydrogen Ions & Acidy • Inside a container of water, occasional collisions between water compounds occur with enough energy to form ions. • HOH + HOH  H3O+ + OH- • H3O+ are known as hydronium ions (can be simplified to H+) • This happens in small occasions and will occur in equal concentrations of ions.

  11. Concentration of Ions • Water acts as the neutral solution when calculating the pH of any solution. • The concentration (mol/L) of the ions (H3O+ and OH-) are each said to be 1 x 10-7 M. • The overall concentration of water is then calculated [H3O+] x [OH-] = 1 x 10-14 M= Kw • This is called the ion-product constant for water – Kw.

  12. Acidic Concentrations • Acidic solutions are when the [H+] is greater than the [OH-]. • The [H+] will be greater than 1 x 10-7 M. • The [OH-] will be less than 1 x 10-7 M. • Ex: HCl  H+ + Cl-

  13. Basic Concentrations • Basic solutions are when the [H+] is less than the [OH-]. (Are also called alkaline solutions.) • The [H+] will be less than 1 x 10-7 M. • The [OH-] will be greater than 1 x 10-7 M. • Ex: NaOH  Na+ + OH-

  14. Calculating the pH value • pH values range from 0 – 14 where 0 is extremely acidic, 7 is neutral and 14 is extremely basic • pH = -log[H+] • (Log is a button on your calculator) • pOH = -log[OH-] • pOH + pH = 14

  15. Calculations • Using [H+] to solve for pH • Using [OH-] to solve for pOH • Using [OH-] to solve for pH • Using [H+] to solve for pOH • Using pH to solve for [H+] • Using pOH to solve for [OH-]

  16. Acid-Base Indicators • Solutions can cause litmus papers to change color indicating acid/base. • Indicators can be added to a solution to produce a color change of the solution at specific pH values. • pH meters can register an accurate measurement of a pH value.

  17. 19.3 Strengths of Acids & Bases • Strength of acids or bases depend on how well the compound ionizes in water. • Strong Acids - completely ionize in water (totally break apart creating a charged particle in the water) HCl • Weak Acids - ionize only slightly (may still have a H-ion that could be removed)

  18. Equilibrium-Constant Expression

  19. Acid Dissociation Constant • Tells whether an acid is strong or weak - stronger acids have higher values.

  20. Sample problem • A 0.1000 M solution of ethanoic acid is only partially ionized. From measurements of the pH of the solution, [H+] is determined to be 1.34 x 10-3 M. What is the Ka?

  21. Solving the problem • [H3O+] = 1.34 X 10-3 M • [ACID] = 0.1000 M • [A-] = (0.1000 - 1.34 X 10-3 M) = 0.0987

  22. Base Dissociation Constant • Strong bases - dissociate completely into metal ions and OH- ions. • Weak bases - dissociate into the conjugate bases and OH- ions.

  23. 19.4 Neutralization Reactions • Acid-base Reactions • Mixing strong acids w/ strong bases will produce a salt and water. • The final product will be neutral as long as the correct mole:mole relationship is followed. • HCl + NaOH --> NaCl + H2O

  24. Titration • Using a buret to add acid to a base to find the equivalence point (where a color change occurs) • The will allow the [base] to be determined because the [A] is already known. • Titration

  25. Steps of Titration • A known [A] is added to a flask (standard solution) • Several drops of an indicator are added • Measured values of base are mixed with the acid until a slight color change is noticed. (end point)

  26. Calculating unknown [B] A 25-mL soln of H2SO4 is completely neutralized by 18 mL of 1.0 M NaOH. What is the concentration of the H2SO4 soln?

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