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Acids and Bases

Acids and Bases. I.  THEORIES OF ACIDS & BASES          A.  Arrhenius theory          B.   Bronsted -Lowry theory                    1.  conjugate pairs                    2.   amphoterism II.   II.  pH SCALE, NEUTRALIZATION & SALTS III .  pH &  pOH IV.  TITRATIONS.

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Acids and Bases

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  1. Acids and Bases I.  THEORIES OF ACIDS & BASES         A.  Arrhenius theory         B.  Bronsted-Lowry theory                   1.  conjugate pairs                    2.  amphoterism II.  II.  pH SCALE, NEUTRALIZATION & SALTS III.  pH & pOH IV.  TITRATIONS

  2. I.  THEORIES OF ACIDS & BASES • The Arrhenius Theory • Acids increase concentration of H+ ions in aqueous solution. • Acids have more H+ ions than OH- ions. • Formulas for acids begin with H • Bases increase concentration of OH- ions in aqueous solution. • Bases have more OH- ions than H+ ions. • Formulas for bases end with OH • The Bronsted-Lowry Theory • An acid is a proton (hydrogen ion) donor • A base is a proton (hydrogen ion) acceptor • The Bronsted-Lowry theory doesn't go • against the Arrhenius theory in any way • Hydroxide ions are still bases because they • accept hydrogen ions from acids and • form water • An acid produces hydrogen ions in solution • because it reacts with the water molecules • by giving a proton to them.

  3. I.  THEORIES OF ACIDS & BASES (Cont’d) Conjugate pairs An acid is a proton donor, so its conjugate base is the formula for the acid minus a proton Ex. Acid (A) = HCl; Conjugate Base (CB) = Cl- A base is a proton accepter, so its conjugate acid is the formula for the base plus a proton Ex. Base (B) = NO3- ; Conjugate Acid (CA) = HNO3

  4. I.  THEORIES OF ACIDS & BASES (Cont’d) Amphoterism Water, as well as other amphoteric substances, act as both an acid and a base Examples: 1. HNO3+ H2O ←→ NO3- + H3O+ Water acts as a base (proton accepter) 2. NH3+ H2O ←→ NH4++ OH- Water acts as an acid (proton donor)

  5. II.  pH SCALE, NEUTRALIZATION & SALTS pH Scale • It is a logrithmatic scale (powers of 10) • Solutions with a pH of 8 have a ten times higher concentration of [OH-] than solutions with a pH of 7 • Solutions with a pH of 6 have a ten times higher concentration of [H+] than solutions with a pH of 7 Neutralization and Salts • Highly basic or acidic solutions must be neutralized before they can be disposed of because it is HORRIBLE for the environment • They are neutralized by combining together to form water and salt for safe disposal • Example: HCl + NaOH→ H2O + NaCl

  6. III.  pH & pOH • pH: the negative log of the hydrogen ion concentration( -log [H+] ) • pOH: the negative log of the hydroxide ion concentration ( -log [OH-] ) • Self-Ionization of Water • H2O + H2O ←→ H3O+ + OH • Water molecules have the ability to self-ionize and do so all the time! • In one liter of pure water at 25⁰C: • the concentration of H3O+ (H+) = the concentration of OH-= 0.0000001M • or • [H3O+] = [H+] = 1 x 10-7 M

  7. III.  pH & pOH (Cont’d) Equations that will be helpful when solving pH and pOH problems: 1. pH = - log [H+] 2. pOH= - log [OH-] 3. [H+] = 10-pH 4. [OH-] = 10-pOH 5. pH + pOH = 14 6. [H+] [OH-] = 1 x 10-14

  8. IV.  TITRATIONS TITRATION: A process used in a Chemistry lab that determines the concentration of an acid or a base • Chemists use a base or an acid with a known concentration to neutralize the acid or base with an unknown concentration naMaVa= nbMbVb n = number of H’s in an acid formula or OH’s in a base formula M = molarity V = volume Example: What is the molarity of a H2SO4 solution if 190 mL of the acid is needed to exactly neutralize 150 mL of a 2.5 M NaOH solution? MaMaMa = 0.99M

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