Acids and Bases: (An Introduction)

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Acids and Bases(An Introduction)

• Chemistry 12 ? Chapter 14

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Properties of Acids

• Turn blue litmus paper red
• Neutralize the properties of bases
• React with certain metals to produce hydrogen gas
• React with carbonate compounds to produce carbon
  dioxide gas
• Have a sour taste
• Are electrolytes
• Have a pH less than 7

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Common Acids

• Sulfuric Acid H2SO4
• Nitric Acid HNO3
• Phosphoric Acid H3PO4
• Hydrochloric Acid HCl
• Acetic Acid CH3COOH
• Carbonic Acid H2CO3

Battery acid
Used to make fertilizers and explosives
Food flavoring
Stomach acid
Vinegar
Carbonated water
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Properties of Bases

• Turn red litmus paper blue
• Turn the indicator phenolphthalein from colorless
  to red
• Neutralize the properties of acids
• Have a bitter taste
• Are electrolytes
• Are slippery to the touch
• Have a pH greater than 7

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Common Bases
Sodium hydroxide NaOH lye (caustic
soda) Potassium hydroxide KOH lye
(caustic potash) Magnesium
hydroxide Mg(OH)2 milk of magnesia Calcium
hydroxide Ca(OH) 2 slaked lime Ammonia
water NH3 H2O household ammonia
.
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Definition of Acids and Bases

• An operational definition is a definition based
  on observed experimental properties.
• Example An acid is that it is a substance that
  turns blue litmus paper red and has a pH less
  than 7.
• A conceptual definition attempts to explain why a
  substance behaves the way it does.
• We will be looking at a couple of conceptual
  definitions of acids
• For example Arrhenius Theory and Bronsted-Lowery
  Acid Base Theory

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1. Arrhenius Theory

• An acid is a substance that produces H ions in
  solution (H always combine with at least one
  water molecule to produce hydronium ions, H3O.
• HCl(g) H2O(l) -- gt H3O (aq) Cl-(aq)
• A base is a substance that produces OH- ions in
  solution.
• NaOH(s) H2O(l) ? Na (aq) OH-(aq)

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Limitations of Arrhenius Theory

• Acids (like HCl) are able to be neutralized by
  NaOH to produce water
• But what about a base like NH3..
• Not all bases have OH- ions, NH3 in H2O does but,
  if NH3 reacts as a gas with HCl gas, NH4Cl is
  still formed, but is not a solution, so this
  theory wont always work!

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2. Bronsted-Lowry Theory

• An acid is a proton (H ion) donor.
• A base is a proton (H ion) acceptor.

H3O is a hydronium ion (aka. THE PROTON)
HCl is the acid (aka the donor)
H2O is the base (aka the acceptor)
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Conjugate Acid-Base Pairs

• Conjugate linked together
• A conjugate acid is the substance that has
• accepted the proton (gained an H)
• A conjugate base is the substance that has lost
  the proton.
• Each acid has a conjugate base and each base has
  a conjugate acid.

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Conjugate Acid-Base Pairs
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How many protons can be lost?

• Bronsted-Lowry acids can be
• Monoprotic acids are capable of losing one
  proton
• HCl(aq) H2O(l) ? H3O(aq) Cl-(aq)
• HNO3(aq) H2O(l) ? H3O(aq) NO3-(aq)
• Diprotic acids are capable of losing two protons
  (in more then one step)
• H2SO4(aq) H2O(l) ? HSO4-(aq) H3O(aq)
  HSO4-(aq) H2O(l) ? SO42-(aq) H3O(aq)

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Amphoteric

• Substances that can act like an acid in one
  reaction and like a base in another type of
  reaction.
• Example hydrogen carbonate ( HCO3-)
• HCO3- OH- lt -- gt CO3-2 H2O
• (donates a H, so acts like an acid)
• HCO3- H3O lt -- gt H2CO3 H2O
• (accepts a H, so acts like a base)

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3. Lewis Theory

• Bases donate pairs of electrons and acids accept
  pairs of electrons. (Acid/base reaction is the
  donation of an electron pair to create a new
  covalent bond)
• A Lewis acid is any substance, such as the H
  ion, that can accept a pair of non-bonding (lone)
  electrons.
• A Lewis acid is an electron-pair acceptor.
• A Lewis base is any substance, such as the OH-
  ion, that can donate a pair of non-bonding
  electrons.
• A Lewis base is an electron-pair donor.

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