ACIDS AND BASES

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ACIDS AND BASES

• Properties of Acids and Bases
• Acid Base Theories
• Strong and Weak Acids and Bases
• Understanding Indicators
• pH Scale
• Buffers and Antacids

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Properties of Acids and Bases

• Macroscopic View
• Acids
• Taste sour
• Produce painful sensation on skin
• React with certain metals (Mg, Zn, Fe) to produce
  H2 gas
• React with limestone and baking soda to produce
  CO2
• Turn litmus paper red
• Bases
• Taste bitter
• Feel slippery on skin
• React with oils and greases
• Turn litmus paper blue
• React with acids to produce salt and water

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Acids in Every Day Life
Common Acid in the Home Chemical Name Common
Name Hydrochloric Acid (HCL) Muratic Acid Acetic
Acid (CH3COOH) Vinegar Sulfuric Acid
(H2SO4) Auto Battery Acid Carbonic Acid
(H2CO3) Carbonated Water Boric Acid
(H3BO4) Antiseptic Eye Drops Acetylsalicylic
Acid (C16H12O6) Aspirin
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Acids
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Acids
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Acid Nomenclature
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Bases in Every Day Life

• Common Bases in the Home
• Chemical Name Common Name or Use
• Ammonia (NH3) Cleaner
• Sodium Hydroxide (NaOH) Lye
• Sodium Bicarbonate (NaHCO3) Baking Soda
• Magnesium Hydroxide (Mg(OH)2) Milk of Magnesia
• Calcium Carbonate (CaCO3) Antacid
• Aluminum Hydroxide (Al(OH)3) Antacid

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Bases
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Properties of Acids and Bases

• Microscopic View
• You may have noticed that all the acids contain
  hydrogen, while most of the bases contain the
  hydroxide ion (OH-). Two main theories use these
  facts in their descriptions of acids and bases
  and their reactions
• Arrhenius Theory
• Bronsted-Lowery Theory

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Arrhenius Theory Must have Water

• This was the first modern acid-base theory, and
  it tells us that when dissolved in water
• An acid yields H ions
• HCl(aq) ? H Cl-
• A base yields OH- ions
• NaOH(aq) ? Na OH-

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Arrhenius Theory

• This theory also classifies the reaction between
  an acid and a base as a neutralization reaction,
  producing a neutral solution composed of a water
  and a salt.
• HCl(aq) NaOH(aq) ? H2O(l) NaCl(aq)
• The water is formed from the combining of the H
  and OH- ions
• Like all theories, this one has its limitations.
  Some bases dont have hydroxide ions. To account
  for this, a new theory was developed.

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Bronstead Lowery Theory

• In this theory, an acid is classified as a proton
  (H) donor.
• A base is classified as a proton acceptor.
• The base accepts the H by furnishing a pair of
  electrons for a coordinate-covalent bond.

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Conjugate Pairs
All acids have a conjugate base, which is formed
when their proton has been donated likewise, all
bases have a conjugate acid, formed after they
have accepted a proton.
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One More The Lewis Theory

• This theory extends well beyond the things you
  normally think of as acids and bases.
• The theory
• An acid is an electron pair acceptor.
• A base is an electron pair donor.

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Lewis Acid-Base Reactions
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Water

• H2O can function as both an acid and a base.
• In pure water, auto-ionization can occur
• In a neutral solution, H3O OH-

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Strong Weak Acids

• It is important to remember that acid-base
  strength is not the same as concentration.
  Strength refers to the amount of ionization or
  breaking apart that a particular acid or base
  undergoes. Concentration refers to the amount of
  acid or base that you initially have. You can
  have a concentrated weak acid or a dilute strong
  acid

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Strong Acids

• Certain acids are considered to be strong, which
  means they are dissociated 100 in solution
• HCl Hydrochloric Acid
• HNO3 Nitric Acid
• H2SO4 Sulfuric Acid
• HBr Hydrobromic Acid
• HI Hydroiodic Acid
• HClO4 Perchloric Acid
• You ought to memorize this list, because almost
  every other acid is weak. The most common example
  is HCl.

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Weak Acids

• A weak acid is one which doesn't ionize fully
  when it is dissolved in water.
• Ethanoic acid is a typical weak acid. It reacts
  with water to produce hydroxonium ions and
  ethanoate ions, but the back reaction is more
  successful than the forward one. The ions react
  very easily to reform the acid and the water.
• At any one time, only about 1 of the ethanoic
  acid molecules have converted into ions. The rest
  remain as simple ethanoic acid molecules.

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Strong Bases

• A strong base is something like sodium hydroxide
  or potassium hydroxide which is fully ionic. You
  can think of the compound as being 100 split up
  into metal ions and hydroxide ions in solution.

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Weak Bases

• A weak base is one which doesn't convert fully
  into hydroxide ions in solution.
• Ammonia is a typical weak base. Ammonia itself
  obviously doesn't contain hydroxide ions, but it
  reacts with water to produce ammonium ions and
  hydroxide ions.
• However, the reaction is reversible, and at any
  one time about 99 of the ammonia is still
  present as ammonia molecules. Only about 1 has
  actually produced hydroxide ions.

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pH Scale

• The pH scale is a way of expressing the strength
  of acids and bases. Instead of using very small
  numbers, we just use the NEGATIVE power of 10 on
  the Molarity (concentration) of the H or OH-
  ion.
• Acidic pH lt 7
• Neutral pH 7
• Basic pH gt 7

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pH of Common Substances
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Calculating pH

• pH -log H
• (Remember that the mean Molarity)
• Example If H 1 X 10-10pH - log 1 X
  10-10
• pH - (- 10)
• pH 10
• Example If H 1.8 X 10-5pH - log 1.8 X
  10-5
• pH - (- 4.74)
• pH 4.74

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pOH

• Since acids and bases are opposites, pH and pOH
  are opposites
• pOH does not really exist, but it is useful for
  changing bases to pH.
• pOH looks at the perspective of a base
• pOH - log OH-
• Since pH and pOH are on opposite ends
• pH pOH 14

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pH
H
OH-
pOH
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pH Testing

• There are several ways to test pH
• Blue litmus paper (red acid)
• Red litmus paper (blue basic)
• pH paper (multi-colored)
• pH meter (7 is neutral, lt7 acid, gt7 base)
• Universal indicator (multi-colored)
• Indicators like phenolphthalein
• Natural indicators like red cabbage, radishes

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pH Indicators

• Indicators are dyes that can be added that will
  change color in the presence of an acid or base.
• Some indicators only work in a specific range of
  pH
• Once the drops are added, the sample is ruined
• Some dyes are natural, like radish skin or red
  cabbage

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Titrations

• Suppose you want to determine the molar
  concentration of an HCl solution.
• You place a known volume in a flask and add a
  base indicator (like phenolphthalein)
• You then add small amounts of a standardized base
  (like NaOH) with a buret
• You keep adding base until the solution turns the
  faintest shade of pink
• This is called the endpoint
• Using the balanced equation and the amounts of
  acid and base used, you can calculate molar
  concentration.
• You can titrate a base with a standard acid
  solution in the same way.

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Titration
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35.62 mL of NaOH is neutralized with 25.2 mL of
0.0998 M HCl by titration to an equivalence
point. What is the concentration of the NaOH?
Ma Va Mb Vb Ma Va Mb Vb (0.0998
M) (25.2 mL)
0.0706 M (35.62 mL)
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Preparing solutions by dilution

• If you want to dilute a solution, use the
  following formula
• M1V1 M2V2

• Example
• You have a stock bottle of hydrochloric acid,
  which is 12.1 M. You need 400 mL of 0.10 M HCl.
  How much of the acid and how much water will you
  need?
• 3.3 mL of HCl and 396.7 mL of water

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Buffers Controlling pH

• Buffers, or buffer solutions, resist a change in
  pH caused by the addition of acids or bases.
• There are two types of buffers
• Mixtures of weak acids and bases these may be
  conjugate acid-base pairs, or nonconjugate
  acid-base pairs
• Amphoteric species these are substances that
  can act either as an acid or a base, like water

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Antacids Good, Basic Chemistry

• The stomach secretes hydrochloric acid to
  activate enzymes that break down proteins.
  Sometimes the stomach produces too much acid and
  it can work its way back up the esophagus leading
  to heartburn.
• Antacids are compounds that neutralize the excess
  acid
• Bicarbonates NaHCO3 and KHCO3
• Carbonates CaCO3 and MgCO3
• Hydroxides Al(OH)3 and Mg(OH)2

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