Acids, Bases, and Salts

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Acids, Bases, and Salts

• Chapter 14, 15

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Some Properties of Acids

• 1. The word acid comes from the Latin word acere,
  which means "sour." All acids taste sour.
• 2. In 1663, Robert Boyle wrote that acids would
  make a blue vegetable dye called "litmus" turn
  red.
• 3. Acids react with bases (they destroy the
  chemical properties of bases).
• 4. Acids conduct an electric current.
• 5. Upon chemically reacting with an active metal,
  acids will evolve hydrogen gas (H2).

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Some Properties of Bases

• 1. The word "base" has a more complex history and
  its name is not related to taste. All bases taste
  bitter.
• 2. Bases are substances which will restore the
  original blue color of litmus after having been
  reddened by an acid.
• 3. Bases destroy the chemical properties of acids
  (will react with acids)
• 4. Bases will conduct an electric current.
• 5. Bases feel slippery (soap, bleach) on your
  skin.

This is because they dissolve the fatty acids and
oils from your skin and this cuts down on the
friction between your fingers as you rub them
together.
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Some Properties of Salts

• 1. A salt is the combination of an anion (- ion)
  and a cation ( ion).
• 2. Salts are products of the reaction between
  acids and bases.
• 3. Solid salts are usually crystalline.
• 4. If a salt dissolves in water solution, it
  usually dissociates into the anions and cations
  that make up the salt (depends on Ksp)

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The Acid Base Theory

• The three main theories regarding acids and bases
  are
• 1. Arrhenius 3. Lewis
• 2. Brønsted-Lowry

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Arrhenius Theory late 1890s

• DEFINITIONS
• Acid - any substance which donates hydrogen ions
  (H) to water (produces hydronium ions, H3O)
• HA ? H A
• Base - any substance which produces hydroxide
  ions (OH) in water.
• XOH ? X OH
• When acids and bases react, they neutralize each
  other, forming water and a salt
• HA XOH ? H2O XA

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Problems with Arrhenius Theory

• The theory did not explain why ammonia (NH3) was
  a base.
• The theory only considers water as a solvent. We
  know that an acid added to benzene will not
  dissociate. Solvents are crucial to acid
  definition.
• The end result of mixing certain acids and bases
  can be a slightly acidic or basic solution.
  Arrhenius had no explanation for this phenomenon
  (degrees of acidity).

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Brønsted Lowry Theory Early 1920s

• Two chemists, independent of one another,
  proposed a new definition of an acid and a base
• An acid is a substance from which a proton can be
  removed (donates protons).
• A base is a substance that can remove a proton
  from an acid (proton acceptor).
• This definition does not require acids and bases
  to be in aqueous solutions.

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Reactions Based on Bronsted - Lowry

• Which are the acids and bases?
• HCl H2O ? H3O Cl
• HCl - this is an acid, because it has a proton
  available to be transferred (it can give a
  proton).
• H2O - this is a base, since it gets the proton
  that the acid lost (it has the capacity to accept
  a proton).

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Conjugate acid-base pairs

• Example HCl H2O ? H3O Cl
• Notice that each pair (HCl and Cl as well as H2O
  and H3O differ by one proton (H). These pairs
  are called conjugate pairs.
• Example HNO3 H2O ? H3O NO3
• The acids are HNO3 and H3O and the bases are H2O
  and NO3. What are the pairs?

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Bases and Conjugate Acid
Base Name Conjugate acid Name
C2H3O2- Acetate ion CH3COOH Acetic acid
NH3 Ammonia NH4 Ammonium
H2PO4- Dihydrogen phosphate ion H3PO4 Phosphoric acid
HSO4- Hydrogen sulfate ion H2SO4 Sulfuric acid
OH- Hydroxide ion H20 water
NO3- Nitrate ion HNO3 Nitric acid
H2O water H30 Hydronium ion
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Lewis Theory Early 1920s

• Remember drawing Lewis Dot Structures for ionic
  and covalent compounds?
• Lewis Theory focuses on the nature of electrons
  rather than proton transfer.
• DEFINITIONS
• An acid is an electron pair acceptor and a base
  as an electron pair donor.
• Lewis Theory is much more general and apply to
  reactions that do not involve hydrogen or
  hydrogen ions.

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Lewis acid-base reaction

• BF3 accepts an electron pair from ammonia
• A Lewis acid must have an empty orbital to accept
  an electron pair.
• A Lewis base must have a pair of unshared
  electrons that can be donated. Typical Lewis
  bases are OH-, H2O, NH3, Cl-, CN- due to lone
  pair electrons.

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Lewis AB reactions and Formation of Coordinate
Complexes

• The metal ion is a Lewis acid and the ligands
  coordinated to the ion are Lewis bases.

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Autoionization of Water
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Strong Acids and Bases

• Strong acids are those that ionized completely in
  water.
• The dissociation of a strong base also looks like
  the diagram at the right in that it dissociates
  into positive and negative ions.

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7 Strong Acids

• HNO3 - nitric acid HCl - hydrochloric
  acidHBr - hydrobromic acidHI - hydroiodic
  acid
• H2SO4- sulfuric acid
• HClO4 - perchloric acid HClO3 - chloric
  acid (wanna be)

• Strong acids are assumed to ionize completely
  (100)
• in water. They exist as H3O ions in water. This
  is known
• as the leveling effect. Water has a greater
  affinity for H
• than the conjugate bases do.

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ANIMATION LINKS

• Acid ionization equilibrium demo

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Weak Acids and Bases

• Some acids and bases ionize only slightly in
  water.
• These are considered weak.
• The most important weak base is ammonia.

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Balance of ions in solutions

• Acidic Neutral
• Solution Solution

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Strong Bases
Uncommon in labs because too expensive

• LiOH - lithium hydroxideNaOH - sodium
  hydroxideKOH - potassium hydroxideRbOH -
  rubidium hydroxideCsOH - cesium
  hydroxideBa(OH)2 - barium hydroxideSr(OH)2 -
  strontium hydroxideCa(OH)2 - calcium hydroxide

GROUP 1 hydroxides
Some GROUP 2 hydroxides

• Strong bases also ionize completely in water,
  except
• for Sr(OH)2 and Ca(OH)2 which are only slightly
  soluble
• (remember Mg(OH)2 is insoluble).

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Polyprotic Acids

• Polyprotic acids dissociate in a stepwise fashion
  with different Ka values for each step In the
  second and subsequent ionizations the acids are
  always weak, whether or not the original is a
  strong or weak acid.
• For most of these acids (ex. H3PO4), the first
  dissociation contributes the significant amount
  of H for pH calculations, and the rest are
  negligible (except for H2SO4 where second
  ionization is significant).

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Naming Acids -REVIEW

• -ide ending (elements) hydro____ic acid
• ex. chloride (HCl) hydrochloric acid
• -ate ending (polyatomics) ______ic acid
• ex. chlorate (HClO3) chloric acid
• -ite ending(polyatomics) ______ous acid
• ex. chlorite (HClO2) chlorous acid

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Net Ionic Equations -REVIEW

• For aqueous acid-base reactions reactions, it is
  common to write equations in the net ionic form.
• Standard form
• NaOH(aq) HCl(aq) ? NaCl(aq) H2O(l)
• Ionic form
• Na(aq) OH-(aq) H(aq) Cl-(aq) ? Na(aq)
  Cl-(aq) H2O(l)
• Net ionic form
• OH-(aq) H (aq) ? H2O(l)
• (No spectator ions are included)

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Things to remember when writing Net Ionic
Equations

• Binary Acids HCl, HBr, and HI are strong all
  other binary acids and HCN are weak. Strong acids
  are written in ionic form weak acids are written
  in molecular form.
• Ternary Acids If the number of oxygen atoms in
  an inorganic acid molecule exceeds the number of
  hydrogen atoms by two or more, the acid is strong
  (complete dissociation). We will consider all
  organic acids as weak.
• Strong HClO3, HClO4, H2SO4, HNO3
• Weak HClO, H3AsO4, H2CO3, H4SiO4, HNO2

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• Polyprotic Acids (acids that contain more than
  one ionizable hydrogen atom. Ex H2SO4, H3PO4,
  H2CO3).
• Bases Hydroxides of Group 1 and 2 elements
  (except Be(OH)2 and Mg(OH)2) are strong bases.
  All others including ammonia, hydroxlamine, and
  organic bases are weak.
• Salts Salts are written in ionic form if
  soluble, and in undissociated form if insoluble.
  Know the solubility rules.
• Oxides Oxides are always written in molecular or
  undissociated form (ex MgO).
• Gases Gases are always written in molecular form
  (ex SO2).

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Practice Net Ionic Equations

• 1. AgNO3 (aq) H2SO4 (aq) ?
• 2. H4SiO4 (aq) NaOH (aq) ?
• 3. HBr (aq) KOH (aq) ?
• 1. Ag HSO4- ? AgHSO4(s)
• 2. H4SiO4 OH- ? H3SiO4- H2O
• 3. H OH- ? H2O

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Weak acids and bases will have Ka or Kb values
less than one, but greater than water
dissociation, Kw
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Relationship between Ka and Kb

• Ka x Kb Kw
• For any acid and its conjugate base, this
  relationship can be used to determine Ka or Kb.
• Ex NH3 H2O ? NH4 OH-
• NH4 H2O ? NH3 H3O
• Kb(NH3)NH4OH- Ka(NH4)NH3H3O
• NH3
  NH4
• Therefore, Ka x Kb OH-H3O Kw

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Strength of Acid-base pairs

• Strong acids yield WEAK conjugate bases they
  have a low affinity for H
• Weak acids yield STRONG conjugate bases
• Strong bases yield WEAK conjugate acids
• Weak bases yield STRONG conjugate acids

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pH Scale
Soren Sorenson (1868-1939) invented the pH scale
while creating a way to test the acidity of beer.
Beer has a pH of about 4.5.

• The pH scale (potential hydrogen scale) is a
  measure of hydronium ion (H3O) concentration.
• Hydronium ion concentration indicates acidity.
  Each increase in pH means a 10-fold decrease in
  H.
• The higher the H3O, the higher the acidity.

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pH scale and H
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Calculating pH

• The concentration (M or mol/L) of H3O is
  expressed in powers of 10, from 10-14 to 100.
• Scientists use pH which is the negative log of
  H3O.
• pH -logH3O
• Note The significant figures for logarithmic
  numbers are given after the decimal, and the
  numbers preceding the decimal give the exponent.

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Calculating pH of a strong acid

• Ex Given a solution of 0.50M HCl, what is the
  pH?
• Step 1 Find H3O in mol/L
• 0.50mol/L 5.0 x 10-1 mol/L
• Step 2 Place value in equation and solve.
• pH -log5.0 x 10-1 0.30

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Practice pH Calculations

• Find pH of the following solutions if H3O is
• 1.00 x 10-3
• 6.59 x 10-6
• 9.47 x 10-10
• Find H3O if the pH is
• 6.678 3. 10.0
• 2.533 4. 2.56

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pOH

• You can calculate the pH of a solution if you
  know the concentration of hydronium ion. OH-
• If we use the ion product constant of water we
  can derive this equation
• pHpOH 1.00 x 10-14
• Working with this equation leads to
• pH pOH 14

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Calculating pH of a strong base

• Ex Find the pH of a solution with an NaOH of
  1.0 x 10-8.
• Step 1 Solve for H3O in equation
• H3O 10-14
• OH-
• Step 2 Place values in
• H3O 10-14 10-6 M
• 1.0 x 10-8

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• Step 3 Solve for pH by placing H3O in pH
  logH3O
• pH -log(1.0 x 10-6)
• pH 6.0

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Practice pH Calculations Using pOH

• Find the pH of the following solutions with OH
  of
• 1.00 x 10-4
• 2.64 x 10-13
• 5.67 x 10-2
• 3.45 x 10-11

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Calculating pH, pOH, H and OH-

• If one of the these values is known, all others
  can be found using the following relationships

pH pOH 14
pH
pOH
pOH-logOH-
OH-10-pOH
pH-logH
H10-pH
OH-
H
H OH- Kw
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Calculating pH of a weak acid

• Ex Find the pH of a 1.00 M HF solution (Ka7.2
  x10-4)
• HF(aq) ? H(aq) F-(aq)
• Ka HF-
• HF
• Use ICE box method
• HF(aq) ? H(aq) F-(aq)
• I 1.00M 0 0
• C -x x x
• E 1.00 -x x x

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• Ka HF- (x)(x) x2
• HF (1.00-x) (1.00)
• x2 Ka(1.00) (7.2 x10-4)(1.00)
• x 2.7 x10-2
• CHECK 5 rule is valid since
• 2.7 x10-2/1.00 2.7
• pH -logH 1.57

Try AP 2005 practice problem!!
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Acid-base properties of salts

• Salt hydrolysis the reaction in which a salt
  dissolved in water produces an acid or basic
  solution (opposite of a neutralization reaction)
• Ex1
• AlCl3 H2O ? Al(OH)3 3HCl
• A salt that contains the conjugate base of a
  strong acid will produce a slightly acidic
  solution when dissolved in water.

weak base
strong acid
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Acid-base properties of salts

• Ex2
• NaC2H3O2 H2O ? NaOH HC2H3O2
• A salt that contains the conjugate base of a weak
  acid will produce a slightly basic solution when
  dissolved in water.
• What will happen when NaCl dissolves in water?

strong base
weak acid
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Percent dissociation

• Percent dissociation amount dissociated x 100
• initial
  concentration

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Practice Problem find Ka given dissociation of
a weak acid

• Find H first using dissociation
• Use formula Kax2/acid0 - x to find Ka

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The effect of structure on acid-base properties

• Relative acidity of oxyacids (hydrogen is
  attached to oxygen, and acid contains one other
  element)
• Acid strength increases as the O-H bond is
  weakened (or with increasing of oxygen atoms on
  central atom).
• HClO4gtHClO3gtHClO2gtHClO
• Bond strength of H-F is too strong (F is so
  small) to make it a strong acid (will not
  dissociate easily). HIgtHBrgtHClgtgtHF

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BUFFERS

• A buffer solution is one which resists changes in
  pH when small quantities of an acid or a base are
  added to it.
• How do buffer solutions work?
• A buffer solution has to contain things which
  will remove any hydrogen ions or hydroxide ions
  that you might add to it - otherwise the pH will
  change.
• buffer demo

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ex HF and NaF ex NH3 and NH4Cl
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Calculating pH of a buffer

• Taking the log of the Ka equation (rearranged to
  solve for H)
• . yields the Henderson-Hasselbalch equation
• This is used to find the pH of a buffer solution

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H-H Special case

• What happens with equimolar amounts of acid and
  conjugate base ion?
• This situation simplifies to pH pKa, since log
  10.

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Preparing buffers

• A buffer can be prepared by adding a common ion
  (the conjugate base) to a weak acid.
• ex CH3COOH NaCH3COO
• A buffer can also be prepared by partially
  neutralizing a weak acid with a strong base to
  produce the conjugate base anion.
• ex CH3COOH OH- ? CH3COO- H2O

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Practice Problem

• A buffer problem is simply a weak acid
  equilibrium problem that involves a common ion..
• Ex Calculate the H, pH and percent
  dissociation of HF in a buffer solution that
  contains 1.0 M HF (Ka 7.2 x10-4) and 1.0 M NaF.
• ANS HF ? H F-
• Ka 7.2 x10-4 HF- x(1.0 x) x(1.0)
• HF 1.0 x
  1.0

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• Solving for x 7.2 x10-4
• H 7.2 x10-4 and pH 3.14
• dissociation is 7.2 x10-4 x100 0.072
• 1.0 M

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Buffer Capacity

• The buffering capacity of a solution represents
  the amount of H or OH- the buffer can absorb
  without significant changes in pH.
• The pH of a buffered solution is determined by
  the ratio A-/HA. The capacity of a buffered
  solution is determined by the magnitudes of HA
  and A-.
• Therefore, a good buffer solution will have
  relatively high concentrations of BOTH components.

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Practice Problem Buffer Preparation

• A chemist needs a solution buffered at pH 4.30
  and can choose from the following acids and their
  sodium salts
• a) CH2ClCOOH, cloroacetic acid (Ka 1.35x10-3)
• b) CH3CH2COOH, propanoic acid (Ka 1.3x10-5)
• c) C6H5COOH, benzoic acid (Ka 6.4x10-5)
• d) HOCl, hypochlorous acid (Ka 3.5x10-8)
• Calculate the ratio HA/A- required for each
  system to yield a pH of 4.30. Which will work
  best?
• A pH of 4.30 corresponds to H 5.0x10-5 M,
  and H KaHA/A-
• Using Ka values above, the HA/A- ratio for
  each of the salts are
• a) cloroacetic acid 3.7x10-2
• b) propanoic acid 3.8
• c) benzoic acid 0.78
• d) hypochlorous acid 1.4 x103

best since HA/A- ratio is closest to 1
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Indicators
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pH and titration curves
Adding strong acid to strong base
Adding strong base to strong acid
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Adding strong acid to weak base
Adding weak acid to strong base
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Adding weak base to strong acid
Adding strong base to weak acid
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More complicated titration curves
Adding weak acid to weak base
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Neutralization Reaction
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Acid-Base Titration