Unit 4: Aqueous Reactions and Solution Chemistry

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Unit 4 Aqueous Reactions and Solution Chemistry

• By Ms. Buroker

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General Properties of Aqueous Solutions

• A solution is made of a solute and solvent.
• Solute Dissolved Material
• Solvent Dissolving Material (Usually found in
  the greatest amount)

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Aqueous solutions

• Dissolved in water.
• Water is a good solvent because the molecules
  are polar.
• 1.)The oxygen atoms have a partial negative
  charge.
• 2.)The hydrogen atoms have a partial
  positive charge.
• 3.)The angle is 105º.

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Solvation

• This is known as HYDRATION when the solvent is
  water.
• The process of breaking the ions of salts apart.
  
• Ions have charges and attract the opposite
  charges on the water molecules.

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Solubility

• How much of a substance will dissolve in a given
  amount of water.
• Varies greatly, but if they do dissolve the ions
  are separated, and they can move around.
• The ionic solid dissociates into its component
  ions as it dissolves.
• Water can also dissolve non-ionic compounds if
  they have polar bonds.

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General Properties of Aqueous Solutions

• Electrolytes and Nonelectrolytes
• Aqueous solution Aqueous solution
• containing ions containing no ions

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Electrolytes

• Strong Electrolytes
• Compounds which dissolve almost completely in
  solution (exist as ions).
• Weak Electrolytes
• Compounds which do not dissolve or ionize in
  solution molecular compounds mostly.

Strong electrolyte Weak electrolyte Nonelectrolyte
Ionic All None None
Molecular Strong Acids Weak Acids All other compounds
Weak Bases
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Acidic Solutions

• Acids-form H ions when dissolved.
• Strong acids disassociate completely.
• You Must Memorize These!!
• H2SO4, HNO3, HCl, HBr, HI, HClO4, HCLO3
• Weak acids do NOT dissociate completely.
• Write the dissociation equation for
• Hydrochloric acid
• Acetic Acid

HCl ? H Cl-
HC2H3O2 ? H C2H3O2-
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Common Reactions

• Synthesis
• A B ? AB
• Decomposition
• AB ? A B
• Single Replacement
• A BX ? AX B or
• A XB ? XA B
• Double Replacement
• AX BY ? AY BX

Dont forget the metal reactivity series and the
non-metal reactivity series!!
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Precipitation Reactions

• When aqueous solutions of ionic compounds are
  poured together a solid forms.
• A solid that forms from mixed solutions is a
  precipitate

Watch this video for examples of precipitation
reactions
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Precipitation reaction

• Called METATHESISGreek word for to transpose.
• We can predict the products
• Can only be certain by experimenting
• The anion and cation switch partners
• AgNO3(aq) KCl2(aq) ?
• Zn(NO3)2(aq) BaCr2O7(aq) ?
• CdCl2(aq) Na2S(aq) ?

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Precipitations Reactions

• Only happen if one of the products is insoluble,
  otherwise all the ions stay in solution-nothing
  has happened.
• Three things drive a double replacement reaction
  the formation of a gas, a solid, or a liquid.

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Solubility Rules

• Molecular Equation
• Complete Ionic Equation
• Net Ionic Equation
• Minus Spectator Ions

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Acid- Base Reactions

• Acids
• Substances that are able to ionize in aqueous
  solutions to form a hydrogen ion and thereby
  increase the concentration of H(aq) ions.
• Bases
• Substances that accept (react with H) ions.
  Substances that increase the OH- concentration
  when added to water.

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Strong and Weak Acids and Bases

• Acids and bases which completely ionize in
  solution are considered strong acids and strong
  bases (strong electrolytes)! The opposite is
  true concerning those which weakly ionize they
  are weak acids and weak bases.

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Strong and Weak Acids and Bases

• Strong Acids
• Hydrochloric, HCl
• Hydrobromic, HBr
• Hydroiodic, HI
• Chloric, HClO3
• Perchloric, HClO4
• Nitric, HNO3
• Sulfuric, H2SO4
• (Most acids are weak)

• Strong Bases
• Group 1A metal hydroxides
• Heavy group 2A metal hydroxides
• (Ca, Sr, Ba)
• (most common weak base, NH3)

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Lets Practice!

• Classify each of the following dissolved
  substances as a strong electrolyte, weak
  electrolyte, or nonelectrolyte
• CaCl2, HNO3, C2H5OH (ethanol), HCHO2 (formic
  acid), KOH
• Answer
• CaCl2 strong , HNO3 strong, C2H5OH (ethanol)
  non, HCHO2 weak (formic acid), KOH strong

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Neutralization Reactions

• Reactions which take place between an acid and a
  metal hydroxide they produce water and a salt.
• Write a balanced complete chemical equation for
  the reaction between aqueous solution of acetic
  acid and barium hydroxide. Write the ionic
  equation for this reaction.

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Acid-Base Reactions With Gas Formation

• The sulfide ion, the carbonate ion, and the
  bicarbonate ion are bases which react with acids
  to form gases that have low solubility in water.
• H2S forms when
• HCl(aq) Na2S(aq) ? H2S(g) NaCl (aq)
• Carbonates Bicarbonates react to form CO2(g)

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• CO32- and HCO3- react with acids first to give
  carbonic acid, then carbonic acid decomposes into
  water and CO2(g).
• HCl(aq) NaHCO3(aq) ? NaCl(aq) H2CO 3(aq)
• Then
• H2CO3(aq) ? H2O(l) CO2(g)
• Actually .
• HCl(aq) NaHCO3(aq) ? NaCl(aq) H2O(l) CO2(g)

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Oxidation- Reduction Reactions

• When an atom or molecule has become more
  positively charged that is, it has lost
  electrons we say that it has been oxidized.
• Oxidation is Loss
• When an atom or molecule has become more
  negatively charged that is, it has gained
  electrons we say that it has been reduced.
• Reduction is Gain
• OIL RIG

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Example of a Non-Redox

• Lets look at a precipitation reaction
• AgCl Na2SO4 ? Ag2SO4 NaCl
• Notice that the charges of each of the ions are
  the same on both sides
• Ag1 ? Ag1
• Na1 ? Na1
• SO4-2 ? SO4-2
• Cl-1 ? Cl-1

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Example of a Redox

• Look at this synthesis reaction
• Fe O2 ? Fe2O3
• What happens to the charge on each of the
  elements?
• Fe0 ? Fe3
• O20 ? O-2

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Which reactions qualify as redox?

• Redox reactions reactions in which the charge of
  elements change from one side to the other
• The element that is Oxidized loses electrons,
  which means a more positive charge (the iron in
  the previous example)
• The element that is Reduced gains electrons,
  which means a more negative charge (the oxygen in
  the previous reaction)

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• When a metal undergoes corrosion, it loses
  electrons and forms cations. For example
• Ca(s) 2H(aq) ? Ca2 H 2(g)

• Corrosion at the terminal of a battery, caused by
  attack of the metal by sulfuric acid from the
  battery.

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Oxidation Numbers

• Oxidation Numbers are basically the
  hypothetical charge we assign to atoms in a
  reaction to keep track of the electrons.
• Rules
• 1.) For an atom in its elemental form, the
  oxidation number is zero.
• 2.) For any monatomic ion, the oxidation number
  equals the charge on the ion.
• 3.)Nonmetals usually have negative oxidation
  numbers, although they can sometimes be positive.
• (The oxidation number of hydrogen is 1 when
  bonded to nonmetals and -1 when bonded to metals.
• 4.) The sum of the oxidation numbers of all atoms
  in a neutral compound is zero. The sum of the
  oxidation numbers in a polyatomic ion equals the
  charge on the ion.

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Examples of Determining Oxidation Numbers

• What are the oxidation numbers of each element in
  KNO3?
• O -2
• K 1
• Since the oxidation numbers of a neutral compound
  must equal 0
• (1) (N) (3-2) 0
• N 5

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Examples of Determining Oxidation Numbers
(continued)

• What are the oxidation numbers of each element in
  (NH4)1?
• H 1
• Since the oxidation numbers of a polyatomic ion
  must equal its charge
• (N) (41) 1
• N -3

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Do these on your own

• Determine the oxidation numbers for each element
  in the following chemicals
• H2CO3
• K3PO4
• Al(NO3)3

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Answers

• H2CO3
• H1
• O-2
• C4

• K3PO4
• K1
• P5
• O-2

• Al(NO3)3
• Al3
• N5
• O-2

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Some Examples

• Assign Oxidation Numbers to the elements in
  the following equations
• Fe(s) Ni(NO3)2(aq) ? Fe(NO3)2(aq) Ni (s)
• Write the balanced molecular and net ionic
  equation for the reaction of aluminum with
  hydrobromic acid. Which element is oxidized in
  the reaction?

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A couple more terms

• From the previous example, we know
• Al H2O ? Al2O3 H2
• Al0 ?Al3
• H1?H0
• The aluminum is oxidized
• The hydrogen is reduced
• H2O is called the oxidizing agent (the e-
  acceptor)
• Al is called the reducing agent (the e- donor)

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Balancing Redox Reactions

• Now that we know what they are, how do we balance
  them?
• Will it be the same as before?
• Of course not.
• Are we excited to learn more techniques that will
  improve our Chemistry knowledge?
• Of course we are.

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Where do redox reactions occur?

• Redox reactions typically take place in solutions
• Because of this redox reactions are most often
  written as net ionic equations
• If the reaction takes place in only water, you
  can balance by inspection (just like normal)
• However, if the reactions involve acidic or basic
  solutions balancing is different

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Why would a redox reaction take place in an acid
or base?

• First, lets define
• Acid H ion donator when dissolved in water
• Base H ion acceptor when dissolved in water
• When these types of solutions are present, it
  creates an opportunity for charges to get messed
  with
• Extra H ions will attract electrons
• Removing H ions creates leftover OH- ions in the
  water

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Writing Half Reactions

• In any case, when balancing a redox reaction it
  will be handy to write the reaction as two half
  reactions
• One showing only the oxidation
• One showing only the reduction
• We will treat the two half reactions separately,
  and then combine them at the end

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Writing Half Reactions Examples

• Write the half reactions for this redox
• CeCl4 Sn(NO3)2 ? SnCl4 Ce(NO3)3
• The half reactions would look like this
• Ce4 ? Ce3 (reduction)
• Sn2 ? Sn4 (oxidation)
• Notice the cancellation of the spectator ions,
  nitrate and chloride (they are net ionic
  equations)

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To balance a redox reaction in an acidic solution

• Write the reaction as two half reactions
• For each half reaction
• Balance the elements except H and O
• Balance the oxygen by adding water
• Balance the hydrogen by adding H ions
• Balance the charge by adding electrons (e-)
• Multiply the half reactions to equalize the
  number of electrons in each
• Combine the half reactions and cancel anything
  identical (which should always include the
  electrons)

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To balance a redox reaction in an acidic
solution Example

• Balance the following redox reaction that occurs
  in an acidic solution
• MnO4-1 Fe2 ? Fe3 Mn2
• Write the reaction as two half reactions
• Oxidation Fe2 ? Fe3
• Reduction MnO4-1 ? Mn2

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To balance a redox reaction in an acidic
solution Example

• Fe2 ? Fe3 MnO4-1 ? Mn2
• For each half reaction
• Balance the elements except H and O
• Not needed in these half reactions
• Balance the oxygen by adding water
• MnO4-1 ? Mn2 4H2O
• Balance the hydrogen by adding H ions
• 8H MnO4-1 ? Mn2 4H2O
• Balance the charge by adding electrons (e-)
• Fe2 ? Fe3 e-
• 5 e- 8H MnO4-1 ? Mn2 4H2O

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To balance a redox reaction in an acidic
solution Example

• Fe2 ? Fe3 e-
• 5 e- 8H MnO4-1 ? Mn2 4H2O
• Multiply the half reactions to equalize the
  number of electrons in each
• 5(Fe2 ? Fe3 e- )
• 5 e- 8H MnO4-1 ? Mn2 4H2O

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To balance a redox reaction in an acidic
solution Example

• 5(Fe2 ? Fe3 e- )
• 5 e- 8H MnO4-1 ? Mn2 4H2O
• Combine the half reactions and cancel anything
  identical (which should always include the
  electrons)
• 5 e- 8H MnO4-1 5Fe2 ? Mn2 4H2O 5Fe3
  5e-

8H MnO4-1 5Fe2 ? Mn2 4H2O 5Fe3
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To balance a redox reaction in an basic solution

• Follow the steps for an acidic solution, then
• To both sides of the reaction, add OH- ions to
  equal the number of H ions remaining
• Form water on the side containing both H and
  OH-, and then cancel water from both sides if
  possible

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To balance a redox reaction in an basic solution
Example

• Balance this redox reaction that takes place in a
  basic solution
• Ag CN-1 O2 ? Ag(CN)2-1 H2O
• Oxidation Ag CN-1 ? Ag(CN)2-1
• Reduction O2 ? H2O

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To balance a redox reaction in an basic solution
Example

• Ag CN-1 ? Ag(CN)2-1 O2 ? H2O
• Ag 2CN-1 ? Ag(CN)2-1 O2 ? H2O
• Ag 2CN-1 ? Ag(CN)2-1 O2 ? 2H2O
• Ag 2CN-1 ? Ag(CN)2-1 4H O2 ? 2H2O

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To balance a redox reaction in an basic solution
Example

• Ag 2CN-1 ? Ag(CN)2-1 4H O2 ? 2H2O
• Ag 2CN-1 ? Ag(CN)2-1 e- 4e- 4H O2 ?
  2H2O
• 4(Ag 2CN-1 ? Ag(CN)2-1 e-) 4e- 4H O2 ?
  2H2O
• 4e- 4H O2 4Ag 8CN-1 ? 4Ag(CN)2-1 4e-
  2H2O

4H O2 4Ag 8CN-1 ? 4Ag(CN)2-1 2H2O
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To balance a redox reaction in an basic solution
Example
4H O2 4Ag 8CN-1 ? 4Ag(CN)2-1 2H2O

• To both sides of the reaction, add OH- ions to
  equal the number of H ions remaining
• 4OH- 4H O2 4Ag 8CN-1 ? 4Ag(CN)2-1 2H2O
  4OH-

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To balance a redox reaction in an basic solution
Example
4OH- 4H O2 4Ag 8CN-1 ? 4Ag(CN)2-1 2H2O
4OH-

• Form water on the side containing both H and
  OH-, and then cancel water from both sides if
  possible
• 4OH- 4H O2 4Ag 8CN-1 ? 4Ag(CN)2-1 2H2O
  4OH-
• 4H2O O2 4Ag 8CN-1 ? 4Ag(CN)2-1 2H2O
  4OH-
• 2H2O O2 4Ag 8CN-1 ? 4Ag(CN)2-1 4OH-

2H2O O2 4Ag 8CN-1 ? 4Ag(CN)2-1 4OH-
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Concentration of Solutions

• Molarity moles of solute
• Liters of solution
• Molality moles of solute
• Kilograms of solvent
• Mole Fraction moles of part
• Moles in total solution
• M1V1 M2V2

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• When expressing the concentration of
  electrolytes, its important to remember that
  ionic compounds break apart completely in
  solution.
• 1.0M NaCl solution 1.0M Na and 1.0M Cl- ions
  in solution

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Titrations

• The process of titration is when we use a
  standard solution (of known concentration) and
  react it with an solution of unknown
  concentration to determine the concentration of
  the solution.
• Equivalence Point The point at which
  stoichiometrically equivalent quantities are
  brought together.

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Lets do a little practice!