CH 4: Chemical Reactions

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CH 4 Chemical Reactions

• Renee Y. Becker
• Valencia Community College
• CHM 1045

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Solutions

• Solute solid in liquid or lowest mass quantity
  of substance
• Solvent- liquid solute is dissolved in or highest
  mass quantity of substance

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Solution Concentrations

• Concentration allows us to measure out a
  specific number of moles of a compound by
  measuring the mass or volume of a solution.
• Molarity(M) Moles of Solute
• Liters of Solution
• moles ML L moles/M

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Example 1 Solution Concentrations

• How many moles of solute are present in 125 mL of
  0.20 M NaHCO3?

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Example 2 Solution Concentrations

• How many grams of solute would you use to prepare
  500.00 mL of 1.25 M NaOH?

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Solution Concentrations

• Dilution the process of reducing a solutions
  concentration by adding more solvent.
• Moles of solute(constant) Molarity x Volume
• Mi Vi Mf Vf Vf (Mi Vi) / Mf
• Mf (Mi Vi) / Vf

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Example 3 Solution Concentrations

• What volume of 18.0 M H2SO4 is required to
  prepare 250.0 mL of 0.500 M H2SO4?

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Example 4 Solution Concentrations

• What is the final concentration if 75.0 mL of
• 3.50 M glucose is diluted to a volume of 400.0
  mL?

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Solution Stoichiometry

• Titration a technique for determining the
  concentration of a solution
• Standard solution known concentration
• If you have a known volume of standard solution
  and use it to titrate a known volume of an
  unknown concentrated solution you can calculate
  to find the number of moles in the unknown and
  therefore find its concentration

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Titration

• When doing a titration you add titrant (standard
  solution) to the analyte (unknown concentration
  solution) until the endpoint or the equivalence
  point is reached. This point is when you have
  equal moles of titrant and analyte, from the
  volume of the titrant and analyte used and the
  molarity of the titrant, you can find the
  molarity of the analyte
• Endpoint- based on an indicator
• Indicator- a substance that changes color in a
  specific pH range
• Equivalence point- not based on an indicator,
  usually a pH meter
• Use Manalyte Vanalyte Mtitrant Vtitrant

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Example 5 Solution Stoichiometry

• A 25.0 mL sample of vinegar (dilute CH3CO2H) is
  titrated and found to react with 94.7 mL of a
  0.200 M NaOH. What is the molarity of the acetic
  acid solution?
• NaOH(aq) CH3CO2H(aq) ? CH3CO2Na(aq) H2O(l)

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OxidationReduction Reactions

• Assigning Oxidation Numbers All atoms have an
  oxidation number regardless of whether it
  carries an ionic charge.
• 1. An atom in its elemental state has an
  oxidation number of zero.

Elemental state as indicated by single elements
with no charge. Exception diatomics H2 N2 O2
F2 Cl2 Br2 and I2
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OxidationReduction Reactions

• 2. An atom in a monatomic ion has an oxidation
  number identical to its charge.

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OxidationReduction Reactions

• 3. An atom in a polyatomic ion or in a molecular
  compound usually has the same oxidation number it
  would have if it were a monatomic ion.
• A. Hydrogen can be either 1 or 1.
• B. Oxygen usually has an oxidation number of 2.
• In peroxides, oxygen is 1.
• C. Halogens usually have an oxidation number of
  1.
• When bonded to oxygen, chlorine, bromine, and
  iodine have positive oxidation numbers.

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OxidationReduction Reactions

• 4. The sum of the oxidation numbers must be zero
  for a neutral compound and must be equal to the
  net charge for a polyatomic ion.
• A. H2SO4 neutral atom, no net charge
• SO42- sulfate polyatomic ion
• SO42- Sx O42- -2
• X -8 -2
• X 6 so sulfur has an oxidation of 6

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OxidationReduction Reactions

• B. ClO4 , net charge of -1
• ClO4-1 Clx O42- -1
• X -8 -1
• X 7 so the oxidation number of chloride is 7

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Example 6 OxidationReduction Reactions

• Assign oxidation numbers to each atom in the
  following
• A. CdS F. VOCl3
• B. AlH3 G. HNO3
• C. Na2Cr2O7 H. FeSO4
• D. SnCl4 I. Fe2O3
• E. MnO4 J. V2O3

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Electrolytes in Solution

• Electrolytes Dissolve in water to produce ionic
  solutions.
• Nonelectrolytes Do not form ions when they
  dissolve in water.

a) NaCl soln conducts electricity, completes
circuit (charged particles)
b) C6H12O6 does not
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Electrolytes in Solution

• Dissociation
• The process by which a compound splits up to
  form ions in the solution.

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Electrolytes in Solution

• Strong Electrolyte Total dissociation when
  dissolved in water.
• Weak Electrolyte Partial dissociation when
  dissolved in water.

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Types of Reactions

1. Precipitation
2. Acid-base neutralization
3. Oxidation-reduction (redox)
4. Double replacement
5. Single replacement
6. Combination
7. Decomposition

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Types of Chemical Reactions

• Precipitation Reactions A process in which an
  insoluble solid precipitate drops out of the
  solution.
• Most precipitation reactions occur when the
  anions and cations of two ionic compounds change
  partners. (double replacement)
• Pb(NO3)2(aq) 2 KI(aq) ? 2 KNO3(aq) PbI2(s)

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Solubility Rules Precipitation

• Allow you to predict whether a reactant or a
  product is a precipitate.
• Soluble compounds are those which dissolve to
  more than 0.01 M.
• There are three basic classes of salts

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Solubility Rules Precipitation

• 1. Salts which are always soluble
• All alkali metal salts Cs, Rb, K, Na, Li
• All ammonium ion (NH4) salts
• All salts of the NO3, ClO3, ClO4, C2H3O2, and
  HCO3 ions

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Solubility Rules Precipitation

• 2. Salts which are soluble with exceptions
• Cl, Br, I ion salts except with Ag, Pb2,
  Hg22
• SO42 ion salts except with Ag, Pb2, Hg22,
  Ca2, Sr2, Ba2

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Solubility Rules Precipitation

• 3. Salts which are insoluble with exceptions
• O2 OH ion salts except with the alkali metal
  ions, and Ca2, Sr2, Ba2 ions
• CO32, PO43, S2, CrO42, SO32 ion salts
  except with the alkali metal ions and the
  ammonium ion
• If not listed the compound is probably insoluble

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Example 7 Solubility Rules Precipitation

• Predict the solubility of the following in water
• CdCO3
• MgO
• Na2S
• (d) PbSO4
• (e) (NH4)3PO4

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Example 8 Solubility Rules Precipitation

• Write the balanced reaction and predict whether a
  precipitate will form for
• NiCl2 (aq) (NH4)2S (aq) ?
• (b) Na2CrO4 (aq) Pb(NO3)2 (aq) ?
• (c) AgClO4 (aq) CaBr2 (aq) ?

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Equations

• Molecular equation Balanced reaction
• 2 FeBr3(aq) 3 Pb(NO3)2(aq) ? 2 Fe(NO3)3(aq)
  3 PbBr2(s)
• Complete ionic equation All broken up into ions
  (only aqueous solutions)
• 2 Fe3(aq) 6 Br-(aq) 3 Pb2(aq) 6
  NO3-(aq) ? 2 Fe3(aq) 6 NO3-(aq) 3
  PbBr2(s)
• Net ionic equation Cancel out spectator ions
• 3 Pb2(aq) 6 Br-(aq) ? 3 PbBr2(s)

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Net Ionic Equations for Precipitation Reactions

• Write net ionic equation for the following
  reaction
• 2 AgNO3(aq) Na2CrO4(aq) ? Ag2CrO4(s) 2
  NaNO3(aq)
• Is it balanced? If not do it! (molecular
  equation)
• Separate all aqueous soln into ions (complete
  ionic equation)
• Cancel out spectator ions on both sides
• Rewrite (net ionic equation)

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Example 9

• Write the ME, CIE, and NIE for the following
  reaction
• Na2CrO4 (aq) Pb(NO3)2 (aq) ? NaNO3(aq)
  PbCrO4(s)

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Types of Chemical Reactions

• AcidBase Neutralization A process in which an
  acid reacts with a base to yield water plus an
  ionic compound called a salt.
• The driving force of this reaction is the
  formation of the stable water molecule.
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)

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AcidBase Concepts

• Arrhenius Acid
• A substance which dissociates in water to form
  hydrogen ions (H).
• Arrhenius Base
• A substance that dissociates in, or reacts with,
  water to form hydroxide ions (OH).
• Limitations Has to be an aqueous solution and
  doesnt account for the basicity of substances
  like NH3.

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AcidBase Concepts

• Brønsted Acid Can donate protons (H) to another
  substance.
• Brønsted Base Can accept protons (H) from
  another substance. (NH3)

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Example 10 Conjugate acid-base pairs

• For the following reactions label the acid, base,
  conjugate acid, and conjugate base.
• CH3CO2H(aq) H2O(l) ? H3O(aq)
  CH3CO2-(aq)
• NH3(aq) H2O(l) ? NH4(aq) OH-(aq)

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AcidBase Concepts

• Lewis Acid Electron pair acceptor. Al3, H,
  BF3.
• Lewis Base Electron pair donor. H2O, NH3, O2.
• Bond formed is called a coordinate bond or
  dative bond.

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Example 11

• Which of the following is a Bronsted-Lowry base
  but not an Arrhenius base?
• NaOH
• NH3
• Mg(OH)2
• KOH

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Acids and Bases

• Strong acid - st. electrolyte, almost completely
  dissociates in water
• HCl, H2SO4, HNO3, HClO4, HI, HBr
• Weak acid - wk. electrolyte, does not dissociate
  well in water
• HF, HCN, CH3CO2H
• Strong base - st. electrolyte, almost completely
  dissociates in water
• Metal hydroxides
• Weak base - does not dissociate well in water

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AcidBase Concepts
Other Weak bases trimethyl ammonia N(CH3)3,
C5H5N pyridine, ammonium hydroxide NH4OH, H2O
water
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ME, CIE, NIE for Acids/Bases

• Strong Acid Strong Base
• ME HCl(aq) NaOH(aq) ? H2O(l) NaCl(aq)
• Complete Ionic Equation
• H Cl- Na OH- ? H2O(l) Na Cl-
• Net Ionic Equation
• H OH- ? H2O(l)
• or
• H3O OH- ? 2 H2O(l)

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ME, CIE, NIE for Acids/Bases

• Weak Acid Strong Base
• ME HF(aq) NaOH(aq) ? H2O(l) NaF(aq)
• Complete Ionic Equation
• HF Na OH- ? H2O(l) Na F-
• Net Ionic Equation
• HF OH- ? H2O(l) F-

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Example 12 ME, CIE, NIE for Acids/Bases

• Write ME, CIE and NIE for the following
• NaOH(aq) CH3CO2H(aq) ?
• (b) HCl(aq) NH3(aq) ?
• NaOH strong base will dissociate well
• CH3CO2H weak acid doesnt dissociate well
• HCl is a strong acid and therefore a strong
  electrolyte
• NH3 is a weak base and is a weak electrolyte

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Types of Chemical Reactions

• Double Replacement These are reactions where
  two reactants just exchange parts. (double
  displacement)
• AX BY ? AY BX
• BaCl2(aq) K2SO4(aq) ? BaSO4(s) 2 KCl(aq)
• This is also a ppt reaction, if I ask you what
  type of reaction is it, what is the best answer??

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Types of Chemical Reactions

• OxidationReduction (Redox) Reaction A process
  in which one or more electrons are transferred
  between reaction partners.
• The driving force of this reaction is the
  decrease in electrical potential.
• Mg(s) I2(g) ? MgI2(s)
• Oxidation Mg0 ? Mg2 2 electrons
• Reduction I20 2 electrons ? I21-

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Example 12

• Which of the following is not an acid-base
  neutralization reaction?
• HCl(aq) NaOH(s) ? NaCl(aq) H2O(l)
• 2 HF(aq) Mg(OH)2(aq) ? MgF2(aq) 2 H2O(l)
• Pb(NO3)2(aq) 2 KI(aq) ? PbI2 (s) 2
  KNO3(aq)

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OxidationReduction Reactions

• Redox reactions are those involving the oxidation
  and reduction of species.
• Oxidation and reduction must occur together. They
  cannot exist alone.
• Fe2 Cu0 ? Fe0 Cu2
• Reduced Iron gained 2 electrons Fe2 2
  e ? Fe0
• Oxidized Copper lost 2 electrons Cu0 ?
  Cu2 2e
• Remember that electrons are negative so if you
  gain electrons your oxidation decreases and if
  you lose electrons your oxidation increases

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OxidationReduction Reactions

• Fe2 Cu0 ? Fe0 Cu2
• Fe2 gains electrons, is reduced, and we call it
  an oxidizing agent
• Oxidizing agent is a species that can gain
  electrons and this facilitates in the oxidation
  of another species. (electron deficient)
• Cu0 loses electrons, is oxidized, and we call it
  a reducing agent
• Reducing agent is a species that can lose
  electrons and this facilitates in the reduction
  of another species. (electron rich)

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Example 13

• Which is a reduction half reaction?
• Fe ? Fe2 2e
• Fe2 ? Fe3 1e
• Fe ? Fe3 3e
• Fe3 1e ? Fe2

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Example 14 OxidationReduction Reactions

• For each of the following, identify which species
  is the reducing agent and which is the oxidizing
  agent.
• Ca(s) 2 H(aq) ? Ca2(aq) H2(g)
• 2 Fe2(aq) Cl2(aq) ? 2 Fe3(aq) 2 Cl(aq)
• C) SnO2(s) 2 C(s) ? Sn(s) 2 CO(g)

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Balancing Redox Reactions

• Half-Reaction Method Allows you to focus on the
  transfer of electrons. This is important when
  considering batteries and other aspects of
  electrochemistry.
• The key to this method is to realize that the
  overall reaction can be broken into two parts, or
  half-reactions. (oxidation half and reduction
  half)

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Balancing Redox Reactions

• Balance for an acidic solution
• MnO4(aq) Br(aq) ? Mn2(aq) Br2(aq)
• 1. Determine oxidation and reduction
  half-reactions
• Oxidation half-reaction Br(aq) ? Br20(aq)
• Reduction half-reaction MnO4(aq) ? Mn2(aq)
• 2. Balance for atoms other than H and O
• Oxidation 2 Br(aq) ? Br2(aq)
• Reduction MnO4(aq) ? Mn2(aq)

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Balancing Redox Reactions

• 3. Balance for oxygen by adding H2O to the side
  with less oxygen
• Oxidation 2 Br(aq) ? Br2(aq)
• Reduction MnO4(aq) ? Mn2(aq) 4 H2O(l)
• 4. Balance for hydrogen by adding H to the side
  with less hydrogens
• Oxidation 2 Br(aq) ? Br2(aq)
• Reduction MnO4(aq) 8 H(aq) ? Mn2(aq) 4
  H2O(l)

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Balancing Redox Reactions

• 5. Balance for charge by adding electrons (e)
• Oxidation 2 Br(aq) ? Br2(aq) 2 e
• Reduction MnO4(aq) 8 H(aq) 5 e ?
  Mn2(aq) 4 H2O(l)
• 6. Balance for numbers of electrons by
  multiplying
• Oxidation 52 Br(aq) ? Br2(aq) 2 e
• Reduction 2MnO4(aq) 8 H(aq) 5 e ?
  Mn2(aq) 4 H2O(l)

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Balancing Redox Reactions

• 7. Combine and cancel to form one equation
• Oxidation 10 Br(aq) ? 5 Br2(aq) 10 e
• Reduction 2 MnO4(aq) 16 H(aq) 10 e ?2
  Mn2(aq) 8 H2O(l)
• 2 MnO4(aq) 10 Br(aq) 16 H(aq) ?2 Mn2(aq)
  5 Br2(aq) 8 H2O(l)
• We will not be balancing in basic solutions!!
  (until CHM 1046)

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Example 15 Balancing Redox Reactions

• Balance the following in an acidic soln
• NO3(aq) Cu(s) ? NO(g) Cu2 (aq)

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Types of Reactions

• A single-replacement reaction is a a reaction
  where a more active metal displaces another, less
  active metal in a compound.
• If a metal precedes another in the activity
  series, it will undergo a single-replacement
  reaction
• Fe(s) CuSO4(aq) ? FeSO4(aq) Cu(s)
• FeSO4(aq) Cu(s) ? NR

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Activity Series

• Metals that are most reactive appear first in the
  activity series.
• Metals that are least reactive appear last in the
  activity series.
• The relative activity series is
• Li gt K gt Ba gt Sr gt Ca gt Na gt Mg gt
• Al gt Mn gt Zn gt Fe gt Cd gt Co gt Ni gt
• Sn gt Pb gt (H) gt Cu gt Ag gt Hg gt Au

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Types of Reactions

• A combination reaction is a reaction where two
  simpler substances are combined into a more
  complex compound.

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Types of Reactions

• In a decomposition reaction, a single compound is
  broken down into simpler substances.
• Heat or light is usually required to start a
  decomposition reaction. Ionic compounds
  containing oxygen often decompose into a metal
  and oxygen gas.
• 2 HgO(s) ? 2 Hg(l) O2(g)