John A. Schreifels

1
CHAPTER 4

• Chemical Reactions

2
Overview

• Ions in Aqueous Solution
• Ion theory in solutions precipitation reactions
• Molecular and ionic equation
• Typical Reactions
• Precipitation
• Acid-Base
• Oxidation-Reduction (Balancing)
• Working with solutions
• Quantitative analysis

3
Ionic Thory of Solutions

• Pure water is a very poor conductor of
  electricity.
• Solutions from dissolving NaCl or KCl in water
  are very conductive.
• Solutions from dissolving substances such as
  sugar (sucrose) C12H22O11 are non-conductive.
• Electrolyte substance that produces ions when
  dissolved in water.
• Strong- good electrical conductor when dissolved
  in water (completely ionized). E.g. NaCl, KNO3,
  Mg(NO3)2, etc.
• Weak-poor conductor when dissolved in water
  (partial ionization)
• Non-electrolyte substance that does not produce
  ions when dissolved in water.

4
Aqueous Reactions and Net Ionic Equations

• Three forms for writing chemical reaction
• Molecular
• AgNO3(aq) NaCl(aq) ? AgCl(s) NaNO3(aq).
• Ionic
• Spectator ions are not directly involved in the
  reaction
• Net ionic exclude spectator ions
• Ag(aq) Cl?(aq) ? AgCl(s).

5
Precipitation Reactions

• Metathesis reaction reaction in which two
  substances react through exchange of their
  components. Driving force is often a
  precipitation.
• AX BY ? AY BX
• E.g. Predict if precipitation occurs for the
  mixture
• AgNO3(aq) KI(aq)
• NaClO4(aq) Pb(NO3)2(aq)
• Na2SO4(aq) BaCl2(aq),
• Ni(NO3)2(aq) (NH4)2S(aq)
• Hint Use the solubility rules to determine if
  either product is insoluble.

6
Solubility Rules
Rule Exception
Soluble
Group 1 elements, NH4
NO3 ?, ClO3?, ClO4 ?
Chlorides, bromides, iodides Ag, Pb2, Hg22
Acetates Ag, Hg22
Sulfates Sr2, Ba2, Pb2, Ca2
Insoluble Compounds
Carbonates, phosphates, oxalates, chromates, sulfides Group 1, NH4
Hydroxides, oxides Group 1, Ba2
7
Metathesis Reaction (cont.)

• Driving force is
• sometimes formation of weak or non electrolyte.
• E.g. acid base reactions
• CuO(s) 2HNO3(aq) ? Cu(NO3)2(aq) H2O(l)
• CuO is normally insoluble in water, but readily
  dissolves in aqueous nitric acid.
• sometimes formation of gas
• most common is CO2 from carbonates or H2S from
  sulfides
• E.g.
• CuCO3(s) 2HNO3(aq) ? CO2(g) Cu(NO3)2(aq)
  H2O(l)
• CuS(s) 2HNO3(aq) ? Cu(NO3)2(aq) H2S(g)

8
Acids, Bases, and Salts

• Arrhenius definition most often used
• Acid a hydrogen containing compound that
  releases hydrogen ions (H) in solution.
• HA(aq) H2O(l) ? H3O(aq) A?(aq) where
• HA HCl, HNO3,etc. and
• H3O hydronium ion often written as H.
• Base compound that releases hydroxide ions
  (OH?) in solution. The general reaction for a
  base is
• MOH(s) ? OH?(aq) M(aq) where
• M some metal such as Na, K, etc.
• Acids and bases can be strong or weak
  electrolytes.
• A base/acid that is a strong electrolyte is a
  strong base/acid.

9
Polyprotic acids and weak bases

• Some acids have more than one acidic proton.
• Sulfuric
• Phosphoric
• Most weak bases produce hydroxide ions by
  reaction with water.
• Ammonia

10
Strong and Weak Acids and Bases

• Organic acids are weak (usually have COOH).
• Amines (containing nitrogen) are weak.
• In water they are completely dissociated
• HCl(aq) H2O(l) ? H3O(aq) Cl?(aq)

Strong Acids Strong Bases
Chloric, HClO3 Grp 1A hydroxides (LiOH, NaOH, KOH, RbOH, CsOH)
Hydrobromic, HBr Grp 1A hydroxides (LiOH, NaOH, KOH, RbOH, CsOH)
Hydroiodic, HI Grp 2A metal hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2
Perchloric, HClO4 Grp 2A metal hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2
Sulfuric, H2SO4
Nitric, HNO3
11
Neutralization Reaction

• Acids react with bases to form a salt and
  possibly water (called the neutralization
  reaction)
• HA(aq)MOH(aq)?M(aq)A?(aq)H2O(l).
• If either the acid or base is a strong
  electrolyte, exclude spectator ions in the ionic
  form.
• E.g. HCN weak acid NaOH strong base
  neutralization reaction is
• HCN(aq) OH?(aq) ? CN?(aq) H2O(l)
• Eg. 2 HCl neutralized by NaOH net ionic
  equation
• H(aq) OH?(aq) ? H2O(l)

12
Oxidation Reduction

• Oxidation loss of at least one electron during
  a reaction..
• Ni(s) H(aq) ? Ni2(aq) H2(g)
• Reduction gain of at least one electron during
  a reaction.
• In above example, H gains an electron to become
  reduced.
• Every reaction must have an oxidation and
  reduction.
• Metals react with acids to form salts and
  hydrogen gas.
• Cu(s) 2HNO3(aq) ? Cu(NO3)2(aq) H2(g)
• Metals also oxidized with salts
• Fe(s) Cu(NO3)2(aq) ? Fe(NO3)2(aq) Cu(s)

13
Oxidation Number

• Oxidation number (state) the charge on an atom
  in a substance or monatomic ion.
• Rules
• Elemental form 0
• Monatomic ions charge of ion
• Oxygen ?2, except in H2O2 and other peroxides.
• Hydrogen 1, except with metal hydrides when it
  is ?1.
• Halogens ?1 (except when bound to oxygen or a
  halide above it)
• Alkali and alkaline earth metal ions have a
  charge of 1 and 2, respectively.
• Compounds and ions sum of the charges on the
  atoms in a compound add up to 0 and to the ion
  charge in the ion.

Ca in CaO 2
Ca2(aq) 2
Cl?(aq) ?1
Cr in CrO3 6
Fe in Fe2O3 3
Cr in K2Cr2O7 6
14
Displacement Reactions Activity series of the
elements
Li Reacts vigorously with acids to give H2 Reacts with H2O to give H2
K Reacts vigorously with acids to give H2 Reacts with H2O to give H2
Ba Reacts vigorously with acids to give H2 Reacts with H2O to give H2
Ca Reacts vigorously with acids to give H2 Reacts with H2O to give H2
Na Reacts vigorously with acids to give H2 Reacts with H2O to give H2
Mg Reacts with acids to give H2 Reacts slowly with H2O to give H2 more vigorous with steam
Al Reacts with acids to give H2 Reacts slowly with H2O to give H2 more vigorous with steam
Zn Reacts with acids to give H2 Reacts slowly with H2O to give H2 more vigorous with steam
Cr Reacts with acids to give H2 Reacts slowly with H2O to give H2 more vigorous with steam
Fe Reacts with acids to give H2 Reacts slowly with H2O to give H2 more vigorous with steam
Cd Reacts with acids to give H2 Reacts slowly with H2O to give H2 more vigorous with steam
Etc.

• A relative reactivity scale allows us to predict
  if reaction will occur when two substances are
  mixed together.
• E.g. Copper ions in solution are reduced to the
  metal when an iron nail is placed in the
  solution.
• Cu2(aq) Fe(s) ? Fe2(aq) Cu(s) ? Iron
  displaces copper.
• Fe2(aq) Cu(s)? NR ? copper will not displace
  iron.
• Iron more reactive than copper.
• E.g. Predict which reaction will occur when
• Li is mixed with K and
• Li is mixed with K.
• E.g. In which of the following mixtures will
  reaction occur
• Li Mg
• Al Mn2
• Fe Cd2
• Cr Zn2

15
Balancing Oxidation-Number Method

• Determine oxidation for each atom- both sides
  of equation.
• Determine change in oxidation state for each
  atom.
• Left side make loss of electrons gain.
• Balance other side.
• Insert coefficients for atoms that don't change
  oxidation state.
• E.g. Balance
• FeS(s)CaC2(s) CaO(s) ? Fe(s) CO(g) CaS(s)
• In acidic or basic solution balance as above,
  then balance charge with H or OH? on one side
  and water on other side.
• E.g. Balance
• Acidic solution

16
Balancing Half-Reaction Method

• Write unbalanced half reactions for the oxidation
  and the reduction
• Balance the number of elements except O and H for
  each.
• Balance O's with H2O to the deficient side.
• Balance H's with H to the hydrogen deficient
  side
• Acidic add H
• Basic add H2O to the deficient side and OH? to
  the other side.
• Balance charge by adding e? to the side that
  needs it.
• Multiply each half-reaction by integers to make
  electrons cancel.
• Add the two half-reactions and simplify.
• E.g. Balance
• Acidic Zn(s) VO2(aq) ? Zn2(aq) V3(aq).
• Basic Ag(s) HS?(aq) CrO42?(aq) ? Ag2S(s)
  Cr(OH)3(s).

17
Solution Composition, Molarity

• Most reactions performed in solution (homogeneous
  mixture) since reactants mobile.
• Solute dissolved substance.
• Solvent substance in which solute is dissolved.
• Concentration amount of solute dissolved in a
  given amount of solvent.
• Concentrated solution large amount of solute in
  solvent.
• Dilute solution very little solute in solvent.
  Often obtained by dilution.
• Molar concentration ( Molarity, M ) moles of
  solute dissolved in a liter of solution.
• E.g. An aqueous solution of 0.25 M NaCl can be
  prepared by dissolving
• 0.25 mol NaCl in a 1-Liter flask
• 0.50 mol NaCl in a 2-Liter flask
• 0.125 mol in 1/2 liter flask (500 mL).
• E.g. 1 Determine mass needed to prepare exactly 2
  liters of 0.150M NaCl.
• E.g. 2 Determine the concentration when 12.5 g
  NaCl is dissolved and diluted to 500.0 mL.

18
Ion Concentrations in Solutions

• Concentrations of ions after dissolution depends
  on formula
• Determine concentration of each ion in the
  following solutions 0.100 M CaSO4, 0.100 M
  Cu(NO3)2, 0.100 M FeCl3

19
Mass To Molarity

• Often manufacturers provide us with mass of a
  compound in solution, but it is more convenient
  to use molarity.
• E.g. Determine the molarity of NH3(aq) if the
  mass 28.0 NH3 and the density 0.898 g/mL.
  
• Assume 100g of solution
• From mass of solute (28 g NH3) in 100 g determine
  .
• From the mass of solution (100 g) and the
  density, determine V the volume of the
  solution.
• From above steps determine the molar
  concentration
• NH3 n(NH3)/V 1.65 mol NH3/ 0.111L 14.96
  M NH3.

20
DILUTIONS

• Dilute solutions are prepared from more
  concentrated ones by adding solvent to the
  concentrated one.
• The concentration of the dilute solution can be
  determined if we know
• The volume of the concentrated solution, Vi.
• The concentration of the concentrated solution,
  Mi.
• The volume of the dilute solution, Vf.
• The relationship between the molarities and
  volumes is
• E.g. Determine the volume needed to prepare 500.0
  mL of a 0.100 M HCl solution from a 12.40 M stock
  solution.

21
REACTIONS IN SOLUTION

• Reactions usually carried out in solution.
• Amounts of reactants and products (m or n) must
  be determined from the volume and molarity of the
  solution.
• Start with the stoichiometric relationship for
  any reaction
• aA bB ? cC
• Depending upon what is given in the problem
  substitute for mol A, B or C. E.g., if we are
  dealing with a solution we substitute MAVA for
  the mol A.
• E.g. Calculate the volume of 0.200 M KI required
  to react with 50.0 mL of 0.300 M Pb(NO3)2.
• Strategy
• Balance reaction Pb(NO3)2 2KI ? PbI2 2KNO3.
• From stoichiometry
• Substitute for mol
• Solve

22
TITRATIONS

• Titration a procedure for determining the
  amount of one substance A by adding a carefully
  measured amount of a solution B until A is just
  consumed.
• Calculations are the same as in the last
  overhead.
• E.g. What is the molarity of HCl if 25.00 mL of
  it was titrated to the equivalence point with
  33.33 mL of 0.1000 M Ba(OH)2?
• The stoichiometric relationship is
• Substitute the given quantities and solve for the
  HCl.
• The ratio of stoichiometric coefficients tells
  how much of one compound will react if we know
  the amount of the other
• aA bB ? cC
• Solutions n MV solids n m/FM