PRINCIPLES OF CHEMICAL REACTIVITY: CHEMICAL REACTIONS

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PRINCIPLES OF CHEMICAL REACTIVITY CHEMICAL
REACTIONS

• AP Chemistry Chapter Four

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Chemical equations

• A chemical equation is the representation of a
  chemical reaction using the symbols of the
  elements and the formulas of the compounds
• The reacting species (starting materials or
  reactants) are listed on the left, and the
  products of the reaction are listed on the right

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Four Things to Know About Equations

• One Reactants are on the left and products are
  on the right of the arrow.
• Two Coefficients are the numbers in front of
  each formula. If no number is shown, a one is
  understood.
• Three The coefficients tell us how many
  molecules (moles) of each reactant used and how
  many molecules (moles) of each product made.
• Four The phases of the reactants and products
  are indicated in parenthesis after the formula.
• gas (g)
• liquid (l)
• solid (s)
• aqueous (in a solution) (aq)

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Balanced Chemical Equation

• C6H12O6(s) 6O2(g) ? 6CO2(g) 6H2O(l)

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Balancing Chemical Equations

• We balance chemical equations because of the Law
  of the Conservation of Mass
• We use coefficients in front of the element or
  compound formula, which are whole numbers that
  adjust the amounts of the species
• Note that we cannot balance an equation by
  changing the subscripts of the compound

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Balancing coefficients

• The subscripts of the compound are fixed they
  cannot be changed in an equation
• The coefficients used should be the smallest
  whole numbers possible
• The coefficient multiplies every number in the
  formula (e.g. 2 MgCl2 means that there are 2 Mg
  atoms and 4 Cl atoms)
• We will balance equations by inspection

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Balance the following

• Si2H3 O2 ? SiO2 H2O
• Al(OH)3 H2SO4 ? Al2(SO4)3 H2O

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Types of chemical reactions

• Combustion reactions - reaction of an element or
  a compound with elemental oxygen (O2)
• These reactions are often accompanied by the
  release of large amounts of heat (and flame)
• Consider 2C2H6 (g) 7O2 (g) ? 4CO2(g) 6H2O(g)

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Combination reactions

• Often called synthesis reactions
• Involves two or more substances combining to form
  a single substance
• Consider
• 2Mg(s) O2 (g) ? 2MgO(s)

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Decomposition reactions

• The reverse of a combination reaction
• One substance breaking apart into two or more
  substances
• Consider
• H2CO3 (aq) ? H2O(l) CO2 (g)

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Examples of Decomposition Reactions

• Metallic carbonates yield metallic oxides and
  carbon dioxide
• Metallic chlorates yield metallic chlorides and
  oxygen
• Ammonium carbonate yields ammonia, water and
  carbon dioxide
• Sulfurous acid yields sulfur dioxide and water
• Carbonic acid yields carbon dioxide and water
• Ammonium hydroxide yields ammonia and water
• Hydrogen peroxide yields water and oxygen
• Binary compounds produce two elements

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Reactions in Aqueous Solution

• Single Replacement
• Double Replacement
• Oxidation-Reduction
• All require an understanding of aqueous solutions!

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Formation of ions in solution

• When we consider a substance dissolving in
  another substance, we define the solute as the
  substance in lesser amount, and the solvent as
  the substance in greater amount
• Solubility is the relative degree to which a
  substance dissolves in a solvent
• Soluble substances dissolve to a considerable
  degree
• Insoluble substances do not dissolve at all or
  only to a small degree

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Dissolution of ionic compounds

• When ionic compounds dissolve in water, they
  dissociate into separate anions and
  cationsNaCl(s) ? Na(aq) Cl-(aq)
• Acids are so named because when they dissolve,
  they ionize to yield H and the associated
  anionHCl(aq) ? H(aq) Cl-(aq)

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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
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Strong Acids

• Strong acids completely dissociate into ions upon
  dissolution in water
• With strong acids the major components in
  solution are the hydrogen ion and the anion.
• Examples are HCl, H2SO4

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Weak Acids

• Weak acids dissolve, but do not completely
  dissociate into H and anions.
• With weak acids the major component in solution
  is the weak acid.
• Examples are CH3CO2H, HF

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Single replacement reactions

• An element replaces another in a compound
• Zn(s) CuCl2 (aq) ? ZnCl2 (aq) Cu(s)
• This particular form of the equation is termed
  the molecular equation, where all reactants and
  products are shown as neutral compounds

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Types of Single Replacemnt Reactions

• Active metals replace less active metals from
  their compounds in aqueous solution
• Active metals replace hydrogen in water
• Active meatls replace hydrogen in acids
• Active metals replace less active metals from
  their compounds in aqueous solution

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Ionic equations

• The total ionic equation writes all of the anions
  and cations separatelyZn(s) Cu2(aq) 2Cl-
  ? Zn2(aq) 2Cl- Cu(s)
• Ions (such as the Cl- ions) that are in the same
  state on both sides of the equation are called
  Spectator ions.

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Net Ionic Equations

• In a net ionic equation the spectator ions are
  omitted.
• Net ionic equations are only used for reactions
  occurring in solution between ionic compounds
  single and double replacement reactions as well
  as oxidation-reduction

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Net ionic equation

• The net ionic equation leaves out the spectator
  ions
• The complete ionic equationZn(s) Cu2(aq)
  2Cl- ? Zn2(aq) 2Cl- Cu(s)
• The net ionic equationZn(s) Cu2(aq) ?
  Zn2(aq) Cu(s)

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Double replacement reactions

• Where the anions and cations in an ionic equation
  swap places.
• AgNO3 (aq) NaCl(aq) ? AgCl(s) NaNO3 (aq)

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What causes double replacements?

• There has to be some process to cause the
  reaction to take place
• Precipitation of a solid
• Evolution of a gas
• Formation of a Molecular Substance
• Neutralization reaction between an acid and a
  base
• If one of these processes does not occur, no
  reaction takes place

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Double replacement reactions

• Precipitation
• AgNO3 (aq) NaCl(aq) ? AgCl(s) NaNO3
  (aq)
• Gas evolution
• 2HCl(aq) Na2CO3 (aq) ? 2NaCl(aq) CO2
  (g)

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Double replacement reactions

• Neutralization
• an acid is a proton donor
• a base is a hydroxide donor
• proton plus hydroxide yields water
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)

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Write the Net Ionic Equation

• AgNO3 (aq) NaCl(aq) ? AgCl(s) NaNO3 (aq)
• 2HCl(aq) Na2CO3 (aq) ? 2NaCl(aq) CO2 (g)
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)

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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.

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Oxidation-Reduction Cont.

• 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)

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Oxidation Number

• Oxidation number (state)
• The charge that an atom in a substance or
  monatomic ion has or would have if all the
  bonding electrons were assigned to the most
  electronegative element.

Ca in CaO 2
Ca2(aq) 2
Cl?(aq) ?1
Cr in CrO3 6
Fe in Fe2O3 3
Cr in K2Cr2O7 6
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Assigning Oxidation Numbers

• 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.

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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

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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.

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Example

• FeS(s) CaC2(s) CaO(s) ? Fe(s) CO(g)
  CaS(s)

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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.

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Half-Reaction Acidic

• Zn(s) VO2(aq) ? Zn2(aq) V3(aq).

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Half Reaction Basic

• Ag(s) HS?(aq) CrO42?(aq) ? Ag2S(s)
  Cr(OH)3(s).