Chemical Equations

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Chemical Equations Reactions

• Chemistry 6.0

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I. Chemical Reactions

• Definition a process by which 1 or more
  substances, called reactants, are changed into 1
  or more substances, called products, with
  different physical chemical properties.
• Evidence of a Chemical Reaction
• Color change
• Formation of a precipitate, ppt
• Release of a gas
• Energy change heat, light, sound
• Odor change
• Reactions are started by the addition of energy

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II. Chemical Equation

• Form
• Reactant Reactant ? Product Product
• Symbols (s), (l), (g), (aq)
• NR

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Writing Chemical Equations

• Two moles of water at room temperature are
  exposed to an electric current and produces two
  moles of hydrogen gas and one mole of oxygen gas.
• When two moles of aluminum pellets are added to
  three moles of a copper(II) chloride solution, 3
  moles of copper precipitate out and two moles of
  aluminum chloride remain in solution.

2 H2O(l) ? 2 H2(g) 1 O2(g)
2 Al(s) 3 CuCl2(aq) ? 3 Cu(s) 2 AlCl3(aq)
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Characteristics of A Balanced Chemical Equations

• The equation must represent known facts. All
  substances have been identified.
• The equation must contain the correct symbols
  and/or formulas for the reactants and products
• Can be either a word equation or a formula
  equation
• The law of conservation of mass must be
  satisfied. This provides the basis for balancing
  chemical equations. 1st formulated by Antoine
  Lavoisier
• TOTAL MASS REACTANTS TOTAL MASS PRODUCTS
• Number of atoms of EACH element is the SAME on
  both sides of the equation.

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Balancing Chemical Equations

• Balance using coefficients after correct formulas
  are written.
• Coefficients are usually the smallest whole
  number required when interpreted at the
  molecular level
• Balance atoms one at a time
• Balance the atoms that are combined and appear
  only once on each side.
• Balance polyatomics that appear on both sides
• Balance H and O atoms last
• NEVER CHANGE SUBSCRIPTS!!!
• Count atoms to be sure that the
• equation is balanced

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BALANCING Examples

• sodium chlorine ? sodium chloride
• CH4 (g) O2 (g) ? CO2 (g) H2O(l)
• K(s) H2O(l) ? KOH(aq) H2(g)
• HOH
• AgNO3(aq) Cu(s) ? Cu(NO3)2(aq) Ag(s)

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Interpretation of a Balanced Equation

• 2Mg(s) O2(g) ? 2MgO(s)
• 2 atoms of solid magnesium react with
• 1 molecule of oxygen gas to form
• 2 formula units of solid magnesium oxide
• OR
• 2 moles of solid magnesium react with
• 1 moles of oxygen gas to form
• 2 moles of solid magnesium oxide
• Reaction Ratios

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Classifying Chemical Reactions

• Pattern for prediction based
• on the kind of reactants
• Combustion or Burning complete combustion
  always produces carbon dioxide and water!
• Hydrocarbons
• CxHy O2 ? CO2 H2O
• Alcohols
• CxHyOH O2 ? CO2 H2O
• Sugars
• C6H12O6 O2 ? CO2 H2O
• C12H22O11 O2 ? CO2 H2O

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Synthesis or Composition

• 2/more reactants ? 1 product
• Element Element ? Compound
• A B ? AB
• 2 Na Cl2 ? 2 NaCl
• 4 Al 3 O2 ? 2 Al2O3

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SynthesisCompound Compound ? Compound

• EXAMPLE 1 metal oxide carbon dioxide ? metal
  carbonate CaO CO2 ? CaCO3
• EXAMPLE 2 metal oxide water ? a base
  (hydroxide)
• Na2O H2O ? 2 NaOH
• H(OH)
• EXAMPLE 3 nonmetal oxide water ? an acid
• SO3 H2O ? H2SO4
• Determine oxidation numbers for molecular
  compounds and oxyacids

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Decomposition Binary Compounds

• 1. Binary Compound ? 2 elements
• AB ? A B
• 2 H2O ? 2 H2 O2
• 2 HgO ? 2 Hg O2

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Decomposition - Ternary CompoundsTernary
Compound ? Compound Element/Compound

• EXAMPLE 1 metal chlorate ? metal chloride
  oxygen
• 2KClO3 ? 2KCl 3O2
• EXAMPLE 2 metal carbonate ? metal oxide
  carbon dioxide
• CaCO3 ? CaO
  CO2
• EXAMPLE 3 metal hydroxide ? metal oxide
  water
• Mg(OH)2 ? MgO H2O (Except
  Group IA metals)
• EXAMPLE 4 acids ? nonmetal oxide water
• H2CO3 ? CO2 H2O
• EXAMPLE 5 Hydrogen Peroxide 2H2O2 ?
  2H2O O2

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Single Replacement or Single Displacement

• Element Compound ? New Compound New Element
• Metals A BC ? AC B
• Active metals displace less active metals or
  hydrogen from their compounds in aqueous
  solution. Refer to the Activity Series.
• a. 2Al 3CuCl2 ? 2AlCl3 3Cu
• b. metal H2O ? metal hydroxide
  H2
• An active metal (top of series to calcium) will
  react with water to form the hydroxide of the
  metal and hydrogen gas.
• 2Na 2HOH ? 2NaOH H2

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Single Replacement or Single Displacement

• 2. Nonmetals
• D EF ? ED F
• Cl2 2NaBr ? 2NaCl Br2
• Many nonmetals displace less active nonmetals
  from combination with a metal or other cation.
  Order of decreasing activity is
• F2 ? Cl2 ? Br2 ? I2 ? S8

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Double Replacement or Double Displacement or
Metathesis

• Compound Compound ? New Compound New
  Compound
• AB CD ? AD CB
• AgNO3 NaCl ? AgCl
  NaNO3
• AgNO3 (aq) NaCl (aq) ? AgCl (s)
  NaNO3 (aq)
• The driving force for these reactions is if it
  produces a
• A precipitate (ppt) See Solubility Table
• Water
• Gas Only HCl and NH3 are soluble in water. All
  other gases (CO2 and H2S) are sufficiently
  insoluble to force a reaction to occur if they
  are found as a product.

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Thermochemistry

• The study of the changes in energy that accompany
  a chemical reaction and physical changes.
• Chemical Reactions involve changes in energy that
  result from
• Bond breaking that requires energy (absorbs) from
  the surroundings.
• Bond making that produces energy (releases) to
  the surroundings.
• Changes in energy result in an energy flow or
  transfer.

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Types of Reactions

• Exothermic Reactions a reaction that releases
  heat into their surroundings.
• Heat is a product of the reaction and temperature
  of the surroundings increase.
• This occurs during bond formation.

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Types of Reactions

• Endothermic Reactions a reaction that absorbs
  heat from the surroundings.
• Heat acts as a reactant and temperature of the
  surroundings decreases.
• This occurs during bond breaking.

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Energy Chemical EquationsCoefficients are
always interpreted as moles. Physical states are
written influences the overall energy
exchanged. Very specific!

• Exothermic release energy E product
• CaCl2(s) ? Ca2 (aq) 2Cl-1(aq) 88.0kJ
• Combustion reactions are ALWAYS exothermic
• C3H8 5O2 ? 3CO2 4H2O 2043 kJ
• Endothermic absorbs energy E reactant
• 2NH4Cl(s) Ba(OH)28H2O(s) 63.9 kJ ?
• BaCl2(s) 2NH3(g) 10H2O
• Rewrite for 1 mole of Cl-1
• ½ CaCl2(s) ? ½ Ca2 (aq) 1 Cl-1(aq)
  44.0kJ

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Heat and Enthalpy Changes

• Enthalpy (H) the heat content of a system at
  constant pressure.
• There is no way to directly measure the enthalpy
  of a substance or system.
• Unit J
• Enthalpy Change (?H) is the heat absorbed or
  released in a physical or chemical change at
  constant pressure.
• ?H Hproducts - Hreactants
• This can be measured.
• ?H is also known as the heat of the reaction.
• Difference between the stored energy of the
  reactants and the products.

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Enthalpy Diagrams
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2
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1 2
Which has a higher enthalpy? Products or Reactants R P
b. Was heat absorbed or released? R A
c. Is this an endothermic or exothermic reaction? Exo Endo
d. Is ?H for this reaction positive or negative? -
e. Would the ?H be on the left or right side of the yield sign? R L
f. Is the reverse reaction exothermic or endothermic? Endo Exo
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Rewrite each equation with the heat term in the
reaction as a reactant or product
THERMOCHEMICAL equation
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Enthalpy Diagrams
BaCl2 2NH3 10H2O
CaCl2
-88.0 kJ
63.9 kJ
Ca2 2Cl-
2NH4Cl Ba(OH)2 8H2O
Exothermic
Endothermic
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Reaction Progress

• Collision Theory
• In order for a reaction to occur, the particles
  must collide
• A successful or effective collision occurs when
• The collision is energetic enough
• The particles collide with the correct
  orientation
• During a collision, kinetic energy is converted
  to potential energy
• The minimum energy needed for a successful
  collision activation energy (Ea)

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Reaction Pathways or Potential Energy (heat
content) Diagrams
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Reaction Pathways or Potential Energy (heat
content) Diagrams
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Answer the following questions based on the
potential energy diagram shown here

• Does the graph represent an endothermic or
  exothermic reaction?
• Label the position of the reactants, products,
  and activated complex.
• Determine the heat of reaction, ?H, (enthalpy
  change) for this reaction.
• Determine the activation energy, Ea for this
  reaction.
• How much energy is released or absorbed during
  the reaction?
• How much energy is required for this reaction to
  occur?

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Practice

• Sketch a potential energy curve that is
  represented by the following values of ?H and Ea.
  You may make up appropriate values for the y-axis
  (potential energy).
• ?Hforward -20 kJ
• Earev 80 kJ
• Activated Complex 120 kJ
• Is this an endothermic or exothermic reaction?

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Enthalpy Diagram - Formative Assessment 1

• Sketch a potential energy curve that is
  represented by the following values of ?H and Ea.
• ?Hreverse -10 kJ
• Eaforward 40 kJ
• Activated Complex 50 kJ
• Is this an endothermic or exothermic reaction?

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Enthalpy Diagram - Formative Assessment 2

• Based on your diagram, determine
• Endo or Exo?
• ?Hforward
• Eaforward
• ?Hreverse
• Eareverse

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Calculating ?H using Bond Energy

• 2 H2 O2 ? 2 H2O
• Bonds Formed
• Bonds Broken
• Using Bond Energy Table, determine ?H.
• ?H -482 kJ (482 kJ released exothermic)

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Hesss Law

• The enthalpy change for a reaction is the sum of
  the enthalpy changes for a series of reactions
  that adds up to the overall reaction.
• This is also called the Law of Heat of Summation
  (S)
• 3. This allows you to determine the enthalpy
  change for a reaction by indirect means when a
  direct method cannot be done.

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Steps for using Hesss Law

• Write a balanced equation.
• Identify the compounds.
• Write the reaction from the table so the compound
  is a reactant or product as it appears in the
  balanced equation.
• Write appropriate ?H for each sub equation.
• If needed, multiply the sub equation and the
  associated ?Hs (coefficients).
• If you reverse the equation, change the sign of
  the enthalpy change.
• Add the sub equations to arrive at the desired
  balanced equation.
• Add ?Hs of each sub equation to calculate the ?H
  for the desired balanced equation.

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Calculate ?H for the following example

• 1) XeF2 F2 ? XeF4 ?H ?
• Xe F2 ? XeF2 ?H -123 kJ
• Xe 2F2 ? XeF4 ?H -262 kJ
• 2) C H2O ? CO H2 ?H ?
• 2CO ? 2C O2 ?H 222 kJ
• 2H2 O2 ? 2H2O ?H -484 kJ

XeF2 F2 ? XeF4 ?H -139 kJ
C H2O ? CO H2 ?H 131 kJ
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Calculate ?H for the following example

• 3) CO O2 ? 2 CO2 ?H ?
• 2C O2 ? 2CO ?H -222 kJ
• CO2 ? C O2 ?H 394 kJ
• 4) H2O2 H2 ? 2 H2O ?H ?
• H2O ½ O2 ? H2O2 ?H 94.6 kJ
• 2H2 O2 ? 2H2O ?H -484kJ

CO O2 ? 2 CO2 ?H -394 kJ
H2O2 H2 ? 2H2O ?H -336.6kJ