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Chemical Equations & Reactions

Chemical Equations & Reactions. Chemistry 6.0. 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

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Chemical Equations & Reactions

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  1. Chemical Equations & Reactions Chemistry 6.0

  2. 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

  3. II. Chemical Equation • Form • Reactant + Reactant  Product + Product • Symbols: (s), (l), (g), (aq) NR

  4. 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)

  5. 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.

  6. 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**

  7. 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)

  8. 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:

  9. 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

  10. 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

  11. SynthesisCompound + Compound  Compound EXAMPLE 1: metal oxide + carbon dioxide  metal carbonateCaO + 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**

  12. Decomposition – Binary Compounds 1. Binary Compound 2 elements AB  A + B 2 H2O  2 H2 + O2 2 HgO  2 Hg + O2

  13. 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

  14. 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

  15. 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

  16. 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.

  17. 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.

  18. surroundings Exothermic Reaction (system) surroundings surroundings surroundings Types of Reactions • Exothermic Reactions: a reaction that releases heat into their surroundings. • Heat is a product of the reaction and temperature of the surroundingsincrease. • This occurs during bond formation.

  19. surroundings Endothermic Reaction (system) surroundings surroundings surroundings Types of Reactions • Endothermic Reactions: a reaction that absorbs heat from the surroundings. • Heat acts as a reactant and temperature of the surroundingsdecreases. • This occurs during bond breaking.

  20. 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)  Ca+2 (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)2·8H2O(s) + 63.9 kJ BaCl2(s) + 2NH3(g) + 10H2O • Rewrite for 1 mole of Cl-1: ½ CaCl2(s)  ½ Ca+2 (aq) + 1 Cl-1(aq) + 44.0kJ

  21. 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.

  22. Enthalpy Diagrams #1 #2

  23. Rewrite each equation with the heat term in the reaction as a reactant or product – THERMOCHEMICAL equation:

  24. products reactants H (kJ) H (kJ) ∆H = - ∆H = + products reactants Course of Reaction Course of Reaction Enthalpy Diagrams BaCl2 + 2NH3 + 10H2O CaCl2 -88.0 kJ +63.9 kJ Ca+2 + 2Cl- 2NH4Cl + Ba(OH)2 8H2O Endothermic Exothermic

  25. 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)

  26. Reaction Pathways or Potential Energy (heat content) Diagrams

  27. Reaction Pathways or Potential Energy (heat content) Diagrams

  28. 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, ΔH, (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?

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

  30. Enthalpy Diagram - Formative Assessment #1 • Sketch a potential energy curve that is represented by the following values of ΔH and Ea. • ΔHreverse = -10 kJ • Eaforward = +40 kJ • Activated Complex = 50 kJ • Is this an endothermic or exothermic reaction?

  31. Enthalpy Diagram - Formative Assessment #2 Based on your diagram, determine: • Endo or Exo? • ΔHforward = • Eaforward = • ΔHreverse = • Eareverse =

  32. 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)

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

  34. Steps for using Hess’s 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 ΔH for each sub equation. • If needed, multiply the sub equation and the associated ΔH’s (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 ΔH’s of each sub equation to calculate the ΔH for the desired balanced equation.

  35. Calculate ΔH for the following example: #1) XeF2 + F2 XeF4 ΔH = ? Xe + F2  XeF2 ΔH = -123 kJ Xe + 2F2 XeF4 ΔH = -262 kJ #2) C + H2O → CO + H2 ΔH = ? 2CO  2C + O2 ΔH = +222 kJ 2H2 + O2 2H2O ΔH = -484 kJ XeF2 + F2 XeF4 ΔH = -139 kJ C + H2O → CO + H2 ΔH = +131 kJ

  36. Calculate ΔH for the following example: #3) CO + O2 → 2 CO2 ΔH = ? 2C + O2  2CO ΔH = -222 kJ CO2 C + O2ΔH = +394 kJ #4) H2O2 + H2 → 2 H2O ΔH = ? H2O + ½ O2  H2O2 ΔH = +94.6 kJ 2H2 + O2 2H2O ΔH = -484kJ CO + O2 → 2 CO2 ΔH = -394 kJ H2O2 + H2 → 2H2O ΔH = -336.6kJ

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