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Chemistry 6.0

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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. 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

  5. B. Energy & Chemical Equations • Exothermicreactions – release energy; energy a product H2O(g)  H2O(l) + 40.7 kJ • Endothermicreactions – absorbs energy; energy a reactant H2O(s) + 6.0 kJ  H2O(l)

  6. C. 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.

  7. D. 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**

  8. BALANCING Examples • sodium + chlorine  sodium chloride • CH4 (g) + O2 (g)  CO2 (g) + H2O(l) • K(s) + H2O(l)  KOH(aq) + H2(g) • AgNO3(aq)+ Cu(s)  Cu(NO3)2(aq) + Ag(s)

  9. Synthesis Reactions • Two or more substances combine to form a more complex product. A + B → AB (only ONE PRODUCT) • A.K.A. Direct Combination Reactions, or composition reactions • Ex. Fe + S → FeS • Ex. CaO + H2O → Ca(OH)2

  10. Synthesis Reaction

  11. Sodium Metal plus Chlorine Gas Video 2 Na + Cl2 2 NaCl Synthesis Reaction

  12. Decomposition Reactions • Single Reactant breaks down to into a simpler substance. AB → A + B (only ONE REACTANT) • The opposite of a synthesis reaction. • Ex. 2HgO → 2Hg + O2 • Ex. CaCO3 → CaO + CO2

  13. Single Replacement Reaction • Atoms of one element replace atoms of another element in a compound. A + BX → AX + B • A more active element will replace a less active element. (See activity series) • Ex. Fe + CuSO4→ FeSO4 + Cu • Ex. Mg + CuSO4 → MgSO4 + Cu

  14. Double-Replacement Reactions • Atoms or ions from 2 different compounds replace each other. AX + BY → AY + BX • Ex. CaCO3 + 2HCl → CaCl2 + H2CO3

  15. Combustion Reactions • One substance reacts with oxygen to produce oxide compounds. • Occurs when burning. • Combustion reactions are often classified as synthesis reactions. • These reactions are usually exothermic, releasing a large amount of energy as light, heat, or sound.

  16. Combustion Reactions (cont.) • When a hydrocarbon is involved in a combustion reaction, H2O and CO2 are the products. • Ex. CH4 + 2O2→ CO2+ 2H2O + 803 kJ • Ex. S + O2 → SO2

  17. Combustion Reaction

  18. 5 Types of Chemical Reactions Video

  19. III. 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

  20. B. Synthesis or Composition 2/more reactants  1 product • Element + Element  Compound A + B  AB 2 Na + Cl2  2NaCl 4 Al + 3 O2 2 Al2O3

  21. 2. Compound + Compound  Compound EXAMPLE 1: metal oxide + carbon dioxide  a carbonate CaO + CO2  CaCO3 EXAMPLE 2: metal oxide + water  a base (hydroxide) Na2O + H2O  2 NaOH EXAMPLE 3: metal oxide + sulfur trioxide  metal sulfate Na2O + SO3  Na2SO4 EXAMPLE 4: metal oxide + sulfur dioxide  metal sulfite Na2O + SO2  Na2SO3 EXAMPLE 5: nonmetal oxide + water  an acid SO3 + H2O  H2SO4

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

  23. Decomposition - Ternary Compounds 2. Ternary Compound Compound + Element/Compound a. metal carbonate  metal oxide + carbon dioxide CaCO3 CaO + CO2 b. metal hydroxide  metal oxide + water (Except Group IA metals) Mg(OH)2 MgO + H2O c. metal chlorate  metal chloride + oxygen 2KClO3  2KCl + 3O2 d. metal nitrate  metal nitrite + oxygen 2NaNO3 2NaNO2 + O2 e. acids  nonmetal oxide + water (HINT: MUST USE CHARGE OF H and O) H2CO3  CO2 + H2O f. Other 2H2O2 2H2O + O2

  24. 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. Zn + CuSO4 ZnSO4 + Cu b. metal + H2O  metal hydroxide + H2 Metals include the alkali metals and calcium. 2Na + 2H2O  2NaOH + H2

  25. Reactivity or Activity of Metals • The reactivity of a metal is based on its ability to replaceanother in a compound. • If a single replacement reaction occurs, the metal that “cuts in” is MOREreactive than the one that was removed or replaced. • An activity series of metals is a listing that ranks metals according to their reactivity. • The most active metal is at the TOPof the list • The least active metal is at the BOTTOMof the list

  26. The ACTIVITY SERIESis listed below: lithium potassium barium strontium calcium sodium magnesium aluminum manganese zinc iron cadmium cobalt nickel tin lead hydrogen copper silver mercury gold The most active metal is LITHIUM The least active metal is GOLD Which is more active nickel or iron? IRON

  27. 3CuCl2 + 2Al  2AlCl3 + 3Cu http://www.youtube.com/watch?v=QQ0eLyBzKYs

  28. 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. Zn + CuSO4 ZnSO4 + Cu b. metal + H2O  metal hydroxide + H2 Metals include the alkali metals and calcium. 2Na + 2H2O  2NaOH + H2

  29. 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

  30. Double Replacement or Double Displacement or Metathesis: Compound + Compound  New Compound + New Compound AB + CD  AD + CB AgNO3 + NaCl  AgCl + NaNO3 Special Notes: If NH4OH is one of the products, it breaks down into NH3 and H2O. If H2CO3 is one of the products, it breaks down into CO2 and H2O.

  31. Double Replacement Reactions and Precipitates • If the reactants are both aqueous in a double replacement reaction, a precipitate may form. • Use solubility rules to determine the identity of the precipitate. • AgNO3 (aq) + NaCl(aq)  AgCl(s)+ NaNO3 (aq)

  32. Exothermic Reactions • Release heat into the surroundings • Heat is a product of the reaction • Combustion reactions are exothermic • C3H8 + 5O2→ 3CO2 + 4H2O + 2043kJ

  33. Endothermic Reactions • Heat is absorbed by the reactants and stored in the chemical bonds of the product. • Heat acts as a reactant. • C + H2O + 113kJ → CO + H2 • Motorcycle Helmet

  34. Molar Heat Capacity • The heat absorbed or released during a reaction depends on a difference in a quantity called enthalpy. (Total energy content of a sample.) • The symbol for enthalpy is H. • When reactions take place at standard temperature and pressure, q = H. • Stand. temp. = 25°C Stand. Pres. = 1 atm • Purest form of a substance = most stable form • Enthalpy change at STP denoted H°

  35. Enthalpy Change∆H = H products – H reactants - > < > + <

  36. Exothermic Reaction Endothermic Reaction

  37. What is the energy of the reactants? • What is the energy of the products? • Is the forward reaction exothermic or endothermic? • What is the ΔH for the forward reaction? • What is the ΔH for the reverse reaction? kJ 150 kJ 50 kJ exothermic -100 kJ 100 kJ

  38. Hess’s Law • Hess’s Law states that if a series of reactions are added together, the enthalpy change of the net reaction will be the sum of the enthalpy changes of the individual steps.

  39. Steps for using Hess’s Law • Identify the compounds • Locate the compounds on the periodic table • Write a reaction from the table. • Write the appropriate “sub equation.” • If needed, multiply equation and enthalpy change. • If you reverse the reaction, change sign of enthalpy change. • Add equations. • Add enthalpy changes.

  40. Bond Energy • Bond energy is the strength of a chemical bond between atoms, expressed as the amount of energy required to break it apart. (Unit kJ/mol) • It is as if the bonded atoms were glued together: the stronger the glue is, the more energy would be needed to break them apart. A higher bond energy, therefore, means a stronger bond. • Ionic bonds are stronger than covalent bonds. Among covalent bonds, triple bonds are stronger than double bonds and double are stronger than single bonds.