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Introduction to Chemical Reactions: Understanding Equations and Balancing Techniques

This overview of chemical reactions provides insights into the fundamental processes where substances transform into new products. Chemical equations represent these transformations through symbols and formulas, adhering to the law of conservation of mass. Key indicators such as heat, light, gas formation, precipitate formation, and color change signify chemical changes. This guide also covers the essential steps to balance chemical equations, ensuring atom counts on both sides of an equation are equal, alongside useful tips for effective balancing.

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Introduction to Chemical Reactions: Understanding Equations and Balancing Techniques

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  1. Ch. 9 – Chemical Reactions Intro to Reactions

  2. BACKGROUND • chemical reaction: process in which 1 or more substances are changed into 1 or more different substances. • chemical equation: uses symbols and formulas to represent the identities and relative amounts of reactants and products in a chemical reaction

  3. A.Signs of a Chemical Reaction • Evolution of heat and light • Formation of a gas • Formation of a precipitate • precipitate: solid produced as result of chemical rxn in a soln. and separates from the soln. • Color change

  4. CHEMICAL EQUATIONS • Three requirements for chemical equations: • a. Equation must represent known facts: actual compounds and actual lab results

  5. CHEMICAL EQUATIONS • b. The equation must contain correct formulas for reactants and products • ex: diatomic elements (7) H2, N2, O2, F2, Cl2, Br2, I2

  6. CHEMICAL REACTIONS c. Law of Conservation of Mass must be satisfied • mass is neither created nor destroyed in a chemical reaction • total mass stays the same • atoms can only rearrange 4 H 2 O 4 H 2 O 36 g 4 g 32 g

  7. C. Chemical Equations A+B  C+D PRODUCTS REACTANTS • word equation: (def)reactants and products represented by words • hydrogen + oxygen  water • formula equation: (def)represents reactants and products of a chemical equation by their symbols or formulas • 2H2 + O2 2H2O

  8. symbols in chemical equations

  9. symbols in chemical equations

  10. Reversible Reaction • reversible reaction: the products re-form the original reactants Reversible Reaction Oscillating Reaction

  11. Significance of Chemical Eqn. • Coefficients show relative amounts of reactants and products 2H2 + O2 2H2O 2 molecule H2: 1 molecule O2: 2 molecules H2O

  12. Significance of Chemical Eqn. • Relative amounts of reactants and products can be determined from reaction coefficients. H2 + Cl2 2HCl 1 molecule H2: 1 molecule Cl2: 2 molecules HCl 1 mol H2: 1 mol Cl2: 2 mol HCl 2.02 g H2: 70.9 g Cl2: 72.92 g HCl

  13. Ch. 8 – Chemical Reactions II. Balancing Equations(p. 250-254)

  14. A. Balancing Steps 1. Write the unbalanced equation. 2. Count atoms on each side. 3. Add coefficients to make #s equal. Coefficient  subscript = # of atoms 4. Reduce coefficients to lowest possible ratio, if necessary. 5. Double check atom balance!!!

  15. B. Helpful Tips • Balance one element at a time. • Update ALL atom counts after adding a coefficient. • If an element appears more than once per side, balance it last. • Balance polyatomic ions as single units. • “1 SO4” instead of “1 S” and “4 O”

  16. C. Balancing Example Al + CuCl2 Cu + AlCl3 Al Cu Cl Aluminum and copper(II) chloride react to form copper and aluminum chloride. 2 3 3 2  2  6 1 1 1 1 2 3 2  3  6   3

  17. Chemical Reactions II. Balancing Equations

  18. A. Balancing Steps 1. Write the unbalanced equation. 2. Count atoms on each side. 3. Add coefficients to make #s equal. Coefficient  subscript = # of atoms 4. Reduce coefficients to lowest possible ratio, if necessary. 5. Double check atom balance!!!

  19. B. Helpful Tips • Balance one element at a time. • Update ALL atom counts after adding a coefficient. • If an element appears more than once per side, balance it last. • Balance polyatomic ions as single units. • “1 SO4” instead of “1 S” and “4 O”

  20. C. Balancing Example Al + CuCl2 Cu + AlCl3 Al Cu Cl Aluminum and copper(II) chloride react to form copper and aluminum chloride. 2 3 3 2  2  6 1 1 1 1 2 3 2  3  6   3

  21. Ch. 9 – Chemical Reactions III. Types of Chemical Reactions

  22. Synthesis A + B  AB • the combination of 2 or more substances to form a compound • only one product

  23. Synthesis H2(g) + Cl2(g)  2 HCl(g)

  24. Synthesis • Products: • ionic - cancel charges • covalent - hard to tell Al(s)+ Cl2(g)  AlCl3(s) 2 3 2

  25. Synthesis • Reactions with oxides • Oxides of active metals react with water to form metal hydroxides • Active metal: Ca metal oxide: CaO reaction: CaO + H2O  Ca(OH)2

  26. Synthesis • Examples: • Mg + O2 → • S8 + O2 → • Co + F2 → • BaO + H2O → MgO 2 2 8 SO2 8 2 2 CoF3 3 Ba(OH)2

  27. Decomposition AB  A + B • a compound breaks down into 2 or more simpler substances • only one reactant

  28. Decomposition 2 H2O(l)  2 H2(g) + O2(g)

  29. Decomposition • Products: • binary - break into elements • ((others - hard to tell – various rules)) KBr(l)  K(s) + Br2(l) 2 2

  30. Decomposition • Products for metal carbonates • MCO3  metal oxide + CO2 • CaCO3  CaO + CO2 • Products for metal hydroxides • NOT GROUP 1 METALS • M(OH)x metal oxide + water • Ca(OH)2 CaO + H2O

  31. Decomposition • Products for metal chlorates • MClO3  metal chloride + O2 • KClO3  KCl + O2 • Products for acids • acid  nonmetal oxide + water (generally) • H2SO4 SO3 + H2O 2 3 2

  32. Single Replacement A + BC  B + AC • one element replaces another similar element in a compound • metal replaces metal • nonmetal replaces nonmetal

  33. Single Replacement Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)

  34. Single Replacement • Activity Series: Part of series Li > Rb > K > Ba > Sr > Ca > Na An element can only replace another if the free element is more reactive than the element in the compound Li + KCl  K + LiCl  LiCl + K NR (no reaction)

  35. Single Replacement • Replacement of metal by metal • Fe2O3+ AlAl2O3+Fe • Replacement of H in water by a metal • Na + H2O  NaOH + H2 • metal + water  metal hydroxide + hydrogen 2 2 2 2 2

  36. Single Replacement • Replacement of hydrogen in an acid by a metal • active metal + acid  metal compound + hydrogen • Mg + HCl  MgCl2 + H2 • Replacement of halogens • Cl2 + NaBr  NaCl + Br2 2 2 2

  37. Double Replacement AB + CD  AD + CB • ions in two compounds “change partners” • cation of one compound combines with anion of the other

  38. Double Replacement Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)

  39. E. Double Replacement • Formation of a precipitate • 2 aqueous compounds form insoluble solid • 2KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2KNO3(aq) • Formation of a gas • 1 product bubbles out of solution • FeS(s) + 2HCl(aq) H2S(g) + FeCl2(aq)

  40. Double Replacement • Formation of water • HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

  41. Combustion A + O2 B • the burning of any substance in O2 to produce heat CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)

  42. Combustion • Products: • contain oxygen • hydrocarbons form CO2 + H2O Na(s)+ O2(g)  Na2O(s) 4 2 C3H8(g)+ O2(g)  CO2(g)+ H2O(g) 5 3 4

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