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Understanding Chemical Reactions: Types, Balancing, and Kinetics**

This comprehensive overview of chemical reactions covers fundamental concepts such as reactants, products, and coefficients. Each type of reaction, including synthesis, decomposition, combustion, and single displacement, is defined with examples and practice equations. The key collision model concept illustrates how molecular interactions drive reactions while discussing factors that affect reaction rates, such as temperature, surface area, catalysts, and concentration. Furthermore, balancing chemical equations adheres to the law of conservation of matter, emphasizing that matter is neither created nor destroyed.

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Understanding Chemical Reactions: Types, Balancing, and Kinetics**

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  1. Chapter 11 Chemical Reactions

  2. Do Now Pg-3 of NotesPkt Label the equation using the following terms: Reactants Products And Yield Coefficient Subscript Solid Liquid Gas Aqueous Yield or produce And or Reacts with Subscript Coefficient 3H2SO4 (aq) + 2Al (s) Al2(SO4)3 (l) + 3H2 (g) Aqueous Liquid Gas Solid Reactants Products

  3. Collision Model Key Idea: • The molecules must touch (or collide) to react with proper orientation. • Collisions must have enough energy (Ea) • Only a small fraction of collisions produces a reaction. Why? Particles lacking the necessary KE to react will bounce apart unchanged when they collide

  4. Activation Energy, Ea Activation energy:is the minimum energy needed for a reaction to occur. Collisions must have enough energyto produce the reaction: must equal or exceed Ea

  5. MnO2 H2O2(aq) H2O(l) + O2(g) Catalyst Substance that speeds up a rxn without being consumed in the rxn • Provides a new pathway with lower Eathan original pathway • Faster rxn • Doesn’t change reactants or products

  6. Exothermic Reaction Pathway Heat of reactants Heat of the products

  7. Endothermic Reaction Pathway Heat of the products Heat of reactants

  8. KINETICS: RATE OF CHEMICAL REACTIONS: What affects the Rate of a Reaction? The rate of a reaction increases when the # of collisions increases. What changes may cause the # of collisions to increase? • Temperature: More KE = particles collide more often • Surface Area (Particle size): More exposed particles allowing more collision to occur • The presence of a catalyst: Lowers Ea (changes rxn pathway) = faster rxn • The nature of the reactants(s, l, g, aq): Size of reactants: high IE or shielding effect • Concentration: Higher concentrated solution have more particles = more collision

  9. Balancing Chemical Equations To follow the law of conservation of matter: Matter is not created or destroyed, ONLY CHANGED! Why balance chemical equations?

  10. Balancing Chemical Equations Coefficients • represent the # of molecules in the reaction (1 is never written) • smallest whole numbers Hydrogen gas (H2) can react with oxygen gas (O2) to form water (H2O) Note! Subscripts should never be changed when trying to balance a chemical equation

  11. Types of Reactions • Synthesis reactions • Decomposition reactions • Combustion reactions • Single displacement reactions • Double displacement reaction Some ways to classify may overlap!!!

  12. 1. Synthesis reactions Synthesis: meaning “to put together” Two or more reactants join together to form one single product reactant + reactant  1 product A + B  AB • Example: 2H2 + O2  2H2O • Example: C+ O2  CO2 or 2C + O2 → 2CO

  13. Synthesis Reactions Magnesium metal and oxygen gas combine to form the compound magnesium oxide. 2Mg(s) + O2→ 2 MgO(s)

  14. Practice Predict the products. Write and balance the following synthesis reaction equations. Sodium metal reacts with chlorine gas Na(s) + Cl2(g)  2Na(s) + Cl2(g)  2NaCl Solid Magnesium reacts with fluorine gas Mg(s) + F2(g)  Mg(s) + F2(g)  MgF2 Aluminum metal reacts with fluorine gas Al(s) + F2(g)  2Al(s) + 3F2(g)  2AlF3 NaCl MgF2 AlF3

  15. 2. Decomposition Reactions Opposite of synthesis rxn One single reactant, a compound, breaks apart into two or more simpler products One Reactant  Product + Product AB  A + B Example: 2 H2O  2H2 + O2 http://youtu.be/OTEX38bQ-2w

  16. heat Decomposition Reactions When mercury(II) oxide is heated, it decomposes or breaks down into two simpler substances. 2HgO(s) 2Hg(l) + O2(g)

  17. Practice Predict the products. Then, write and balance the following decomposition reaction equations: • Solid Lead (IV) oxide decomposes PbO2(s)  PbO2(s)  Pb(s) + O2 (g) • Aluminum nitride decomposes AlN(s)  2AlN(s)  2Al(s) + N2(g) Pb(s) + O2 (g) Al(s) + N2(g)

  18. 3. Combustion Reactions • oxygen reacts with another substance often producing energy in the form of heat and light • Burning!!! 1) A Fuel (hydrocarbon)2) Oxygen to burn it with3) Something to ignite the reaction (spark)

  19. Combustion Reactions of hydrocarbon Combustion of hydrocarbons will produce carbon dioxide and water In general: CxHy+ O2  CO2 + H2O Products: CO2 + H2O(incomplete burning does cause some by-products like CO)

  20. Combustion Practice: Predictthe products. Write and balance the following combustion reaction equations. C5H12 + O2 ANS: C5H12 + 8 O2 5 CO2 + 6 H2O C10H22 + O2  ANS: 2 C10H22 + 31 O2  20 CO2 + 22 H2O CO2 + H2O CO2 + H2O

  21. Do Now (pg-11) Write down the hints to ID the reaction. O2 is a reactant One product One reactant

  22. What happens when ionic compounds dissolve in water? Strong electrolyte: • each unit dissolved in water produces separated ions Ba(NO3)2 = Ba2+ & NO3-

  23. 4. Single Displacement Reactions Occur when one element replaces another in a reaction. A metal can replace another metal or H (+) OR A nonmetal can replace another nonmetal (-) element1 + compound1 element2+ compound2 A + BC  AC + B (if A is a metal/H+) OR A + BC  BA + C (if A is a nonmetal) (remember the cation always goes first!) When H2O splits into ions, it splits into H+ and OH-

  24. Single-Replacement Reactions • Dropping a small piece of potassium into a beaker of water creates the vigorous reaction. • 2K(s) + 2H2O(l) → 2KOH(aq) + H2

  25. Examples metals: Cu + 2AgNO3 Cu(NO3)2 + 2Ag metal + acid: Zn + 2HCl ZnCl2 + H2 metal + water: 2K + 2H2O  2KOH + H2 halogens: 2NaI + Cl2 2NaCl + I2 DON’T FORGET DIATOMIC ELEMENTS BrINClHOF

  26. Practice 1. Zinc metal reacts with aqueous hydrochloric acid Zn(s) + HCl(aq) Note: Zinc replaces the hydrogen ion in the rxn 2. Sodium chloride reacts with fluorine gas 2NaCl(aq) + F2(g)  2NaF(aq) + Cl2(g) Note: fluorine replaces chlorine in the compound 3. Aluminum metal reacts with aqueous copper (II) nitrate 2Al(s) + 3Cu(NO3)2(aq) 3Cu(s) + 2Al(NO3)3(aq) Note: Aluminum replaces copper 2 ZnCl2 (aq) + H2(g)

  27. 5. Double Replacement Reactions an exchange of ions between two ionic compounds (aqueous ONLY) Exchanges ions → forms new compounds AB + CD → AD + CB **A and C are the METALS, they must be written first

  28. Double Replacement Reactions Think about it like “foil”ing in algebra: Outer: first and last ions go together Inner: inside ions go together • Example: AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • Another example: K2SO4(aq) + Ba(NO3)2(aq)  2 KNO3(aq) + BaSO4(s)

  29. Practice Predict the products then balance the equation HCl(aq) + AgNO3(aq)  CaCl2(aq) + Na3PO4(aq)  FeCl3(aq) + NaOH(aq)  H2SO4(aq) + NaOH(aq)  AgCl(s) + HNO3(aq) NaCl(aq) + Ca3(PO4)2(s) 3 2 6 NaCl(aq) + Fe(OH)3(s) 3 3 Na2SO4(aq) + H2O(l) 2 2

  30. Closure Write down the hints to ID the reaction. Reactants are element + ionic compound/acid Reactants are ionic compound + ionic compound/acid

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