1 / 54

What is the difference between a chemical reaction and physical change?

What is the difference between a chemical reaction and physical change? When you watch a reaction occur, what are some hints that it is a chemical reaction?. Ch. 11 Chemical Equations Reactions. Describing Chemical Reactions. Objectives.

imelda
Télécharger la présentation

What is the difference between a chemical reaction and physical change?

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. What is the difference between a chemical reaction and physical change? • When you watch a reaction occur, what are some hints that it is a chemical reaction?

  2. Ch. 11 Chemical Equations Reactions Describing Chemical Reactions

  3. Objectives • List three observations that suggest that a chemical reaction has taken place. • List three requirements for a correctly written chemical equation. • Write a word equation and a formula equation for a given reaction. • Balance a formula equation by inspection.

  4. Chemical Reactions • when a substance changes identity • reactants- original • products- resulting • law of conservation of mass • total mass of reactants = total mass of products

  5. Chemical Reactions • chemical equation • represents identities and relative amounts of reactants and products in the chemical reaction • uses symbols and formulas

  6. Hints of Chemical Rxn • heat or light • can also happen with physical changes • gas bubbles • means a gas is being created as product • precipitate • solid is being created • color change

  7. Writing Chemical Equations • most pure elements • written as elemental symbol • diatomic molecules • molecule containing only 2 atoms • some elements normally exist this way • H2, O2, N2, F2, Cl2, Br2, I2 • other exceptions • sulfur: S8 • phosphorus: P4

  8. Word Equations • uses names instead of formulas • helps you to write formula equation

  9. Example • Description: Solid sodium oxide is added to water at room temperature and forms sodium hydroxide. • Word Equation: sodium oxide + water  sodium hydroxide • Formula Equation: Na2O + H2O  NaOH

  10. yields reversible above arrow: or heat heated MnO2 or Pt catalyst 25°C specific T requirement 2 atm specific P requirement after a formula: (s) solid (l) liquid (aq) aqueous: dissolved in water (g) gas Symbols Used in Equations

  11. Text Pg. 323 Chart of symbols used in chemical equations

  12. List three observations that suggest that a chemical reaction has taken place.

  13. Acids you have to know! • HCl hydrochloric acid • H2SO4 sulfuric acid • HNO3 nitric acid • H3PO4 phosphoric acid • HC2H3O2 acetic acid

  14. Write the chemical equation from the following description: Zinc metal is added to hydrochloric acid to create zinc chloride and hydrogen gas.

  15. Aluminum reacts with oxygen to produce aluminum oxide • Al + O  Al2O3 • Al + O2 Al2O3 • Al3 + O  Al2O3

  16. Aluminum reacts with oxygen to produce aluminum oxide • Al + O  Al2O3 • Al + O2 Al2O3 • Al3 + O  Al2O3

  17. Phosphoric acid is produced through the reaction between tetraphosphorus decoxide and water • H3PO4 P4 + H2O • H3PO4 + H2O  P4 • P4O10 + H2O  H3PO4

  18. Phosphoric acid is produced through the reaction between tetraphosphorus decoxide and water • H3PO4 P4 + H2O • H3PO4 + H2O  P4 • P4O10 + H2O  H3PO4

  19. Iron(III)oxide reacts with carbon monoxide to produce iron and carbon dioxide • FeO + CO Fe + CO2 • Fe2O3 + CO  Fe + CO2 • Fe + CO Fe2O3 + CO2

  20. Iron(III)oxide reacts with carbon monoxide to produce iron and carbon dioxide • FeO + CO Fe + CO2 • Fe2O3 + CO  Fe + CO2 • Fe + CO Fe2O3 + CO2

  21. Coefficients • whole numbers in front of formula • distributes to numbers of atoms in formula • specifies the relative number of moles and molecules involved in the reaction • used to balance the equation

  22. Balancing Equations • ONLY add/change coefficients- NEVER subscripts!!! • balance one type of atom at a time • balance polyatomic ions first • balance atoms that appear only once second • balance H and O last • simplify if you can • Check at end!

  23. Rules for writing and balancing equations – Pg. 327 in text.

  24. Writing Equations • Write Word equations to help you organize reactants and products • Be sure to include symbols showing states of each reactant and product • Be sure to write the correct formula for each (crossing over for ionic compounds!) • Check your balancing of the equation when you are finished

  25. Example 1 • Description: • Aqueous iron III oxide reacts with hydrogen gas to produce iron metal and liquid water Word Equation: Iron III oxide + hydrogen gas  iron + water

  26. Example 1 • Formula Equation: Fe2O3 (aq) + H2(g) Fe (s) + H2O (l) • Balanced Formula Equation Fe2O3(aq) + 3H2(g) 2Fe (s) + 3H2O (l)

  27. Example 2 • Solid calcium metal reacts with water to form aqueous calcium hydroxide and hydrogen gas. • calcium + water  calcium hydroxide + hydrogen • Ca(s) + H2O(l) Ca(OH)2(aq) + H2(g) • Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)

  28. Example 3 • solid zinc metal reacts with aqueous copper (II) sulfate to produce solid copper metal and aqueous zinc sulfate • zinc + copper (II) sulfate  copper + zinc sulfate • Zn(s) + CuSO4 (aq)  Cu (s) + ZnSO4 (aq) • Zn(s) + CuSO4 (aq)  Cu(s) + ZnSO4 (aq)

  29. Example 4 • Hydrogen peroxide in an aqueous solution decomposes to produce oxygen and water • hydrogen peroxide  oxygen + water • H2O2 (aq)  O2 (g) + H2O (l) • 2H2O2 (aq)  O2 (g) + 2H2O (l)

  30. Example 5 • Solid copper metal reacts with aqueous silver nitrate to produce solid silver metal and aqueous copper (II) nitrate • copper + silver nitrate  silver + copper (II) nitrate • Cu (s) + AgNO3 (aq)  Ag (s) + Cu(NO3)2 (aq) • Cu (s) + 2AgNO3 (aq)  2Ag (s) + Cu(NO3)2 (aq)

  31. Example 6 • Carbon dioxide gas is bubbled through water containing solid barium carbonate, creating aqueous barium bicarbonate • carbon dioxide + water + barium carbonate  barium bicarbonate • CO2 (g) + H2O (l) + BaCO3 (s) Ba(HCO3)2 (aq) • CO2 (g) + H2O (l) + BaCO3 (s) Ba(HCO3)2 (aq)

  32. Example 7 • Acetic acid solution is added to a solution of magnesium bicarbonate to create water, carbon dioxide gas, and aqueous magnesium acetate. • acetic acid + magnesium bicarbonate  water + carbon dioxide + magnesium acetate • HCH3COO (aq) + Mg(HCO3)2 (aq)  H2O(l) + CO2 (g) + Mg(CH3COO)2 (aq) • 2HCH3COO (aq) + Mg(HCO3)2 (aq)  2H2O(l) + 2CO2 (g) + Mg(CH3COO)2 (aq)

  33. Write the balanced formula equation for: Lithium metal is added to a solution of aluminum sulfate to make aqueous lithium sulfate and aluminum metal.

  34. Types of Chemical Reactions

  35. Types of Chemical Reactions • 5 basic types discussed here • not all reactions fall in these categories • you should be able to: • categorize a reaction • predict the product(s)

  36. 1. Synthesis • also called combination reaction • reactants: • more than one • can be elements or compounds • products: only one compound A + X  AX where A is the cation and X is anion

  37. 1. Synthesis • Rubidium and sulfur Rb(s) + S8 (s) Rb2S (s) • Magnesium and oxygen Mg (s) + O2 (g) MgO(s) • Sodium and chlorine Na (s) + Cl2 (g) NaCl(s) • Magnesium and fluorine Mg (s) + F2 (g) MgF2 (s)

  38. 1. Synthesis • calcium oxide and water CaO(s) + H2O(l) Ca(OH)2 (aq) • sulfur dioxide and water SO2 (g) + H2O (l) H2SO3 (aq) • calcium oxide and sulfur dioxide CaO(s) + SO2 (g) CaSO3 (s)

  39. 2. Decomposition • opposite of synthesis • usually require energy • reactants: only one compound • products: more than one • usually elements but can be compounds AX  A + X

  40. 2. Decomposition • water H2O (l) H2 (g) + O2 (g) • calcium carbonate CaCO3 (s) CaO(s) + CO2 (g) • calcium hydroxide Ca(OH)2 (s) CaO(s) + H2O (l) • carbonic acid H2CO3 (aq)  CO2 (g) + H2O (l)

  41. 3. Single Replacement • an element replaces a similar element in a compound • reactants: 1 element & 1 compound • products: 1 element & 1 compound A + BX  B + AX Y + AX  X + AY

  42. 3. Single Replacement • zinc and hydrochloric acid Zn (s) + HCl(aq)  ZnCl2 (aq) + H2 (g) • iron and water Fe (s) + H2O (l)  FeO(aq) + H2 (g) • magnesium and lead (II) nitrate Mg (s) + Pb(NO3)2 (aq)  Mg(NO3)3 (aq) + Pb(s) • chlorine and potassium bromide Cl2 (g) + KBr(s)  KCl(s) + Br2 (g)

  43. 4. Double Replacement • two similar elements switch places • reactants: 2 compounds • products: 2 compounds AX + BY  BX + AY

  44. 4. Double Replacement • barium chloride and sodium sulfate BaCl2 (aq) + Na2SO4 (aq)  NaCl(aq) + BaSO4 (s) • iron sulfide and hydrochloric acid FeS(aq) + HCl(aq)  FeCl2 (aq) + H2S (g) • hydrochloric acid and sodium hydroxide HCl(aq) + NaOH (aq)  NaCl(aq) + H2O (l) • potassium iodide and lead (II) nitrate KI (aq) + Pb(NO3)2 (aq)  KNO3 (aq) + PbI2 (s)

  45. 5. Combustion • Only responsible for one type • releases energy in form of heat/light • reactants: hydrocarbon + O2 • H2O and CO2 as the only products Ex: CH4 + O2 CO2 + H2O

  46. Combustion • propane and oxygen C3H8(g) + O2(g)  CO2(g) + H2O(g)

  47. Practice Classify each of the following reactions one of the five basic types: • Na2O + H2O  NaOH • synthesis • Zn (s) + 2HCl (aq)  ZnCl2 (aq) + H2 (g) • single replacement • Ca(s) + 2H2O (l) Ca(OH)2 (aq) + H2 (g) • single replacement

  48. Practice • 2H2O2 (aq)  O2 (g) + 2H2O (l) • decomposition • Cu (s) + 2AgNO3 (aq)  2Ag (s) +Cu(NO3)2 (aq) • single replacement • C2H4 (g) + O2 (g)  CO2 (g) + H2O (g) • combustion • ZnO(s) + C (s)  2Zn (s) + CO2 (g) • single replacement

  49. Practice • Na2O (s) + 2CO2 (g) + H2O (l) NaHCO3 (s) • synthesis • Ca(s) + H2O(l)  Ca(OH)2(aq) + H2 (g) • single replacement • KClO3 (s)  KCl(s) + O2 (g) • decomposition • H2SO4 (aq) + BaCl2 (aq)  HCl(aq) + BaSO4 (s) • double replacement

  50. Activity Series

More Related