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CHEMICAL REACTIONS AND EQUATIONS

CHEMICAL REACTIONS AND EQUATIONS

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CHEMICAL REACTIONS AND EQUATIONS

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  1. CHEMICAL REACTIONS AND EQUATIONS CHEMISTRY MS. WACK

  2. CHEMICAL REACTIONS In a chemical reaction, bonds are broken, atoms are rearranged, and new bonds are formed. • Something “new” is produced. H2 + O2 H2O

  3. Is it a chemical reaction?

  4. Clues that a chemical reaction has occurred: • Odor • Formation of a gas (may see bubbles) • Precipitate formation • Color Change (not always) • New substances formed • Energy is absorbed or released (not always)

  5. General Description of a Chemical Reaction Reactants  Products The starting substances in a chemical reaction The substances formed in a chemical reaction Word Equations : Use words and symbols to represent chemical reactions. • Uses words and symbols to describe a chemical reaction • Common symbols used: • (s) solid (l) liquid (g) gas (aq) aqueous—dissolved in water •  heat is added to the reaction reversible reaction • a catalyst is added to the reaction + separates reactants or separates products yields/produces catalyst

  6. Word Equations Write the word equations for the following chemical reactions: • Hydrogen gas reacts with chlorine gas to produce hydrogen chloride, a gas with a sharp odor. • Calcium oxide is produced when a piece of calcium is burned in oxygen. HYDROGEN + CHLORINE  HYDROGEN CHLORIDE CALCIUM + OXYGEN  CALCIUM OXIDE

  7. CHEMICAL EQUATIONS • Formula Equations: Use chemical formulas and symbols to represent a chemical reaction. • Keys to writing formula equations: • Write the word equation first—then, replace the words with chemical formulas. • Make sure your formulas are correct!!!!! • Make sure that each element in the reactants is also in the products and vice versa. • The arrow must always go to the right.

  8. CHEMICAL EQUATIONS Write the formula equation for the reaction of calcium burning in oxygen to produce calcium oxide.

  9. LAW OF CONSERVATION OF MASS • How does this law affect chemical equations? • The mass of the reactants must equal the mass of the products • The # of atoms of each element must be equivalent on both sides of the reaction • Chemical reactions must be balanced

  10. What are the #’s called? Mg+2 superscript • 2H2O  2H2 + O2 coefficient subscript

  11. BALANCING EQUATIONS Cd + HCl  CdCl2 + H2

  12. BALANCING EQUATIONS MnSO4  MnO + SO3

  13. BALANCING EQUATIONS H2SO4 + NaOH  Na2SO4 + H2O + +

  14. BALANCING EQUATIONS Co + O2  Co2O3

  15. BALANCING EQUATIONS Write a balanced chemical equation for the reaction in which magnesium reacts with nitrogen to produce magnesium nitride.

  16. BALANCING EQUATIONS Write a balanced chemical equation for the combustion of methane (CH4) in oxygen to produce carbon dioxide and water.

  17. BALANCING EQUATIONS Sodium phosphate is used to cut grease. Write a balanced equation for the reaction in which iron(II) chloride reacts with sodium phosphate to produce sodium chloride and iron(II) phosphate.

  18. BALANCING EQUATIONS Magnesium metal and water combine to form solid magnesium hydroxide and hydrogen gas.

  19. HEAT IN CHEMICAL REACTIONS • Most chemical reactions involve changes in energy. • The unit for energy is the Joule (J) • This is because bond breaking requires energy and bond forming releases energy. • Almost all chemical reactions either release or absorb energy • This energy flow results in heat, either being absorbed or released.

  20. EXOTHERMICREATIONS • To the touch an exothermic reaction would feel HOT because heat is being released to the surroundings • In a chemical reaction that is exothermic, the energy would be represented in the chemical equation on the product side of the reaction (heat is formed/released). • C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) + 2043 kJ

  21. ENDOTHERMIC REACTIONS • To the touch an endothermic reaction would feel cold because heat is being absorbed from the surroundings. • In a chemical reaction that is endothermic, the energy would be represented in the chemical equation on the reactants side of the reaction (since heat is being absorbed/taken in). • C(s)+H2O(g) +113 kJ  CO(g)+H2(g)

  22. Identify each of the following reactions as either an endothermic reaction or an exothermic reaction. • ) C2H4 2C + 2H2 + 52.3 kJ b) B2H6 + 6H2O  2H3BO3 + 6H2 + 493.4 kJ c) 2Fe + 3CO2 + 26.8 kJ  Fe2O3 + 3CO d) Br2 + Cl2 + 29.4 kJ  2BrCl

  23. 5 MAIN CLASSIFICATIONS OF CHEMICAL REACTIONS • Synthesis Reactions • Decomposition Reactions • Combustion Reactions • Single-Replacement Reactions • Double-Replacement Reactions

  24. Synthesis Reactions • Always Have: 1 Product • General Format: A + B  AB • Predicting the Product: • The reactants come together to form one product • If the reactants are a metal & a nonmetal then they will form an ionic compound—get the charges & crisscross • If they are not, the product will be given

  25. Examples of Synthesis Reactions • Na + Cl2

  26. Examples of Synthesis Reactions • H2 + O2 • Al + S 

  27. Decomposition Reactions • Always Have: 1 Reactant • General Format: AB  A + B • Predicting the Product: Separate the reactant into its elements • Example of Decomposition: C6H12O6  C + H2O

  28. Examples of Decomposition • MgS  • Li2O 

  29. Combustion Reactions • Always have: Oxygen (O2) as a reactant • General Formats: • CxHy + O2 CO2 + H2O • A + O2  AxOy • Predicting the product: • If its combustion between a hydrocarbon & oxygen the reaction will produce CO2 & H2O • If its combustion between an element and oxygen it is a synthesis reaction—get your charges and crisscross

  30. COMBUSTION REACTIONS • Always has elemental oxygen (O2) as one of the reactants • Example: • Burning of methane: • CH4 + O2 CO2 + H2O • Burning of propane • C3H8 + O2  CO2 + H2O • Synthesis of sodium oxide • Na + O2 Na2O

  31. Examples • C2H2 + O2 • Na + O2 

  32. Single Replacement Reactions • Always have: 1 element and 1 compound as the reactants and as the products • General Format: A + BC  B + AC • Activity Series: Lists the elements in order of their reactivity. The more reactive elements can replace the less reactive elements • Elements higher on the list can replace elements lower on the list

  33. Which element can replace the other in a chemical reaction? • Tin or Barium • Iodine or Bromine

  34. Single Replacement Reactions • General Format: A + BC  B + AC • Predicting the Product: If the single element is more reactive than the similar element in the compound, those two elements will switch spots. This will form an element and a new ionic compound.

  35. Examples of Single Replacement Reactions • Zn + Cu(NO3)2 • Cu + Zn(NO3)2  • Cu + AgNO3 

  36. More examples of Single Replacement Reactions • Zn(s) + H2SO4(aq)  • Na(s) + H2O(l)  • Sn(s) + NaNO3(aq)  • Cl2(g) + NaBr(aq) 

  37. Double Replacement Reactions • Always have: 2 compounds in the reactants (typically ionic compounds) and 2 compounds in the products • General Format: AX + BY  BX + AY • Double replacement reactions occur if one of the products will be a molecular compound (such as H2O), a precipitate or a gas. • Predicting the product: The metals will switch places—get the charges and crisscross • Types of products: water, gas, or precipitate  + +

  38. Double Replacement Reactions • Precipitate: An insoluble solid formed from two aqueous solutions.  + +

  39. SOLUBILITY RULES Solubility Rules: On the solubility chart on back of your periodic table: “s” = precipitate, “aq” = aqueous (not a precipitate) GENERAL SOLUBILITY RULES

  40. Examples of Double Replacement Reactions • BaCl2(aq) + K2CO3(aq)  • FeS(s) + HCl(aq)  • CaCO3 + HCl 

  41. DISSOCIATION • Double replacement reactions occur between 2 ionic compounds in aqueous solution. • When ionic compounds dissolve in water, they dissociate into the cation and anion of the compound. • DISSOCATION: The separation of ions that occurs when an ionic compound dissolves

  42. Complete Ionic Equation An equation that shows dissolved ionic compounds as dissociated free ions. Chemical Equation: AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) Complete Ionic Equation: Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq)

  43. Net Ionic Equation An equation for a reaction in solution that shows only those particles that are directly involved in the chemical change. A net ionic equation shows only the particles involved in the reaction and is balanced with respect to both mass and charge. Spectator Ion: An ion that appears on both sides of an equation and is not directly involved in the reaction.

  44. Chemical Equation: AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) Complete Ionic Equation: Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq) Any ions that appear on both sides, cross out, they are spectator ions and are not involved in the chemical reaction. Net Ionic Equation: The ionic equation rewritten without the spectator ions. Ag+(aq) + Cl-(aq)  AgCl(s) Once the net ionic equation is written, make sure all atoms are balanced and all charges are balanced.

  45. Example 1 Chemical Equation: Pb(s) + AgNO3(aq)  Ag(s) + Pb(NO3)2(aq) Complete Ionic Equation: Net Ionic Equation: Once the net ionic equation is written, make sure all atoms are balanced and all charges are balanced.

  46. Example 2 Chemical Equation: FeCl3(aq) + KOH(aq)  Fe(OH)3(s) + KCl(aq) Complete Ionic Equation: Net Ionic Equation: Once the net ionic equation is written, make sure all atoms are balanced and all charges are balanced.

  47. Example 3 Word Equation: sodium carbonate + barium nitrate  Chemical Equation: Complete Ionic Equation: Net Ionic Equation: Once the net ionic equation is written, make sure all atoms are balanced and all charges are balanced.