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Chemical Reactions

Chemical Reactions. Unit 7 Part 1. Chemical Equations and Reactions. Chemical equations are used to show chemical reactions. Indicators of a Chemical Reaction Include: Change in color (unexpected) Production/release of a gas (bubbles) Change in temperature Exothermic Endothermic

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Chemical Reactions

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  1. Chemical Reactions Unit 7 Part 1

  2. Chemical Equations and Reactions • Chemical equations are used to show chemical reactions. • Indicators of a Chemical Reaction Include: • Change in color (unexpected) • Production/release of a gas (bubbles) • Change in temperature • Exothermic • Endothermic • Formation of a precipitate (solid)

  3. Law of Conservation of Mass • Matter cannot be created nor destroyed, it can only change forms. • Chemical equations must ALWAYS be balanced to obey the law of conservation of mass. • Mass has to be the same on both sides of the reaction. • The number of atoms has to be the same for the products and reactants. • What you start with, you must end with

  4. Components of a Chemical Equation • Reactants • Listed on the LEFT side of the equation. • Starting substances (ingredients) • Products • Listed on the RIGHT side of the equation. • Ending substances (produced substance) • Arrow  “Yields” • States of Matter • Solid (s) • Liquid (l) • Gas (g) • Aqueous (aq) • Coefficients • Whole number that proceeds the chemical substance (number of moles of the substance)

  5. Steps to Follow to Balance Equations • Determine the number of atoms for each element. • Pick an element that is not equal on both sides of the equation. • Add a coefficient in front of the formula with that element and adjust your counts. • Continue adding coefficients to get the same number of atoms of each element on each side. **Note: LEAVE THE SUBSCRIPTS ALONE! (you cannot change the chemical formula to balance atoms)

  6. Coefficients • Coefficients go in front of a compound 4H2O -All atoms in a compound are multiplied by the coefficient Number of atoms: Hydrogen- Oxygen-

  7. Practice With Coefficients *List the amount of each atom in the 4 substances below: 4CH43Ni(NO3)2 3NH3 4Ca3(PO4)2

  8. Class Examples - Model • See hand out for chemical formulas. Balancing Word Equations: (See additional steps on next slide) • Potassium chlorate  potassium chloride + oxygen • When solid copper reacts with aqueous silver nitrate, the products are aqueous copper (II) nitrate and silver metal. • Nitrogen and oxygen combine to form dinitrogenpentaoxide

  9. Steps to Write Equations • When writing formulas don’t forget to look up charges if the compound is ionic!!! • Diatomic Molecules: Always travel in pairs • H.O.F.Br.I.N.Cl. (Dr. HOFINBrCl) • These elements need a subscript 2 after them if they are by themselves! • Once the equation is written, then balance the equation by adding coefficients.

  10. Check for Understanding Balance the following Chemical Equations: • CH4 + O2 CO2 + H2O • Na2O2 + H2SO4  Na2SO4 + H2O2 • N2 + H2 NH3 • When solid copper reacts with aqueous silver nitrate, the products are aqueous copper (II) nitrate and silver metal.

  11. Steps to Follow to Balance Equations • Determine the number of atoms for each element. • Pick an element that is not equal on both sides of the equation. • Add a coefficient in front of the formula with that element and adjust your counts. • Continue adding coefficients to get the same number of atoms of each element on each side. **Note: LEAVE THE SUBSCRIPTS ALONE! (you cannot change the chemical formula to balance atoms)

  12. Steps to Write Equations • When writing formulas don’t forget to look up charges if the compound is ionic!!! • Diatomic Molecules: Always travel in pairs • H.O.F.Br.I.N.Cl. (Dr. HOFINBrCl) • These elements need a subscript 2 after them if they are by themselves! • Once the equation is written, then balance the equation by adding coefficients.

  13. Types of Chemical Reactions Unit 7 Part 2

  14. Types of Chemical Reactions – WHY? • Recognizing patterns allows us to predict future behavior. • Weather experts use patterns to predict dangerous storms so people can get their families to safety. • Political analysts use patterns to predict election outcomes. • Similarly, chemists classify chemical equations according to their patterns to help predict products of unknown but similar chemical reactions.

  15. Synthesis Reaction Synthesis Reaction: A reaction in which two or more substances combine to form ONE new compound. A + B AB

  16. Examples of Synthesis Rxn A + B AB Water: H2+ O2H2O Table Salt: Na + Cl2 NaCl Given the following reactants, predict the products. S8 + O2

  17. Synthesis: 2Mg + O2→ 2MgO

  18. Decomposition Reactions Decomposition Reaction: A reaction in which a single compound breaks down to form two or more simpler substances. AB A + B

  19. Examples of Decomposition Rxn Digestion is a series of decomposition reactions that break down food for fuel for your body. The production of gasoline is done by “cracking” crude oil where you break down carbon and hydrogen molecules. Electrolysis of water: H2OH2+ O2

  20. Decomposition: C12H22O11→ 12C + 11H2O

  21. Single Displacement Reaction Single Displacement Reaction: A reaction in which one element takes the place of another element in a compound. * A more reactive element will take the place of a less reactive element. (see page 286 Table 3) AX+ B BX+ A

  22. Examples of Single Displacement Zn+ HClZnCl2 + H2 Q: In the example above, which element is more reactive? CuCl2 + Al Q: In the above example, predict the products.

  23. Single Replacement: Cu + 2AgNO3 → Cu(NO3)2 + 2Ag

  24. Double Displacement Reaction Double Displacement Reaction: A reaction where there is an apparent exchange of atoms or ions between two compounds. AX+ BY AY + BX Indicators of a double displacement reaction: • Formation of a Precipitate (solid) • Formation of a Gas (bubbles) • Formation of Water

  25. Example of Double Displacement The yellow lines on the roads are made from a double displacement reaction. Pb(NO3)2 + K2CrO4PbCrO4 + KNO3

  26. Double Replacement: 3CuCl2 + 2Na3PO4 → Cu3(PO4)2 + 6NaCl

  27. Combustion Reactions Combustion Reactions: A reaction where an organic molecule is combined with oxygen. CxHy + O2CO2 + H2O • Oxygen (O2) is a reactant. • CO2 and H2O are produced

  28. Uses for Combustion Combustion reactions are used in your home everyday in stoves, water heaters, and furnaces. Example of Combustion Reaction: 2CH4 + 4O22CO2 + 4H2O

  29. Combustion: CH4 + 2O2 → CO2 + 2H2O

  30. Practice (in your notes) • Balance and Classify the following chemical reactions: • __Mg + ___HCl ____MgCl2 + ____H2 • __C4H8 + ___ O2 ____CO2 + ____ H2O • ___Ca(OH)2 + ___H2 SO4 ____CaSO4 + ___H2O • ___N2 + ___O2 ____N2O5

  31. Types of Reactions DEMO Activity • LAB Pictures • Feel free to share – • Email: abechtum@waukeeschools.org • Keys to success- • Be sure your hand is VERY soapy!! • Keep your hand flat • Keep all your fingers together and “tuck your thumb in” • Do NOT move your hand once it is on fire. • Do not forget to SMILE for your picture! 

  32. Chemical Equations/Reactions Exam

  33. Acid-Base Reactions Acid/Base Rxn: The reactants will include an acid combining with a base. The products will include water and a salt. HX+ B(OH) H2O+ BX

  34. Acid-Base Reactions • The Bronsted-Lowry Definitions • Acid is a proton donor. • Base is a proton acceptor. • An acid-base reaction is often called a neutralization reaction.

  35. Examples of Acid/Base Reaction HCl+ Ca(OH)2H2O+ CaCl2 HCN+ K(OH) H2O+ KCN

  36. Ions in Aqueous Solutions and Colligative Properties Chemistry 1 (Chapter 13)

  37. Questions From Readings • Explain what happens when an ionic substance is dissolved in water. (examples to strengthen response)

  38. Dissociation reactions … • AgNO3 • KBr  • BaSO4

  39. Questions from Reading • What is a precipitation reaction? How can you determine if the reaction occurs?

  40. Precipitate Reactions (ppt) • Precipitate: An insoluble solid compound is formed during a reaction. • Anions are exchanged between two cations. • (This is a double displacement reaction) • To be a ppt. rxn, both must occur: • 1. Both reactants must be aqueous (aq) • 2. At least one product must be a solid (s) • You MUST include phase labels with your equation. (See solubility table)

  41. Predicting Solubility of Compounds • Use a solubility table to determine if a substance is going to be soluble (aqueous) or insoluble (solid) in water a) Hg2Cl2 b) KI c) lead (II) nitrate

  42. Practice Solubility 5 Minutes!

  43. Rules for ppt. reactions • Write a balanced chemical equation. • Use the solubility table to place phase labels to each formula. • If one of the products is a solid and the reactants are aqueous the reaction is classified as a precipitate reaction. • If all of the products are (aq) then the reaction is NOT a pptrxn and is classified as double displacement.

  44. Q: For each of the following decide if a ppt. will occur. Aqueous solutions of sodium chloride and iron (II) nitrate are mixed. B) Aqueous solutions of aluminum sulfate and sodium hydroxide are mixed.

  45. Check for Understanding For the following reactions, predict the identity of the precipitate formed. Write the correct formula of the precipitate on the space. If no precipitate is likely, write No Reaction. • BaCl2 and K2SO4 ______________________ • CuCl2and AgNO3______________________ • (NH4)3PO4 and CaS ______________________ • KCland Ca(NO3)2 ______________________

  46. Ionic Equations • Molecular Equation: Chemical equation in which the reactants and products are written as if they were molecular substances, even though they may exist in solutions as ions. • Must include phase labels (s, l, g, aq) • This provides you with the big picture Example: Al2(SO4)3(aq)+ 6NaOH(aq) ----- 2Al(OH)3(s)+ 3Na2SO4(aq)

  47. The next 2 types of equations are only completed for precipitate reactions!!! • Complete ionic equation: Shows all the particles in the solution as they realistically exist. • Break apart the aqueous substances into their ions. • Do NOT break apart s, l, or g!! • When writing a complete ionic equation include the Amount, Symbol and CHARGE!

  48. Example of a Complete Ionic Equation Molecular Equation: Al2(SO4)3(aq)+ 6NaOH(aq)--- 2Al(OH)3(s)+3Na2SO4(aq) Complete Ionic Equation: Amount, Symbol, Charge

  49. Last step for precipitate reaction: • Net ionic equation: Ionic equations that include only the particles that participate in the reaction. • This tells us what substances actually formed something new in the reaction. • Cross out the spectator ions

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