1. Chapter 7 Chemical reactions

2. Day 1&2 Types of reactions lab and write up

3. Day 3 7.1-7.4

4. 7.1- K. volcanoes, Autos, and Detergents  Gas Evolution reactions- rxn in a liquid forming a gas  Oxidation- Reduction reactions- rusting, combustion rxn, electrons are transferred  Precipitation rxn- solids form in water  Chemical rxn- a change in arrangement of molecules!

5. 7.2 -Evidence of Chemical Reactions  Color Change  Formation of a solid- precipitate  Formation of a gas- bubbles or odor  Energy absorption or emission- heat or light  Changes in state are not chemical changes!

6. Evidence of Chemical Change Color Change Formation of Solid Precipitate Formation of a Gas Emission of Light Release or Absorption of Heat

7. 7.3- Chemical Equations  Short-hand way of describing a reaction.  Provides information about the reaction.  Formulas of reactants and products.  States of reactants and products.  Relative numbers of reactant and product molecules that are required.  Can be used to determine masses of reactants used and products that can be made.

8. Symbols Used in Equations  Symbols used to indicate state after chemical.  (g) = gas; (l) = liquid; (s) = solid.  (aq) = aqueous = dissolved in water.  Energy symbols used above the arrow for decomposition reactions.   = heat.  h = light.  shock = mechanical.  elec = electrical.

9. Conservation of Mass  Matter cannot be created or destroyed.  And the total mass of the reactants will be the same as the total mass of the products.  In a chemical reaction, all the atoms present at the beginning are still present at the end.

10. The Combustion of Methane  Methane gas burns to produce carbon dioxide gas and gaseous water.  Whenever something burns it combines with O 2 (g).

11. Chemical Equations  CH 4 and O 2 are the reactants, and CO 2 and H 2 O are the products.  The (g) after the formulas tells us the state of the chemical.  The number in front of each substance tells us the numbers of those molecules in the reaction.  Called the coefficients. CH 4 (g) + 2 O 2 (g)  CO 2 (g) + 2 H 2 O(g)

12. Chemical Equations, Continued  This equation is balanced, meaning that there are equal numbers of atoms of each element on the reactant and product sides.  To obtain the number of atoms of an element, multiply the subscript by the coefficient. 1  C  1 4  H  4 4  O  2 + 2 4  O  2 + 2 CH 4 (g) + 2 O 2 (g)  CO 2 (g) + 2 H 2 O(g)

13. Day 4 7.4 Balancing equations

14. 7.4- Writing Balanced Chemical Equations 1.Write an equation with the formula of each reactant and product. 2.Count the number of atoms of each element on each side of the equation.  Polyatomic ions may be counted as if they are one “element” if they appear as a reactant and product. 3.Pick an element to balance.  If an element is found in only once balance it first  Metals before nonmetals.  Leave elements that are free elements somewhere in the equation until last.

15. Writing Balanced Chemical Equations, Continued 4.Use the least common multiple as a coefficient in the equation.  If there is already a coefficient there, multiply it by the common multiple.  coefficients must go in front of entire molecules, not between atoms within a molecule. 5.Recount and repeat until balanced. If you get stuck- increase the number in front of the first reactant by one and repeat.If you get stuck- increase the number in front of the first reactant by one and repeat.

16. 16 Example: When magnesium metal burns in air, it produces a white, powdery compound magnesium oxide. Mg(s) + O 2 (g)  MgO(s) 3.Pick an element to balance.  Avoid element in multiple compounds.  Do free elements last.  Since Mg already balanced, pick O. 4.Find the LCM of both sides 5.and multiply each side by factor so it equals LCM.  LCM of 2 and 1 is 2. Mg(s) + O 2 (g)  MgO(s) 1  Mg  1 1  Mg  1 2  O  1 1 x x 2

17. Another Example: Under appropriate conditions at 1000°C, ammonia gas reacts with oxygen gas to produce gaseous nitrogen monoxide and steam 1.write the equation a)first in words  identify the state of each chemical ammonia(g) + oxygen(g)  nitrogen monoxide(g) + water(g) b)then write the equation in formulas  identify diatomic elements  identify polyatomic ions  determine formulas NH 3 (g) + O 2 (g)  NO(g) + H 2 O(g)

18. Practice #1 When aluminum metal reacts with air, it produces a white, powdery compound called aluminum oxide.  Reacting with air means reacting with O 2 : Aluminum(s) + oxygen(g)  aluminum oxide(s) Al(s) + O 2 (g)  Al 2 O 3 (s)

19. Practice #2 Acetic acid reacts with the metal aluminum to make aqueous aluminum acetate and gaseous hydrogen.  Acids are always aqueous.  Metals are solid except for mercury.

20. Day 5 7.5- Aqueous solutions and solubility 7.6- Precipitation Reactions

21. Aqueous Solutions  Many times, the chemicals we are reacting together are dissolved in water.  Mixtures of a chemical dissolved in water are called aqueous solutions (aq).  Dissolving the chemicals in water helps them to react together faster.  The water separates the chemicals into individual molecules or ions.  The separate, free-floating particles come in contact more frequently so the reaction speeds up.

22. Dissociation  When ionic compounds dissolve in water, the anions and cations are separated from each other. This is called dissociation.  However, not all ionic compounds are soluble in water!  When compounds containing polyatomic ions dissociate, the polyatomic group stays together as one ion.

23. Electrolytes  Electrolytes are substances whose water solution is a conductor of electricity.  All electrolytes have ions dissolved in water.

24. Electrolytes, Continued  In strong electrolytes, all the electrolyte molecules or formula units are separated into ions.  In nonelectrolytes, none of the molecules are separated into ions.  In weak electrolytes, a small percentage of the molecules are separated into ions.

25. Types of Electrolytes  Salts = Water soluble ionic compounds.  All strong electrolytes.  Acids = Form H +1 ions and anions in water solution.  In binary acids, the anion is monoatomic. In oxyacids, the anion is polyatomic.  Sour taste.  React and dissolve many metals.  Strong acid = strong electrolyte, weak acid = weak electrolyte.  Bases = Water-soluble metal hydroxides.  Bitter taste, slippery (soapy) feeling solutions.  Increases the OH -1 concentration.

26. When Will a Salt Dissolve?  A compound is soluble in a liquid if it dissolves in that liquid.  NaCl is soluble in water, but AgCl is not.  A compound is insoluble if a significant amount does not dissolve in that liquid.  AgCl is insoluble in water.  Though there is a very small amount dissolved, but not enough to be significant.

27. Compounds containing the following ions are generally soluble Exceptions (when combined with ions on the left the compound is insoluble) Li +, Na +, K +, NH 4 + none NO 3 –, C 2 H 3 O 2 – none Cl –, Br –, I – Ag +, Hg 2 2+, Pb 2+ SO 4 2– Ca 2+, Sr 2+, Ba 2+, Pb 2+ Solubility Rules: Compounds that Are Generally Soluble in Water (p.210)

28. Compounds containing the following ions are generally insoluble Exceptions (when combined with ions on the left the compound is soluble or slightly soluble) OH – Li +, Na +, K +, NH 4 +, Ca 2+, Sr 2+, Ba 2+ S 2– Li +, Na +, K +, NH 4 +, Ca 2+, Sr 2+, Ba 2+ CO 3 2–, PO 4 3– Li +, Na +, K +, NH 4 + Solubility Rules: Compounds that Are Generally Insoluble (p.210)

29. 7.6- Precipitation Reactions  Many reactions are done by mixing aqueous solutions of electrolytes together.  When this is done, often a reaction will take place from the cations and anions in the two solutions that are exchanging.  If the ion exchange results in forming a compound that is insoluble in water, it will come out of solution as a precipitate.

30. No Precipitate Formation = No Reaction KI(aq) + NaCl(aq)  KCl(aq) + NaI(aq) All ions still present,  no reaction.

31. Process for Predicting the Products of Precipitation Reaction 1.Write the formula for the reactants and Determine what ions each aqueous reactant has. 2.Exchange ions.  (+) ion from one reactant with (-) ion from the other. 3.Balance charges of combined ions to get formula of each product. 4.Balance the equation.  Count atoms. 5.Determine solubility of each product in water.  Use the solubility rules.  If product is insoluble or slightly soluble, it will precipitate.  If neither product will precipitate, no reaction.

32. Example: When an Aqueous Solution of Sodium Carbonate Is Added to an Aqueous Solution of Copper(II) Chloride, a White Solid Forms. 1.Write the formulas of reactants Na 2 CO 3 (aq) + CuCl 2 (aq)  (Na + + CO 3 2- ) + (Cu +2 + Cl - )  2.Exchange the ions. (Na + + CO 3 2- ) + (Cu +2 + Cl - )  (Na + + Cl - ) + (Cu +2 + CO 3 2- ) 3.Write the formulas of the products.  Cross charges and reduce. Na 2 CO 3 (aq) + CuCl 2 (aq)  NaCl + CuCO 3 4.Balance the equation. Na 2 CO 3 (aq) + CuCl 2 (aq)  NaCl + CuCO 3

33. Practice–Predict the Products and Balance the Equation  KCl(aq) + AgNO 3 (aq)   Na 2 S(aq) + CaCl 2 (aq) 

34.  KCl(aq) + AgNO 3 (aq)   KCl(aq) + AgNO 3 (aq)  KNO 3 (aq) + AgCl(s)  Na 2 S(aq) + CaCl 2 (aq)   Na 2 S(aq) + CaCl 2 (aq)  2 NaCl(aq) + CaS(aq)  No reaction. Practice–Predict the Products and Balance the Equation, Continued

35. Day 6 Quiz: 7.1-7.4 Solubility Lab

36. Day 7 7.7- Writing reactions in solutions

37. Ionic Equations  Equations that describe the chemicals put into the water and the product molecules are called molecular equations. 2 KOH(aq) + Mg(NO 3 ) 2 (aq)  2 KNO 3 (aq) + Mg(OH) 2 (s)  Equations that describe the actual dissolved species are called complete ionic equations.  Aqueous electrolytes are written as ions.  Soluble salts, strong acids, strong bases.  Insoluble substances and nonelectrolytes written in molecule form.  Solids, liquids, and gases are not dissolved, therefore, molecule form. 2K +1 (aq) + 2OH -1 (aq) + Mg +2 (aq) + 2NO 3 -1 (aq)  K +1 (aq) + 2NO 3 -1 (aq) + Mg(OH) 2(s)

38. Writing Complete Ionic Equations  Rewrite the molecular equation, but dissociate strong electrolytes into individual ions.  Strong electrolytes must be aqueous.  Solids, liquids, or gases cannot be electrolytes.  All soluble ionic compounds are strong electrolytes.  Strong acids are strong electrolytes.  HCl, HNO 3, H 2 SO 4..  Weak acids are not written in the dissociated ion form.  Molecular compounds do not have ions, leave in the molecular form.

39. Day 8 7.8- Acid/base and gas evolution reactions

40. Day 9 7.9- Oxidation-reduction reactions

41. Day 10 7.10- Classifying reactions Types of reactions activity

42. Days 11-13 Solutions Lab

43. Days 14&15 Review and Test