1. Unit 6 Chemical Reactions and Equations Slides adapted from Nivaldo Tro

2. 2 Experiencing Chemical Change chemical reactions are happening both around you and in you all the time some are very simple, others are complex in terms of the pieces – even the simple ones have a lot of interesting principles to learn from chemical reactions involve changes in the structures of the molecules, and many times we can experience the effects of those changes What are some examples of chemical reactions you experience?

3. 3 Chemical Reactions Reactions involve chemical changes in matter resulting in new substances Reactions involve rearrangement and exchange of atoms to produce new molecules Elements are not transmuted during a reaction Reactants  Products

4. 4 Evidence of Chemical Reactions look for evidence of a new substance visual clues (permanent) color change precipitate formation  solid that forms when liquid solutions are mixed gas bubbles large energy changes  container becomes very hot or cold  emission of light other clues new odor whooshing sound from a tube permanent new state

5. 5 Evidence of Chemical Change Color Change Formation of Solid Precipitate Formation of a Gas Emission of Light Release or Absorption of Heat

6. 6 Evidence of Chemical Change in order to be absolutely sure that a chemical reaction has taken place, you need to go down to the molecular level and analyze the structures of the molecules at the beginning and end Is boiling water a chemical change?

7. 7 Chemical Equations Shorthand 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 weights of reactants used and products that can be made

8. 8 Conservation of Mass Matter cannot be created or destroyed Therefore the total mass cannot change 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 if all the atoms are still there, then the mass will not change

9. 9 The Combustion of Methane methane gas burns to produce carbon dioxide gas and gaseous water whenever something burns it combines with O 2 (g)

10. 10 Combustion of Methane methane gas burns to produce carbon dioxide gas and gaseous water whenever something burns it combines with O 2 (g) CH 4 (g) + O 2 (g)  CO 2 (g) + H 2 O(g) H H C H H OO + O O C + O HH 1 C + 4 H + 2 O1 C + 2 O + 2 H + O 1 C + 2 H + 3 O

11. 11 Combustion of Methane Balanced to show the reaction obeys the Law of Conservation of Mass it must be balanced CH 4 (g) + 2 O 2 (g)  CO 2 (g) + 2 H 2 O(g) H H C H H + O O C + OO OO + O HH O HH + 1 C + 4 H + 4 O

12. 12 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)

13. 13 Chemical Equations 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 CH 4 (g) + 2 O 2 (g)  CO 2 (g) + 2 H 2 O(g)

14. 14 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

15. 15 Writing Balanced Chemical Equations 1.Write a skeletal equation by writing 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 often be counted as if they are one “element” 3.Pick an element to balance if an element is found in only one compound on both sides, balance it first leave elements that are free elements somewhere in the equation until last

16. 16 Writing Balanced Chemical Equations 4.Find the Least Common Multiple of the number of atoms on each side the LCM of 3 and 2 is 6 5.Multiply each count by a factor to make it equal to the LCM 6.Use this factor as a coefficient in the equation if there is already a coefficient there, multiply it by the factor it must go in front of entire molecules, not between atoms within a molecule 7.Recount and Repeat until Balanced

17. 17 Examples when magnesium metal burns in air it produces a white, powdery compound magnesium oxide Mg(s) + O 2 (g)  MgO(s)

18. 18 Examples when magnesium metal burns in air it produces a white, powdery compound magnesium oxide 2Mg(s) + O 2 (g)  2MgO(s)

19. 19 Examples Under appropriate conditions at 1000°C ammonia gas reacts with oxygen gas to produce gaseous nitrogen monoxide and gaseous water NH 3 (g) + O 2 (g)  NO(g) + H 2 O(g)

20. 20 Examples Under appropriate conditions at 1000°C ammonia gas reacts with oxygen gas to produce gaseous nitrogen monoxide and gaseous water 2(2NH 3(g) + 2.5O 2(g)  2NO (g) + 3H 2 O (g) ) 4NH 3(g) + 5O 2(g)  4NO (g) + 6H2O (g)

21. 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 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. 22 Predicting Whether a Reaction Will Occur in Aqueous Solution “forces” that drive a reaction 1)formation of a solid 2)formation of water 3)formation of a gas 4)transfer of electrons when chemicals (dissolved in water) are mixed and one of these 4 things can occur, the reaction will generally happen

23. 23 Classifying Reactions one way is based on the process that happens precipitation, neutralization, formation of a gas, or transfer of electrons

24. 24 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 exchanging if the ion exchange results in forming a compound that is insoluble in water, it will come out of solution as a precipitate

25. 25 Precipitation Reactions Pb(NO 3 ) 2 (aq) + 2 KI(aq)  2 KNO 3 (aq) + PbI 2 (s)

26. 26 Precipitation Reactions Pb(NO 3 ) 2 (aq) + 2 KI(aq)  2 KNO 3 (aq) + PbI 2 (s)

27. 27 No Precipitate Formation = No Reaction KI(aq) + NaCl(aq)  KCl(aq) + NaI(aq) all ions still present,  no reaction

28. 28 Process for Predicting the Products of a Precipitation Reaction 1.Determine what ions each aqueous reactant has 2.Exchange Ions (+) ion from one reactant with (-) ion from 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

29. 29 Acid-Base Reactions also called neutralization reactions because the acid and base neutralize each others properties in the reaction of an acid with a base, the H +1 from the acid combines with the OH -1 from the base to make water the cation from the base combines with the anion from the acid to make the salt *acid + base  salt + water 2 HNO 3 (aq) + Ca(OH) 2 (aq)  Ca(NO 3 ) 2 (aq) + 2 H 2 O(l) the net ionic equation for an Acid-Base reaction is H +1 (aq) + OH -1 (aq)  H 2 O(l) as long as the salt that forms is soluble in water

30. 30 Gas Evolving Reactions Some reactions form a gas directly from the ion exchange K 2 S(aq) + H 2 SO 4 (aq)  K 2 SO 4 (aq) + H 2 S(g) Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water K 2 SO 3 (aq) + H 2 SO 4 (aq)  K 2 SO 4 (aq) + H 2 SO 3 (aq) H 2 SO 3  H 2 O(l) + SO 2 (g)

31. 31 Oxidation-Reduction Reactions We say that the element that loses electrons in the reaction is oxidized and the substance that gains electrons in the reaction is reduced you cannot have one without the other

32. 32 Combustion Reactions Reactions in which O 2 (g) is a reactant and heat is needed initially are called Combustion Reactions Combustion reactions release a significant amount of energy Combustion reactions are a subclass of Oxidation-Reduction reactions 2 C 8 H 18 (g) + 25 O 2 (g)  16 CO 2 (g) + 18 H 2 O(g)

33. 33 Combustion Products to predict the products of a combustion reaction, combine each element in the other reactant with oxygen Remember: These reactions all need heat to occur ReactantCombustion Product *contains CCO 2 (g) *contains HH 2 O(g) contains SSO 2 (g) contains NNO(g) or NO 2 (g) contains metalM 2 O n (s)

34. 34 Classifying Reactions another scheme classifies reactions by what the atoms do Type of ReactionGeneral Equation Synthesis A + B  AB Decomposition AB  A + B Displacement A + BC  AC + B Double Displacement AB + CD  AD + CB

35. Rules for Predicting Reactions 1)Form your products first  Hint: DO NOT write any numbers when doing this step, unless it is part of a polyatomic ion 2)Balance the charges Make sure the charges cancel out  Hint: Use polyatomic charges handout for help with charges 3)Balance your equation Follow rules for balancing (make sure diatomics are written as H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 )  Hint: Balance oxygen last, followed by hydrogen 4)Write in the state of matter (s), (l), (g), (aq)  Hint: use solubility chart 35

36. 36 Synthesis Reactions also known as Combination reactions two (or more) reactants combine together to make one product 2 form 1 simpler substances combining together 2 CO + O 2  2 CO 2 2 Mg + O 2  2 MgO

37. Hints for Synthesis Reactions Common synthesis examples Metal Oxide + H 2 O → Metal Hydroxide CaO + H 2 O →Ca(OH) 2 Metal Oxide + CO 2 → Metal Carbonate CaO+ CO 2 → CaCO 3 Covalent Compound(nonmetals) + H 2 O→ Acid SO 3 + H 2 O → H 2 SO 4 Metal + Nonmetal → Ionic Compound 2Na + Cl 2 →2NaCl 37

38. 38 Decomposition Reactions a large molecule is broken apart into smaller molecules or its elements caused by addition of energy into the molecule 1 forms 2 have only one reactant, make 2 or more products Example: 2FeCl 3  2FeCl 2 + Cl 2

39. 39 Decomposition of Water

40. Hints for Decomposition Reactions Common decomposition examples Metal Hydroxide →Metal Oxide + H 2 O Ca(OH) 2 →CaO + H 2 O Metal Carbonate →Metal Oxide + CO 2 CaCO 3 →CaO+ CO 2 Acid →Covalent Compound(nonmetals) + H 2 O H 2 SO 4 → SO 3 + H 2 O Ionic Compound →Metal + Nonmetal 2NaCl →2Na + Cl 2 40

41. 41 Displacement Reactions also known as single-replacement reactions X  Y  + A   X + A  Y  Examples of cation exchange: metals in general or H+ Zn(s) + 2 HCl(aq)  ZnCl 2 (aq) + H 2 (g) 2Na(s) + H 2 O(aq) → Na 2 O(aq) + H 2 (g) Examples of anion exhange: nonmetals Br 2 + MgI 2 → MgBr 2 + I 2 Cl 2 + 2KI → 2KCl + I 2

42. 42 Displacement of Copper by Zinc

43. 43 Double Displacement Reactions two ionic compounds exchange ions Either cation or anion exchange  Hint:It doesn’t matter in this case; just don’t swith both the cations and anions X  Y  (aq) + A  B  (aq)  XB + AY precipitation, acid-base and gas-evolving reactions are also double displacement reactions Example: 2AlBr 3(aq) + Ca 3 (PO 4 ) 2(s) → 2AlPO 4(s) + 3CaBr 2(aq) The Al and the Ca switch places

44. Rates of a Reaction 4 ways to speed up the rate of a reaction Based on the kinetic molecular theory 1.Increase the concentration (think more kool aid powder in your kool aid) 2.Increase the surface area (think little twigs instead of big logs when starting a fire)) 3.Increase the temperature 4.Add a catalyst (not used up) 44