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Chemical Reactions

Chemical Reactions. Effects of chemical reactions:. Chemical reactions rearrange atoms in the reactants to form new products. The identities and properties of the products are completely different from that of the reactants.

chandler-sanders
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Chemical Reactions

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  1. Chemical Reactions

  2. Effects of chemical reactions: • Chemical reactions rearrange atoms in the reactants to form new products. • The identities and properties of the products are completely different from that of the reactants. • Production of gases and color changes are signs of chemical reactions.

  3. Energy and Reactions • Energy must be ADDED to BREAK bonds. • Energy is RELEASED when bonds are FORMED. • Chemical energy is CONSERVED in chemical reactions.

  4. Exo- vs. Endo- • EXOTHERMIC REACTIONS: release energy(More energy is released as the products form bonds than is absorbed to break the bonds in the reactants.) • ENDOTHERMIC REACTIONS: absorb energy

  5. Chemical Equations Chemical equations are used to represent or describe chemical reactions. For example when hydrogen H2 burns, it reacts with oxygen, O2, in the air to form water. We write the chemical equation for this reaction as follows: 2H2 + O2 —> 2H2O

  6. Chemical Equations An equation shows… Formulas of reactants Formulas of products Molar ratios of all compounds in the reaction.

  7. Chemical Equations We read the (+) sign as “reacts with” and the arrow (—>) as “produces” or “yields”. 2H2 + O2 —> 2H2O Reactants Products

  8. To show physical states of each substance: • (s) or  - solid • (l) - liquid • (g) or  - gas • (aq) - aqueous • aqueous means dissolved in water

  9. To show physical states of each substance: • Consider the reaction of iron with oxygen to form iron (III) oxide, or rust. • Fe(s) + O2(g)  Fe2O3(s) (unbalanced)

  10. Coefficients & Subscripts COEFFICIENTS: numbers in front of compound that represents the number of molecules/moles of that compound SUBSCRIPTS: small numbers that help define the compound. 2H2SO4 Coefficient Subscript

  11. H2O One molecule of water 2H2O Two molecules of water H2O2 One molecule of Hydrogen Peroxide

  12. During a chem. rxn.; atoms are rearranged (NOT created or destroyed!) • Chemical equations must be balanced to show the relative amounts of all substances. • Balanced means: each side of the equations has the same # of atoms of each element. • CH4 + O2 —> H2O + CO2Unbalanced • CH4 + 2O2 —> 2H2O + CO2 Balanced

  13. In order to balance… • Write correct formulas for all reactants and products • Reactants  Products • Count the number of atoms of each element in reactants & products. • Balance one at a time using coefficients. • Check for balance • Are the coefficients in the lowest possible ratio?

  14. Balancing Equations NOTE: When balancing equations, you may change coefficients as much as you need to, but you may never change subscripts because you can’t change what substances are involved.

  15. Balancing equations involves a great deal of “trial and error” at first, but there are some tricks…

  16. For example….. Sodium metal reacts with water to produce sodium hydroxide and hydrogen gas. Na + H2O —> NaOH + H2 Note that on the product side (right side) there are an odd number of hydrogens. On the reactant side (left side) there is an even number. This implies there must be an even coefficient in front of the NaOH. Lets start with 2

  17. _Na + _H2O —> 2NaOH + _H2 Now lets balance sodium; we need a 2 in front of the Na… 2Na + _H2O —> 2NaOH + _H2 Now consider hydrogen… 2Na + 2H2O —> 2NaOH + H2

  18. 2Na + 2H2O —> 2NaOH + H2 Check to see if it balances… 2 Na on the left 2 Na on the right 4 hydrogen 2 + 2 = 4 hydrogen 2 oxygen 2 oxygen the equation is balanced.

  19. Examples CuCl2(aq) + Al(s)  Cu(s) +AlCl3(aq) 3CuCl2(aq) + 2Al(s) 3Cu(s) +2AlCl3(aq) (3:2:3:2)

  20. Examples Propane, C3H8, burns in oxygen, O2, to form carbon dioxide and water. C3H8 + O2CO2 + H2O Balance C – then H – then O C3H8 + 5O23CO2 + 4H2O (1:5:3:4)

  21. Examples Pentane, C5H12, burns in oxygen, O2, to form carbon dioxide and water. C5H12 + O2CO2 + H2O Balance C – then H – then O C5H12 + 8O25CO2 + 6H2O (1:8:5:6)

  22. Examples Silver nitrate reacts with copper to produce silver and copper (II) nitrate. AgNO3 + Cu Ag + Cu(NO3)2 2AgNO3 + Cu 2Ag + Cu(NO3)2 (2:1:2:1)

  23. Examples Phosphorus reacts with oxygen gas to produce diphosphorus pentoxide. P + O2 P2O5 4P + 5O2 2P2O5 (4:5:2)

  24. Examples C7H14 + O2CO2 + H2O Balance C – then H – then O C7H14 + 10½O27CO2 + 7H2O 2C7H14 + 21O214CO2 + 14H2O (2:21:14:14)

  25. Types of Reactions

  26. Types of Chemical Reactions • Synthesis / Combination • Decomposition • Single Replacement • Double Replacement • Combustion

  27. Synthesis / Combination Reactions Definition: Reaction where two or more substances react to form a single substance. A + BAB Examples: 2K(s) + Cl2(g) 2KCl(s) SO2(g) + H2O(l) H2SO3(aq)

  28. Decomposition Reactions Definition: Reaction where a single compound is broken down into two or more products. AB A + B Examples: 2H2O(l) 2H2(g) + O2(g) CaCO3 CaO + CO2

  29. Single-Replacement Reactions Definition: Reaction where atoms of one element replace atoms of a second element in a compound. XA + B BA + X Note: A reactive metal will replace any metal listed below it in the activity series. Generally, nonmetal replacement is limited to the halogens. The activity of the halogens decreases as you go down Group 7A of the periodic table. See handout. Examples: 2AgNO3 + MgMg(NO3)2+2Ag Mg+LiNO3 no reaction

  30. Li K Ca Na Mg Al Zn Fe Pb (H)* Cu Hg Ag Activity Series Increasing Activity Any element will replace any element below it. *Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only

  31. For Example… Ca + MgO  CaO + Mg The Ca will replace the Mg because Ca is more active than Mg. That is to say…Ca is above Mg on the activity list.

  32. Double-Replacement Reactions Definition: Reaction that involves an exchange of positive ions between two compounds. XA + BYBA + XY Note: These reactions generally take place between two ionic compounds in aqueous solution, and are often characterized by one of the products coming out of solution in some way. Examples: 2NaCN(aq)+H2SO4(aq) 2HCN(g)+Na2SO4(aq) Na2S(aq)+Cd(NO3)2(aq) CdS(s)+2NaNO3(aq)

  33. Combustion Reactions Definition: Reaction where an element or a compound reacts with oxygen, often producing energy in the form of heat and light. Examples: CH4+2O2 CO2+2H2O + heat + light 2Mg(s)+O2(g) 2MgO(s)

  34. Combustion of Hydrocarbons If the reactant is a hydrocarbon, the products are always carbon dioxide and water. CH4 + 2O2  CO2 + 2H2O

  35. Ionic Equations • When a soluble substance is dissolved in water, the substance often breaks into ions. This solution is said to be an aqueous solution. • Pb(NO3)2(aq) Pb2+ + 2NO3- • NaI(aq)  Na+ + I-

  36. Ionic Equations • Consider the reaction… • Pb(NO3)2(aq) + NaI(aq) PbI2(s) + NaNO3(aq) • What is really going on is… • Pb2+ + NO3- + Na+ + I- PbI2(s) + Na+ + NO3- • Note that the Na+ ion and the NO3- ion are not reacting. They are said to be spectator ions.

  37. Net Ionic Equations • It is often useful to write an equation showing only the species that are actually reacting. This is called a net ionic equation. It does not show the spectator ions. Pb2+ + NO3- + Na+ + 2I- PbI2(s) + Na+ + NO3- becomes…. Pb2+ + 2I- PbI2(s)

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