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Chapter 5 Chemical Reactions

Chapter 5 Chemical Reactions. CHE 101 Sleevi. Chemical Reactions. The process of chemical change Substances are transformed to different substances Indicators: formation of precipitate unexpected color change evolution of gas release or absorption of energy. Chemical Reaction.

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Chapter 5 Chemical Reactions

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  1. Chapter 5Chemical Reactions CHE 101 Sleevi

  2. Chemical Reactions • The process of chemical change • Substances are transformed to different substances • Indicators: • formation of precipitate • unexpected color change • evolution of gas • release or absorption of energy

  3. Chemical Reaction A + B  C + D

  4. The Language of Reactions • Reactant • starting material of a chemical reaction • appears to the left of the reaction arrow • Product • substance formed in chemical reaction • appears to the right of the reaction arrow • Reaction Conditions • solvent, energy applied, catalysts, etc.

  5. Chemical Equations Unbalanced chemical equation: Na (s) + H2O (l) H2(g) + NaOH (aq) s = solid l = liquid g = gas aq = aqueous solution

  6. Reactions Conditions Heat = Δ Electricity = e- Light = hv or γ Catalyst = Pt Specific Temperature = 50oC

  7. Chemical Reaction • Calcium carbonate is heated strongly and carbon dioxide gas is driven off, leaving a residue of calcium oxide. CaCO3(s) CO2(g) + CaO (s)

  8. Descriptions of Chemical Reactions • Identify reactants by language such as: • is heated strongly • decomposes • is combined with • when added to • reacts with • neutralizes • x is converted to…

  9. Descriptions of Chemical Reactions • Identify products by language such as: • is formed • produced • precipitates • is given off • is evolved • leaving a residue of • is converted to y

  10. Writing Chemical Equations from Descriptions of Reactions • Identify reactants and products • Write correct chemical formulas • Identify and document states of each substance • Record reaction conditions (if given) above/below reaction arrow

  11. Elements as Diatomic Molecules • Seven elements occur as diatomic molecules

  12. Examples • Bubbling chlorine gas through a solution of potassium iodide gives elemental iodine and a solution of potassium chloride. • Solid silver oxide can be heated to give silver and oxygen gas.

  13. Balancing Chemical Reactions • Law of Conservation of Matter: • matter is neither created nor destroyed • Representation of chemical reactions must account for the laws of nature… • the number and type of each atom in the reactants and products is the same on both sides of the equation, regardless of the initial or final substances

  14. Why Equations are Balanced • Conservation of matter/mass • # atoms of each element same on both sides of equation • Mass of reactants = mass of products • Number of moles or particles NOT conserved only mass and matter • Type and number of each atom, not the arrangement of them

  15. Balancing Chemical Equations • Chemical reactions occur on a particle basis H2 + O2 = H2O Molecules of hydrogen and oxygen react to form molecules of water. Need to balance the equation to satisfy law of conservation of matter  atoms are conserved!

  16. Balancing Chemical Equations • Balanced chemical reactions represent the proportion in which substances undergo chemical change • Coefficients represent lowest whole number ratio of substances involved in the reaction

  17. Interpreting Balanced Chemical Equations 2 H2 + O2 = 2 H2O • two molecules of H2 and one molecule of O2 combine to form 2 molecules of water OR • two moles of H2 and one mole of O2 combine to form 2 moles of water

  18. Steps to Balancing Equations • Write the unbalanced equation that describes the reaction (correct formulas!) • Use coefficients to balance the number of atoms of each element on both sides • Start with elements that appear only once on each side of the equation

  19. Steps to Balancing Equations • For elements found in two substances on one side of the equation add the two to ensure the number on the other side of the equation is correct • Reduce coefficients to lowest whole number ration • Double check all elements when done.

  20. Tips for Balancing Equations… • Never alter a correct formula in a chemical equation to balance the equation. • Coefficients go in front of the formula and multiply each subscript of each element in the compound.

  21. Tips for Balancing Equations… • Polyatomic ions that occur on both sides of the equation can be balanced as ions rather than balancing individual atoms. • If there is an odd number of atoms of an element on one side of the reaction and an even number on the other side before balancing, both sides will need to have an even number coefficient to balance the equation.

  22. Examples Al + O2 Al2O3 NaCl + H2SO4  Na2SO4 + HCl H2 + Fe2O3  Fe + H2O C4H10 + O2  CO2 + H2O

  23. Driving Forces for Chemical Reactions • Formation of a solid • Formation of water • Formation of a gas • Transfer of electrons (Reactions are spontaneous if the products are favored)

  24. Types of Chemical Reactions • Decomposition • Combination • Single Replacement • Double Replacement • Complete Combustion of Hydrocarbons

  25. Decomposition Reactions • Single substance broken down into two or more simpler substances • Reactant must be compound • Products can be elements or compounds • Require input of energy to occur (light, heat, electricity) A  B + C

  26. Decomposition Reactions 2H2O  2H2 + O2 2H2O2  2H2O + O2 K2CO3  K2O + CO2 2KOH  K2O + H2O

  27. Decomposition Reactions • Binary compound decomposes to elements • Metal carbonate decomposes to metal oxide and carbon dioxide • Base decomposes to metal oxide and water

  28. Predict the Products OF2 NaOH  NiCO3  FeCl3 

  29. Combination Reactions • Two or more substances form one product • Reactants can be elements or compounds • Product is always a compound A + B  C

  30. Combination Reactions Na (s) + Cl2 (g) NaCl (s) SO3(g) + H2O (l)  H2SO4(aq) K2O (s) + H2O (l)  KOH (aq)

  31. Combination Reactions • Two elements combine to form a binary compound • metal + nonmetal  ionic compound • nonmetal + nonmetal  molecular compound • Nonmetal oxide + water  acid • Metal oxide + water  base

  32. Predict the Products Be + O2 CO2 + H2O  Mg + N2 CuO + H2O 

  33. Single Replacement Reactions • Substitution reactions in which an element replaces the element in an ionic compound • metal replaces metal • halogen replaces halogen • metal replaces hydrogen (in an acid) • Not all reactions occur A + BX  AX + B

  34. Single Replacement Reactions Mg + ZnCl2  MgCl2 + Zn Fe + CuSO4  FeSO4 + Cu Na + H2O  NaOH + H2 F2 + KCl  KF + Cl2

  35. Single Replacement Reactions • Activity Series provides reference for which single replacement reactions occur

  36. Predict the Products Zn + CuSO4 Mg + PbCl2 Na + H2O 

  37. Double Replacement Reactions • Exchange of positive ions between two compounds • Usually occur between ionic compounds in aqueous solutions • When reaction occurs • a precipitate forms • water or other molecular compound is formed • evolution of a gas

  38. Double Replacement Reactions Na2S (aq) + Cd(NO3)2(aq) NaNO3 (aq) + CdS (s) NaCN (aq) + H2SO4(aq)  HCN (g) + Na2SO4(aq) AgNO3(aq) + KCl (aq) AgCl (s) + KNO3(aq) NaOH (aq) + H2SO4(aq) Na2SO4(aq) + H2O (l)

  39. When Will DR Reactions Occur? • Look at solubility of expected products • If cations are the same (or both cations are 1A metals) – NR • If anions are the same – NR • Nitrates and Alkali metal salts are soluble • Look at reactants – acid + base neutralization is special case of DR

  40. Predict the Products HCl + Ca(OH)2 Ag2SO4 + AlCl3 BaBr2 + K2CO3 

  41. Ways to Count

  42. Ways to Count • Each of the relationships on the previous page can be written as an equality and used to create conversion factors • These relationships are common place and we perform these calculations in our head, hardly thinking about it

  43. Ways to Count Class Exercise

  44. Ways to Count Methane Molecules

  45. Counting Atoms and Molecules • Atoms and molecules are very small • Reactions happen at the particle level so number of atoms and molecules is important • Term introduced to represent a specific quantity of atoms or molecules • The chemists “dozen”

  46. The Mole 1 mole = 6.02 x 1023 particles Avogadro’s Number (N)

  47. How big is a mole? • a mole of…. • sand grains would cover the US in ~ 1 cm of sand • cells in the human body is ~1 mole • grapefruit would have approximately the same volume as the Earth (~500 large grapefruit will fit in 1 cubic meter) • dollars would take 20 billion years to spend at 1 million dollars a second data from Wikipedia 12/28/06, updated 12/22/06 and Flinn ChemFax, 10/27/03

  48. Mass Ratios

  49. Conclusion • Average atomic mass of any element in grams is the mass of one mole of that element 1 mole C = 12.01 g 1 mole Mg = 24.31 g 1 mole S = 32.06 g

  50. Terms • Synonyms: • gram atomic mass • gram atomic weight • molar mass • Definitions: • Gram atomic mass is the mass, in grams, of one mole of the element • Molar mass is the mass, in grams, of one mole of a substance

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