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Chapter 7

Chapter 7. Chemical Reactions. Homework. Assigned Problems (odd numbers only) “Questions and Problems” 7.1 to 7.31 (begins on page 200) “Additional Questions and Problems” 7.41 to 7.49 (page 221) “Challenge Questions” 7.51-7.57 (page 222). Chemical Reactions. Physical changes:

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Chapter 7

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  1. Chapter 7 Chemical Reactions

  2. Homework • Assigned Problems (odd numbers only) • “Questions and Problems” 7.1 to 7.31 (begins on page 200) • “Additional Questions and Problems” 7.41 to 7.49 (page 221) • “Challenge Questions” 7.51-7.57 (page 222)

  3. Chemical Reactions • Physical changes: • Involves no changes in chemical identity of a substance • No changes in physical properties (color, physical state, freezing point, boiling point) • Chemical changes: • A chemical reaction in which one or more substances changes to a different substance • Properties that matter exhibits as it undergoes changes in chemical composition

  4. Chemical Reactions • Chemical properties determine whether or not a substance can be changed to another substance • Reactions involve chemical changes in matter resulting in new substances • Reactions involve rearrangement and exchange of atoms to produce new molecules • Elements are not changed during a reaction Reactants  Products

  5. Changes During Chemical Reactions • A chemical change occurs when new substances are made • Conversion of material(s) into one or more new substances • These substances will have different properties from the original material • New properties are visible (visual clues) • Color change, precipitate formation, gas bubbles, flames, heat release

  6. Changes During Chemical Reactions Fe Fe2O3 Li LiOH, H2 HCO3-CO2 Na NaOH, H2

  7. Changes During Chemical Reactions • In a chemical reaction: • At least one new substance is produced • Atoms are never created or destroyed • Every atom present as a reactant has to be present as a product • The atoms in reactants rearrange to form new products

  8. Chemical Equations • A chemical equation is a written statement that uses symbols and formulas (no words) to describe the changes during a chemical reaction • It shows substances at the beginning of a reaction (reactants) • It shows substances formed in the reaction (products)

  9. Writing a Chemical Equation • Chemical reactions can be written as: • Word equations • Formula equations reactants products

  10. Balancing Chemical Equations • A balanced chemical reaction has the same number of atoms of each element on both sides of the arrow • Atoms are neither created nor destroyed • Every atom must be accounted for • Equations are balanced by placing a coefficient in front one or more of the substances in the equation

  11. Symbols Used in Equations • Symbols used after chemical formula to indicate physical state • (g) = gas • (l) = liquid • (s) = solid • (aq) = aqueous, dissolved in water

  12. Writing Chemical Equations • When magnesium metal burns in air it produces a white, powdery compound magnesium oxide • Burning in air means reacting with O2 • Write the word equation • The reactants are to the left of the arrow • The products are to the right of the arrow • Two or more reactants or products are separated by a plus sign magnesium + oxygen magnesium oxide

  13. Writing Chemical Equations • Indicate the physical state of each substance • Use the correct chemical symbol to indicate liquids and solids • Metals are solids, except for Hg which is liquid • Use molecular form for gases (H2, O2, N2, all halogens) • Identify polyatomic ions magnesium(s) + oxygen(g) magnesium oxide(s)

  14. Writing Chemical Equations • Convert the word equation into a formula equation • Use the correct chemical symbol to indicate liquids and solids • There must be the same number of each kind of atom on the reactant and product side of the equation • Determine if the equation is balanced • If not equal, must BALANCE ___Mg (s) +___O2 (g) ___MgO(s)

  15. Balancing Chemical Equations • Balance equations by the use of a coefficient placed to the left of a substance • NEVER change the subscripts of a compound to balance an element • It changes the identity of the compound • Can change coefficients but never subscript numbers

  16. 2 2 2 Balancing a Chemical EquationExample 1 2 2 Coefficient 1 Mg 1 Mg 2 O 1 O

  17. Balancing a Chemical Equation: Example 2 • When solid ammonium nitrite is heated it produces nitrogen gas and water vapor • Write the formula equation

  18. 2 4 Balancing a Chemical Equation Example 2 2 2 x N 2 x N 2 x O 1 x O 4 x H 2 x H

  19. Balancing a Chemical Equation: Example 3 • Nitrogen monoxide gas decomposes to produce dinitrogen monoxide gas and nitrogen dioxide gas • Write the formula equation

  20. 3 3 Balancing a Chemical Equation Example 3 3 1 x N 3 x N 1 x O 3 x O

  21. Balancing a Chemical Equation Example 4 • Liquid nitric acid decomposes to reddish-brown nitrogen dioxide gas, liquid water and oxygen gas. • Write the formula equation

  22. 4 4 4 2 2 4 4 7 6 12 12 4 2 Balancing a Chemical Equation Example 4 2 2 2 1 x N 1 x N 3 x O 5 x O 1 x H 2 x H

  23. Types of Reactions • Reactions are separated into groups of similar reactions • Based on the form of the equation for the reaction • Synthesis (combination) • Decomposition • Single replacement • Double replacement • Combustion

  24. Types of Reactions • Synthesis Reactions • Reactions in which two or more substances combine to form a third substance • (one product forms) • General form of equation: A + B AB

  25. Synthesis Reactions • The combinations can include • Two elements • An element and a compound • Two compounds • Examples

  26. Types of Reactions • Decomposition Reactions • Reactions in which one reactantbreaks down into simpler (smaller) substances • Generally initiated by addition of energy (electric current or heating substances to high temperature) • Opposite of a Synthesis Reaction • General Form of Equation A + B AB

  27. Decomposition Reactions • Can be broken down to: • Smaller compounds • Elements • Both • Examples

  28. Types of Reactions • Single replacement reactions • One element replaces another element • Forms a new compound which frees the replaced element • Most reactions occur in an aqueous solution • General Form of Equation A + BC AC + B

  29. Single Replacement Reactions • Three types • Metal replaces a metal • Metal replaces hydrogen • Nonmetal replaces nonmetal • Examples metal metal replaces metal replaces hydrogen nonmetal nonmetal replaces

  30. Type of Reactions • Two compounds exchange ions or atoms to form new compounds • Also called exchange reactions • Shows the exchange of “associates” when comparing the reactants and products • General Form of Equation AB + CD AD + BC

  31. Double Replacement Reactions • Most of these reactions occur in aqueous solution • Most involve acids, bases, and ionic compounds • Products formed • Precipitate (a solid that is insoluble) • A gas • Water

  32. Double Replacement Reactions • Examples precipitate gas water

  33. Summary of Reaction Types

  34. Combustion Reactions • Occurs when a hydrocarbon combines with oxygen which produces carbon dioxide, water and heat (flame) • The reaction of oxygen with any substance • If a combustion reaction is possible then the substance will burn

  35. Combustion Reactions • Examples • The combustion of propane gas • Produces carbon dioxide and water • Produces heat (flame) • The combustion of sulfur • Also a combination reaction • Also produces heat (flame) hydrocarbon

  36. Energy in Chemical Reactions • In a chemical reaction • A change in energy occurs as bonds are broken (reactants) and new ones form (products) • Nearly all chemical reactions absorb or produce heat • Measured by the heat of reaction or enthalpy • Enthalpy change is the amount of heat produced or consumed in a process (∆H )

  37. Heat of Reaction • Endothermic reactions absorb heat as they occur • If (∆H ) is positive, then heat is added to the reaction • Exothermic reactions produce heat as they occur • If (∆H ) is negative, then heat is evolved by the reaction

  38. Heat of Reaction • Photosynthesis reaction • Carbon dioxide reacts with water to produce glucose and oxygen • Cell metabolism • Glucose reacts with oxygen to produce carbon dioxide and water ∆H = +2801 kJ ∆H = -2801 kJ

  39. Calculation of Heat in Reactions • The combustion of sulfur dioxide • It reacts with oxygen to produce sulfur trioxide • Calculate the heat produced when 75.2 g of sulfur trioxide is produced ∆H = -99.1 kJ Given 75.2 g SO3 Heat in kJ produced when SO3 is formed

  40. Calculation of Heat in Reactions Relation between g of SO3 and heat released Grams of SO3 Heat of rxn Moles of SO3 Molar mass kj Write the necessary conversion factors Set up the problem 46.5 kJ

  41. End

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