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Understanding Chemical Reactions: Types, Equations, & Balancing

Learn about physical & chemical changes, reactions, balancing equations, and types of reactions in chemistry. Practice with problems and challenges to enhance your understanding.

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Understanding Chemical Reactions: Types, Equations, & Balancing

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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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