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Chapter 11

Chapter 11. Chemical Reactions. Chemical Equation. Describes chemical reaction. Chemical equation: reactants yield products Reactants  Products Much easier to write symbols and formulas instead of words. Examples. Solid Iron reacts with oxygen gas to form the solid IronIIIoxide .

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Chapter 11

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  1. Chapter 11 Chemical Reactions

  2. Chemical Equation • Describes chemical reaction. • Chemical equation: reactants yield products • Reactants  Products • Much easier to write symbols and formulas instead of words

  3. Examples Solid Iron reacts with oxygen gas to form the solid IronIIIoxide. iron(s) + oxygen(g)  ironIIIoxide(s) Fe(s) + O2(g)  Fe2O3(s) Carbon tetrahydride gas BURNS to form carbon dioxide gas and water vapor. Carbon tetrahydride(g) + oxygen(g)  carbon dioxide(↑) + water(↑) CH4(g) + O2(g)  CO2(↑) + H2O(↑) Skeleton Equation: chemical equation that tells you what the reactants and products are but NOT how much of each you have. First step in writing a chemical equation.

  4. Symbols Used • (s) solid • (l) liquid • (g) gas • (↑) gas as a product • (aq) aqueous (in water solution) • ()ppt(precipitate)solid product from 2 aqueous reactants • D means with heat • Pt means with Platinum catalyst: speeds up a reaction without being used. •  reversible reaction

  5. Balancing Chemical Equations • Balanced equations have: • the same # of atoms of each element on BOTH sides of the equation. • Law of Conservation of Mass – atoms can neither be created nor destroyed, simply rearranged.

  6. Rules for Balancing Equations • Get the correct formulas for reactants and products. • (USE ION CHART AND DON”T FORGET DIATOMIC ELEMENTS!) • Write reactants on left, products on right. Use plus signsto separate compounds and yield sign to separate the reactants from products.

  7. Rules Continued Count the # of atoms of each element in reactants and products. (Polyatomic atoms on both sides count as one.) Balance # of each element using coefficients. Coefficient – small whole # in front of a formula. NEVER CHANGE FORMULA SUBSCRIPTS

  8. Rules for Balancing Equations • Balance elements appearing 3 or more places LAST. • Check each element to make sure equation is balanced. • Make sure all coefficients are in the lowest whole number ratio. Do not change subscripts!!!

  9. Diatomic Molecules Diatomic Molecules- a molecule made up two atoms of the same element. They are only diatomic when they are alone. • There are 7 naturally occurring • diatomic molecules. HONClBrIF

  10. Balancing Examples ___ C(s) + ___ O2(g)  ___ CO2 (g) ___ C(s) + ___ O2(g)  ___ CO (g) ___ AgNO3(aq) + ___Cu(s)  ___ Cu(NO3)2(aq) + ___ Ag(s) ___ Al(s) + ___ O3(g)  ___ Al2O3(s) *___ C2H6(g) + ___ O2(g)  ___ CO2(g) + ___ H2O(g) *___ H3PO3  ___ H3PO4 + PH3

  11. 5 Types of Chemical Reactions • Combination Reaction – elements combine to form a compound. • A + B AB • element + element  compound • Ex. Sodium + chlorine  sodium chloride • ___Na(s) + ___ Cl2(g)  ___ NaCl(s) 2 • 2

  12. 5 Types of Reactions • Decomposition Reaction – compound breaks down into its element. • AB A + B • compound  element + element • Ex: MercuryII oxide  mercury + oxygen • ___ HgO  ___Hg + ___O2 2 • 2

  13. 5 Types of Reactions - 3 Single Replacement Reaction – one element replaces another element in a compound. AB + C A + CB or AB + D AD + B + - + + + - + - - + - -

  14. Examples of Single Replacement Reactions • Must use Activity Series to see if reaction works • Zinc + sulfuric acid  zinc sulfate + hydorgen • Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(↑) • Periodic table is activity series for halogens • Sodium bromide + chlorine  sodium chloride + bromine • ___NaBr(s) + ___Cl2(g)  ___NaCl(s) + ___Br2(↑) 2 • 2

  15. 5 Types of Reactions Double Replacement Reaction – two compounds react and exchange positive ions to form two new compounds. AB + CD AD + CB • Barium Chloride(aq) + potassium carbonate(aq)  barium carbonate() + potassium chloride(aq) • BaCl2(aq) + K2CO3(aq)  BaCO3() + ___ KCl(aq) + - + - + - + - 2

  16. 5 Types of Reactions Combustion Reaction – oxygen reacts with a compound composed of C and H. CxHy + O2 CO2 + H20 Also called Burning (exothermic) The products are always CO2 and H2O.

  17. Examples of Combustion Reactions 1. C6H6 + O2 CO2 + H2O 7½ 6 3 2 15 12 6 C6H6 + O2 CO2 + H2O 2. CH3OH + O2 CO2 + H2O 1½ 2 2 3 2 4 CH3OH + O2 CO2 + H2O

  18. Special Decomposition Reactions: • Decomposition of a Carbonate: • Metal carbonate  metal oxide + carbon dioxide XCO3 XO + CO2 ex. Na2CO3 Na2O + CO2

  19. Special Decomposition Reactions: • Decomposition of a Hydroxide: • Metal hydroxide  metal oxide + water XOH XO + H2O ex. 2NaOH Na2O + H2O

  20. Special Decomposition Reactions: • Decomposition of a Chlorate: (ClO3) • Metal chlorate  metal chloride + oxygen • XClO3XCl + O2 ex. ___NaClO3 ___NaCl+ ___O2 2 • 3

  21. Special Decomposition Reactions: 4 • Special single Replacement Reaction: • Group IA or IIA metal and H2O X + HOH XOH + H2 ex. 2Na + 2HOH 2NaOH + H2

  22. How to ID types of reactions. Combination Reactions – given 2 items that form 1 new compound. Decomposition Reactions – given a single compound that breaks into parts. Single Replacement – given a single element plus a single compound, forms a new compound a a different element. Double Replacement – given two compounds (+’s change places). Combustion Reaction – given CH compound with Oxygen, always forms water and carbon dioxide.

  23. The End

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