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Chemical Equations

Chemical Equations. Chp 7 and 8. Chemical and Physical Changes. Physical Change No new substance is formed Can often be reversed or undone Ex. Melting, freezing, crushing, etc Chemical Change A new substance is produced as compounds switch atoms Usually are permanent

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Chemical Equations

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  1. Chemical Equations Chp 7 and 8

  2. Chemical and Physical Changes • Physical Change • No new substance is formed • Can often be reversed or undone • Ex. Melting, freezing, crushing, etc • Chemical Change • A new substance is produced as compounds switch atoms • Usually are permanent • Ex. Cooking, decaying, rusting, etc

  3. Evidence of a Chemical Change • Color change • Bubbles • Solid forms • Flames • Heat • Other visual cues

  4. The Chemical Equation • A written representation of a chemical change • Reactants – what compounds/atoms exist before the reaction (before the arrow) • Products – what compounds/atoms exist after the reaction has occurred (after the arrow) • Subscripts – lowerset # after an atom, indicates the amount of that atom present in the compound • Coefficient – large # in front of a compound, indicates that amount of that molecule needed for the chemical reaction to occur

  5. Conservation of Matter • Atoms cannot be created or destroyed, they can only switch between compounds • That means, whatever types and #s of atoms are present before a reaction occurs (reactants) must equal the types and #s that are present afterward (products) • Or, equations must be balanced

  6. Balancing Equations • Place coefficients in front of compounds to make all atoms equal on each side of the equation (conserve the matter) • The coefficient multiplies by the subscript of each atom within that compound • You may NEVER change a subscript when balancing, that would change the compound • Ex. H2O is water, H2O2 is hydrogen peroxide

  7. An Example ____ H2 + _____ O2 ____ H2O ReactantsProducts H = 2 H = 2 O = 2 O = 1

  8. An Example, Continued ____ H2 + _____ O2 2 H2O ReactantsProducts H = 2 H = 2 4 O = 2 O = 1 2

  9. An Example 2 H2 + _____ O2 2 H2O ReactantsProducts H = 2 4 H = 2 4 O = 2 O = 1 2

  10. Tips • If an atom appears in more than one compound on the same side, leave it until the end (it will often work itself out) • Start with atoms that have an odd number of one side and an even number on the other • Use a pencil (it’s okay to go back and erase when things don’t work out) • If you get to the end, and everything but one atoms works, try multiplying all the existing coefficients by 2.

  11. Types of Reactions • Synthesis – 2 reactants (or more) form 1 product • A + B C • Decomposition – 1 reactant forms 2 products (or more) • C A + B • Single Displacement – 1 lone element on the reactant side combines with part of a compound, leaving another element alone • A + BC AC + B

  12. Types of Reactions, Cont’d • Double Displacement – 2 compounds swap atoms • AB + CD AD + CB • Has 2 subtypes: • Acid base – water and a salt are formed • Precipitation – a solid is formed • Combustion – carbon dioxide and water are produced • Oxidation-Reduction – electrons are exchanged (doesn’t fit anywhere else)

  13. Writing Equations from Words • Decide which are your products and which are your reactants • Look for words like creates, yields, forms, etc • Words like and, also, in the presence of usually indicate a plus • Swap and drop your formulas • Check charges from the table, swap if not equal • Remember your diatomic elements (group 17 and all non-noble gases) • Balance your equation • Remember, you can’t have an element on one side, but not on the other

  14. An example • Calcium metal reacts in the presence of oxygen gas to form solid calcium oxide. Write a balanced equation for this. • Start with Ca and O, end with CaO • Ca + O -> CaO • Remember that O is diatomic though so its O2 • Check charges: Ca+2, O-2 so we don’t need subscripts • Then balance so: • 2 Ca + O2 -> 2 CaO

  15. Aqueous Solution • When most ionic compounds are placed in water, the two ions separate from each other and act independently • Ex. NaCl(aq) Na+ + Cl- • Those who don’t can be determined by viewing the rules of solubility on your purple sheet • Ex. Ba(CrO4)(s) Ba(CrO4)(s)

  16. Writing a Complete Molecular Equation • All reactants are soluble (ie they dissolve in water) • Reactants swap partners (need to remember to swap and drop on new compounds) • One of the products will be insoluble (check your purple sheet) and one will be soluble • Soluble compounds have (aq) after them because they separate into separate ions in water • Insoluble compounds have (s) after them because they remain together as a solid (or precipitate) • Need to balance the equation if its not balanced

  17. An Example ? • K2(CrO4)(aq) + Ba(NO3)2 (aq) • K2(CrO4) (aq) + Ba(NO3)2 (aq) K(NO3) + Ba(CrO4) • K2(CrO4) (aq) + Ba(NO3)2 (aq) 2 K(NO3) (aq) + Ba(CrO4) (s) +1 - 1 +2 - 2

  18. Complete Ionic Equation • Split all (aq) compounds up into their separate ions since they don’t actually stay together in solution. • Ex. K2(CrO4) (aq) + Ba(NO3)2 (aq) 2 K(NO3) (aq) + Ba(CrO4) (s) becomes 2K+ + CrO4-2 + Ba+2 + 2NO3- 2K+ + 2NO3- + Ba(CrO4) (s)

  19. Net Ionic Equation • Eliminate all spectator ions (the ones that always stay apart and are never part of a compound) • Ex. 2K+ + CrO4-2 + Ba+2 + 2NO3- 2K+ + 2NO3- + Ba(CrO4) (s) Becomes CrO4-2 + Ba+2 Ba(CrO4) (s)

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