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Chemical Reactions

Chemical Reactions. When one or more substances are changed into different substances a chemical reaction has occurred. Reactants  Products Word equations use names of chemicals Chemical equations use the chemical formulas. Example. Example.

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Chemical Reactions

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  1. Chemical Reactions

  2. When one or more substances are changed into different substances a chemical reaction has occurred. • Reactants  Products • Word equations use names of chemicals • Chemical equations use the chemical formulas

  3. Example

  4. Example • Iron + sulphur  iron (II) sulphide + energy • Fe (s) + S (s)  FeS (s) + energy • NOTE: energy is released in this reaction and is therefore a product...sometimes more energy is absorbed and it is a reactant • States of matter are identified in brackets beside each chemical.

  5. (s) = solid • (l) = liquid • (g) = gas • (aq) = aqueous • Another example: • Fe (s) + CuSO4 (aq)  FeSO4 (aq) + Cu (s) + energy

  6. Practice • Page 227 #2

  7. Homework • Page 227 #3-6

  8. Conservation of Mass

  9. Matter and mass can not be created nor can it be destroyed • The mass of the reactants is equal to mass of the products

  10. Balancing Equations

  11. Counting Atoms • subscripts indicate the number of atoms for the element preceding it • Ex. CaF2 has 1 calcium atom and 2 Fluorine atoms • Coefficients in front of a molecule indicate that number of molecules • Ex. 4H2O is 4 molecules of water: therefore there are 8 hydrogen and 4 oxygen atoms • Polyatomic ions can be counted together (do not separate the elements • Ex. Fe(NO3)3 there are 3 nitrates with iron: therefore there are 1 iron atom, and 3 nitrate ions

  12. Balancing • Count atoms on the left and then the right • For elements which are not balanced you may use COEFFICIENTS ONLY and not change the subscripts on any of the molecules

  13. Example: Ca + O2 CaO • Ca = 1 Ca = 1 • O = 2 O = 1 • Calcium is balanced however the Oxygen is not...place a 2 in front of CaO. • Ca + O2 2CaO • Oxygen is now balanced with 2 on each side however now there are 2 Calcium atoms on the right...so that means there are 2 on the left. • Balanced equation is • 2Ca + O2 2CaO

  14. Types of Reactions

  15. Synthesis • Elements react and combine to form a new compound: • A + B  AB • Example: Iron + Oxygen • Fe + O2 Fe2O3 • Balanced??? 4Fe + 3O2 2Fe2O3 • Example: H2 + O2 H2O • Balanced?? 2H2 + O2  2H2O

  16. Decomposition • Molecules are broken down into smaller molecules or individual elements • AB  A + B • Example: hydrogen peroxide • H2O2  H2O + O2 • Balanced???? 2H2O2  2H2O + O2

  17. Classwork • Page 239 #1,2 • Homework #5

  18. Single Displacement • Exchange of an ion with an ion from another compound • A + BX  B + AX element A has displaced element B from the compound BX • Example: Fe + CuCl2  Cu + FeCl2 • In the reaction Iron displaces the copper from the copper chloride in solution...recall in this reaction the iron turns copper in colour and the blue solution turns clear (due to the iron (II) chloride now in solution).

  19. Double Displacement • The exchange of ions between two different compounds...often produces a solid precipitate • AX + BY  BX + AY in this equation cations A and B have displaced each other • Example: • Fe(NO3)3 + NaSCN  Fe(SCN)3 + NaNO3 • Recall in the above reaction two clear solutions resulted in a dark red precipitate being produced (the sodium thiocyanate)

  20. Classwork • Page 243#2,3,4,7

  21. Combustion • reactions with oxygen to produce oxides • Fuel + oxygen  oxide + energy • Energy produced is mainly heat and light • Most fuels are hydrocarbons which also produce water vapour and as a byproduct • Complete combustion • Propane + oxygen  carbon dioxide + water • 2C3H8 + 7O2  4CO2 + 6H2O • NOTE: every carbon from the hydrocarbon is converted into carbon dioxide

  22. Incomplete Combustion: • When there is not enough oxygen • Carbon monoxide (poisonous gas) and carbon (soot) are also produced • Example: • 2C2H6 + O2 CO2 + 2CO + C + 6H2O • Other Combustion reactions: elements will also react to produce oxides. • Example: magnesium burns to produce magnesium oxide similar to carbon producing carbon dioxide • 2Mg + O2  2MgO • ** combustion of elements are synthesis reactions

  23. Combustion • Page 251#

  24. Corrosion • Page 254

  25. Chapter Review • Page 258

  26. Acids and Bases

  27. Acids • Are sour tasting, water soluble and highly reactive • conduct electricity due to ionic properties • HCl - H+ and Cl- ions • Common acids are recognized from the H at the beginning of chemical formulas • HCl (hydrochloric acid), H2CO3 (carbonic acid), H2SO4 (sulfuric acid)

  28. Bases • Are bitter tasting, water soluble and feel slippery to the touch in solution • Also good conductors of electricity due t ionic nature • Most bases are recognized from the hydroxide in the formula • NaOH (sodium hydroxide), KOH (potassium hydroxide)

  29. pH • 0 -----------------------7---------------------------14 • High H+ Equal OH- and H+ High OH- neutral weak acid Weak Base Very strong base Very strong acid

  30. pH scale is logarithmic • Each change in 1 pH is equivalent to 10X difference in strength • Example a pH of 3 is 10x more acidic than an acid with pH of 4...and an acid with pH 3 would be 1000X more acidic than an acid with pH of 6 • Similarly a base with pH of 9 is 10x stronger than a base with pH of 8 • note that for bases the strength increases with higher pH...whereas acid strength increases with lower pH

  31. Neutralization • HCl + NaCl NaCl + HOH • H2CO3 + KOH  K2CO3 + HOH • Acid + Base  Salt + Water

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