2. Indicators of chemical reactions Formation of a gas Emission of light or heat Formation of a precipitate Color change Emission of odor

3. All chemical reactions: Two parts: Reactants – starting substances Products- ending substances The reactants turn into the products

4. Describing chemical reaction Combination of atoms changes Atoms are neither created nor destroyed. Can be described several ways Copper reacts with chlorine to form copper (II) chloride. Copper + chlorine  copper (II) chloride Cu(s) + Cl 2 (g)  CuCl 2 (aq)

5. Symbols used in equations (s) after the formula – solid Cu (s) (g) after the formula – gas H 2 (g) (l) after the formula – liquid H 2 O (l) (aq) after the formula – dissolved in water, an aqueous solution. CaCl 2 (aq)

6. Summary of Symbols

7. What is a catalyst? A substance that speeds up a reaction without being changed by the reaction. Enzymes are biological or protein catalysts.

8. All chemical reactions are accompanied by a change in energy. Exothermic - reactions that release energy to their surroundings (usually in the form of heat) ΔH (enthalpy) is negative – energy leaving system Endothermic - reactions that need to absorb heat from their surroundings to proceed. ΔH (enthalpy) is positive – energy coming into the system Reaction Energy

9. Spontaneous Reactions - Reactions that proceed immediately when two substances are mixed together. Not all reactions proceed spontaneously. Activation Energy – the amount of energy that is required to start a chemical reaction. Once activation energy is reached the reaction continues until you run out of material to react. Reaction Energy

11. Formula Equation Use formulas and symbols to describe a reaction. Doesn’t indicate how many All chemical equations are sentences that describe reaction.

12. Diatomic elements 8 elements that never want to be alone, form diatomic molecules. H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, and At 2 The –ogens and the –ines 1 + 7 pattern on the periodic table

14. Balancing Equations ___ H 2 (g) + ___ O 2 (g) ---> ___ H 2 O(l) 22

15. What Happened to the Other Oxygen Atom? This equation is not balanced! Two hydrogen atoms from a hydrogen molecule (H 2 ) combines with one of the oxygen atoms from an oxygen molecule (O 2 ) to form H 2 O. Then, the remaining oxygen atom combines with two more hydrogen atoms (from another H 2 molecule) to make a second H 2 O molecule. __2_ H 2 (g) + ___ O 2 (g) ---> __2_ H 2 O(l)

16. Aluminum metal reacts with liquid bromine to form solid aluminum bromide Translate Equation ___ Al(s) + ___ Br 2 (l) →___ AlBr 3 (s) 2 3 2

17. Types of Reactions There are millions of reactions. Fall into several categories. Focus on 6 types. Be able to predict the products. Be able to predict if they will happen at all.

18. 1. Synthesis (combination) Reactions 2 elements, or compounds, combine to make one compound. A + B  AB Na (s) + Cl 2 (g)  NaCl (s) Ca (s) +O 2 (g)  CaO (s) SO 3 (s) + H 2 O (l)  H 2 SO 4 (s) Can predict the products if they are two elements. Mg (s) + N 2 (g)  Mg 3 N 2 (s)

19. A simulation of the reaction: 2H 2 + O 2  2H 2 O

20. 2. Decomposition Reactions decompose = fall apart one compound (reactant) falls apart into two or more elements or compounds. Usually requires energy AB  A + B 2NaCl  2Na + Cl 2 CaCO 3  CaO + CO 2

21. Can predict the products if it is a binary compound Made up of only two elements Falls apart into its elements H 2 O  H 2 (g) + O 2 (g) 2HgO  2Hg (s) + O 2 (g) Decomposition Reactions

22. If the compound has more than two elements you must be given one of the products The other product will be from the missing pieces NiCO 3 (aq)  CO 2 (g) + 2Ni (s) H 2 CO 3(aq)  H 2 (g) + CO 2 (g) Decomposition Reactions

23. 3. Single Replacement One element replaces another Reactants must be an element and a compound. Products will be a different element and a different compound. A + BC  AC + B 2Na + SrCl 2  Sr + 2NaCl F 2 + LiCl  LiF + Cl 2

24. Single Replacement We can tell whether a reaction will happen Some are more active than others More active replaces less active

25. Reactivity Series

26. 4. Double Replacement Two things replace each other. Reactants must be two ionic compounds or acids. Usually in aqueous solution AB + CD  AD + CB AgNO 3 + NaCl  AgCl + NaNO 3 ZnS + 2HCl ® ZnCl + H 2 S

27. 5. Combustion A reaction in which a compound (often carbon) reacts with oxygen CH 4 + O 2 ® CO 2 + H 2 O C 3 H 8 + O 2 ® CO 2 + H 2 O C 6 H 12 O 6 + O 2 ® CO 2 + H 2 O

28. 6. Acid/Base Reaction An acid and a base react to form a salt and water. Always in aqueous solution Acid (H + ) + Base (OH - ) → Salt + H 2 O NaOH + HCl → NaCl + H 2 O NH 4 OH + H 2 SO 4 → (NH 4 ) 2 SO 4 + H 2 O

29. How to recognize which type Look at the reactants Element(E), Compound(C) E + E C E + C C + C Acid + Base Look at the Products CO 2 + H 2 O Redox Synthesis Decomposition Single replacement Double replacement Acid/Base reaction Combustion

30. Examples Synthesis Decomposition Single replacement Double replacement H 2 + O 2 ® H 2 O ® AgNO 3 + NaCl ® Zn + H 2 SO 4 ® HgO ® KBr +Cl 2 ® Mg(OH) 2 + H 2 SO 3 ®

31. Examples Acid/Base Decomposition Single replacement Synthesis Acid/Base Single replacement Double replacement HNO 3 + KOH ® CaPO 4 ® AgBr + Cl 2 ® Zn + O 2 ® HgO + Pb ® HBr + NH 4 OH ® Cu(OH) 2 + KClO 3 ®