Classification of Chemical Reactions

1. Classification of Chemical Reactions Physical Science Sleevi

2. Chemical Reactions • The process of chemical change • Substances are transformed to different substances • Indicators: • formation of precipitate • unexpected color change • evolution of gas • release or absorption of energy

3. Chemical Reaction A + B  C + D

4. The Language of Reactions • Reactant • starting material of a chemical reaction • appears to the left of the reaction arrow • Product • substance formed in chemical reaction • appears to the right of the reaction arrow • Reaction Conditions • solvent, energy applied, catalysts, etc.

5. Chemical Equations 2Na (s) + 2H2O (l) H2(g) + 2NaOH (aq) s = solid l = liquid g = gas aq = aqueous solution

6. Reactions Conditions Heat = Δ Electricity = e- Light = hv or γ Catalyst = Pt Specific Temperature = 50oC

7. Chemical Reaction • Calcium carbonate is heated strongly and carbon dioxide gas is driven off, leaving a residue of calcium oxide. CaCO3(s) CO2(g) + CaO (s)

8. Descriptions of Chemical Reactions • Identify reactants by language such as: • is heated strongly • decomposes • is combined with • when added to • reacts with • neutralizes • x is converted to…

9. Descriptions of Chemical Reactions • Identify products by language such as: • is formed • produced • precipitates • is given off • is evolved • leaving a residue of • is converted to y

10. Writing Chemical Equations from Descriptions of Reactions • Identify reactants and products • Write correct chemical formulas • Identify and document states of each substance • Record reaction conditions (if given) above/below reaction arrow

11. Elements as Diatomic Molecules • Seven elements occur as diatomic molecules

12. Examples • Bubbling chlorine gas through a solution of potassium iodide gives elemental iodine and a solution of potassium chloride. • Solid silver oxide can be heated to give silver and oxygen gas.

13. Balancing Chemical Equations • Conservation of matter • Conservation of mass • Rearrangement of atoms • Same number and type of atoms on each side of the equation

14. Balancing Chemical Equations • Write the unbalanced equation that describes the reaction (with correct formulas) • Use coefficients to balance the number of atoms of each element on both sides • Start with the elements that appear only once on each side of the equation

15. Balancing Chemical Equations • For elements found in two substances on ONE side of the equation: sum the number of atoms on that side of the equation • Reduce coefficients to lowest whole number ration • Double check!

16. Driving Forces for Chemical Reactions • Formation of a solid • Formation of water • Formation of a gas • Transfer of electrons (Reactions are spontaneous if the products are favored)

17. Types of Chemical Reactions • Decomposition • Combination • Single Replacement • Double Replacement • Complete Combustion of Hydrocarbons

18. Decomposition Reactions • Single substance broken down into two or more simpler substances • Reactant must be compound • Products can be elements or compounds • Require input of energy to occur (light, heat, electricity) A  B + C

19. Decomposition Reactions 2H2O  2H2 + O2 2H2O2  2H2O + O2 K2CO3  K2O + CO2 2KOH  K2O + H2O

20. Decomposition Reactions • Binary compound decomposes to elements • Metal carbonate decomposes to metal oxide and carbon dioxide • Base decomposes to metal oxide and water

21. Combination Reactions • Two or more substances form one product • Reactants can be elements or compounds • Product is always a compound A + B  C

22. Combination Reactions Na (s) + Cl2 (g) NaCl (s) SO3(g) + H2O (l)  H2SO4(aq) K2O (s) + H2O (l)  KOH (aq)

23. Combination Reactions • Two elements combine to form a binary compound • metal + nonmetal  ionic compound • nonmetal + nonmetal  molecular compound • Nonmetal oxide + water  acid • Metal oxide + water  base

24. Single Replacement Reactions • Substitution reactions in which an element replaces the element in an ionic compound • metal replaces metal • halogen replaces halogen • metal replaces hydrogen (in an acid) • Not all reactions occur A + BX  AX + B

25. Single Replacement Reactions Mg + ZnCl2  MgCl2 + Zn Fe + CuSO4  FeSO4 + Cu Na + H2O  NaOH + H2 F2 + KCl  KF + Cl2

26. Single Replacement Reactions • Activity Series provides reference for which single replacement reactions occur Note: All examples we will use are reactions that occur. You will not need to use the activity series to determine whether or not a reaction occurs

27. Double Replacement Reactions • Exchange of positive ions between two compounds • Usually occur between ionic compounds in aqueous solutions • When reaction occurs • a precipitate forms • water or other molecular compound is formed • evolution of a gas

28. Double Replacement Reactions Na2S (aq) + Cd(NO3)2(aq) NaNO3 (aq) + CdS (s) NaCN (aq) + H2SO4(aq)  HCN (g) + Na2SO4(aq) AgNO3(aq) + KCl (aq) AgCl (s) + KNO3(aq) NaOH (aq) + H2SO4(aq) Na2SO4(aq) + H2O (l)

29. Combustion Reactions • Hydrocarbon burns in the presence of oxygen • For complete combustion the products are ALWAYS carbon dioxide and water

30. Combustion Reactions CH4 + O2  CO2 +H2O C6H12O6+ O2  CO2 +H2O C6H6 + O2  CO2 +H2O