1. Chemical Reactions & Equations Chapter 8 – Modern Chemistry Text

2. Chemical Reaction Defined  A process in which substances are changed into different substances with different chemical and physical properties.  Reactants: substances that react, left side of equation  Products: substances that form, right side of equation

3. Why Chemical Reactions Occur  Only valence electrons are involved in chemical bonding.  Atoms with full outer E levels will not bond.  Atoms with partially filled outer E levels will gain, lose or share electrons in order to complete the outer E level (= The Octet Rule.)  Elements that lose (or gain) valence electrons easily are more reactive (higher activity).

4. Energy in Chemical Reactions  Energy is either absorbed or released in a chemical change.  Endothermic: energy absorbed so heat is a reactant.  Exothermic: energy released so heat is a product.

5. What is a Chemical Equation?  An expression in which symbols and formulas are used to represent a chemical reaction.  It describes the chemical change.

6. Steps for writing a balanced equation 1.Read the reaction/problem & determine reactants and products. 2.Write formulas for the reactants on the left of the arrow and products on the right. 3.Cover the reactant formulas when writing the product formulas – they no longer exist – all the atoms become rearranged. 4.Use the subscript 2 when writing the formula for any diatomic gas. 5.Check formulas.  DO NOT CHANGE THEM!  6.Use coefficients ONLY to balance the equation.

7. Symbols used in chemical equations +reacts with +reacts with ↔reaction is reversible → “yields”, forms, produces ↓ precipitate forms as a product ↑ gas forms as a product (s) solid (l) liquid (g) gas (aq) aqueous (in solution with water) ∆substances are heated

8. COEFFICIENTS  Whole numbers written before the formulas for reactants and products.  Use only coefficients to balance an equation.  NEVER change a correctly written chemical formula (or add substances) to balance an equation.  NEVER change subscripts!

9. why chemical equations must be balanced…  Law of Conservation of atoms/mass: –Atoms are the units of chemical change –Atoms are never created or destroyed in a chemical Rxn, they are simply rearranged  When the number of atoms of each element is the same on both sides of the equation, the equation is BALANCED.

10. Balancing Hints & Tips 1.Start with metals or any elements that appear only once on both sides of the equation. 2.Balance polyatomic ions as single units if they appear on both sides of the equation. 3.Write water as H-OH in single and double replacement reactions. 4.Do H and O last; they often appear in more than one compound. 5.Look for 3/2 combos.

11. Evidence of chemical change 1)The evolution of a gas: Observed by bubbling or emission of an odor. 2)The formation of a precipitate: A precipitate is an insoluble solid formed from a solution. 3)The evolution or absorption of heat: If the reaction system increases in temp., exothermic. If the reaction system decreases in temp., endothermic. 4)The emission of light. 5)A color change.

12. Reference Materials Needed 1)pencils 2)Periodic Table 3)Ion Tables – see Table 7-1, p.221 & Table 7-2, p.226. 4)The seven diatomic molecules – see Table 8-1, p.263. 5)The Hydrocarbons 6)Acids (see table 7-5, p. 230) 7)Activity Series - for single replacement reactions– see table 8-3, p.286) 8)Solubility Table - for double replacement reactions – see Appendix A, p.860)

13.  There are 5 basic types of chemical reactions: 1.Synthesis 2.Decomposition 3.Single-Displacement 4.Double-Displacement 5.Combustion Types of Chemical Reactions Chapter 8 – Section 2: Types of Chemical Reactions

14.  In a synthesis reaction 2 or more substances combine to form a new compound.  This type of reaction is represented by the following general equation: A + X AX Synthesis Reactions Chapter 8 – Section 2: Types of Chemical Reactions

15. 2Mg(s) + O 2 (g) 2MgO(s) 2Mg(s) + O 2 (g) 2MgO(s) S 8 (s) + 8O 2 (g) 8SO 2 (g) S 8 (s) + 8O 2 (g) 8SO 2 (g) 2H 2 (g) + O 2 (g) 2H 2 O(g) 2H 2 (g) + O 2 (g) 2H 2 O(g) 2Na(s) + Cl 2 (g) 2NaCl(s) 2Na(s) + Cl 2 (g) 2NaCl(s) Mg(s) + F 2 (g) MgF 2 (s) Mg(s) + F 2 (g) MgF 2 (s) Synthesis – Examples NOTICE ONLY ONE PRODUCT Chapter 8 – Section 2: Types of Chemical Reactions

16.  In a decomposition reaction, a single compound breaks apart to form 2 or more simpler substances.  Decomposition is the opposite of synthesis.  This type of reaction is represented by the following general equation: AX A + X Decomposition Reactions Chapter 8 – Section 2: Types of Chemical Reactions

17. 2H 2 O(l) 2H 2 (g) + O 2 (g) CaCO 3 (s) CaO(s) + CO 2 (g) H 2 CO 3 (aq) CO 2 (g) + H 2 O(l) 2HgO(s) 2Hg(l) + O 2 (g) Ca(OH) 2 (s) CaO(s) + H 2 O(g) Decomposition – Examples NOTICE ONLY ONE REACTANT Chapter 8 – Section 2: Types of Chemical Reactions electricity ∆ ∆ ∆

18.  In a single-displacement reaction (also called single- replacement) one element replaces a similar element in a compound.  They often take place in aqueous solution.  This type of reaction is represented by the following general equation: A + BX AX + B Single-Displacement Reactions Chapter 8 – Section 2: Types of Chemical Reactions

19. 1.Metal + Water: hydrogen is displaced 2Na(s) + 2H-OH (l) 2NaOH(aq) + H 2 (g) 2.Metal + Acid: hydrogen is displaced Mg(s) + 2HCl(aq) H 2 (g) + MgCl 2 (aq) 3. HALOGENS! …are nonmetals, replace halogens! Cl 2 (g) + 2KBr(aq) 2KCl(aq) + Br 2 (l) 4.More active metal replaces less active metal: Zn(s) + CuCl 2 (aq) ZnCl 2 (aq) + Cu(s) Single Displacement - Examples Chapter 8 – Section 2: Types of Chemical Reactions

20.  The ability of an element to react is referred to as the element’s activity.  The more easily an element reacts with other substances, the greater its activity is. Chemical Activity Chapter 8 – Section 3: Activity Series of the Elements Li Au

21.  An activity series is a list of elements organized by their chemical activity. The most-active element is placed at the top in the series.The most-active element is placed at the top in the series.  It can replace each of the elements below it in a single- displacement reaction.  Activity series can also be used to predict whether a chemical reaction will occur with acids, water and oxygen.  Activity series are based on experiment. Activity Series – Use for single replacement reactions Chapter 8 – Section 3: Activity Series of the Elements

22.  In double-displacement reactions, the ions of 2 compounds exchange places in an aqueous solution to form 2 new compounds.  One of the compounds formed is usually either a precipitate, a gas, or a molecular compound like water.  Represented by the following general equation: AX + BY AY + BX Double-Displacement Reactions Chapter 8 – Section 2: Types of Chemical Reactions

23. Formation of a Precipitate 2KI(aq) + Pb(NO 3 ) 2 (aq) PbI 2 (s) + 2KNO 3 (aq) 2KI(aq) + Pb(NO 3 ) 2 (aq) PbI 2 (s) + 2KNO 3 (aq) Formation of a Gas FeS(s) + 2HCl(aq) H 2 S(g) + FeCl 2 (aq) FeS(s) + 2HCl(aq) H 2 S(g) + FeCl 2 (aq) Formation of Water HCl(aq) + NaOH(aq) NaCl(aq) + H-OH(l) HCl(aq) + NaOH(aq) NaCl(aq) + H-OH(l) This reaction is a neutralization reaction! Double Displacement - Examples Chapter 8 – Section 2: Types of Chemical Reactions

24.  In a combustion reaction, a fuel combines with oxygen, releasing a large amount of energy in the form of light and heat.  Products of combustion reactions are always carbon dioxide and water vapor.  Example: Combusion of propane C 3 H 8 (g) + 5O 2 (g) 3CO 2 (g) + 4H 2 O(g) Combustion Reactions Chapter 8 – Section 2: Types of Chemical Reactions

25. Classify as a syn, decomp, SR, DR, or combustion reaction. N 2 (g) + 3H 2 (g) → 2NH 3 (g) 2Li(s) + 2H 2 O(l) → 2LiOH (aq) + H 2 (g) 2NaNO 3 (s) → 2NaNO 2 (s) + O 2 (g) 2C 6 H 14 (l) + 19O 2 (g) → 12CO 2 (g) + 14H 2 O(l) HCl (aq) + NaOH (aq) → NaCl (aq) + H-OH (l) Types of Reactions - Sample Problems Chapter 8 – Section 2: Types of Chemical Reactions synthesis single-displacement decomposition combustion double-displacement

26. Use the activity series to predict whether or not there will be a reaction for the possibilities below. If a reaction will occur, write the products and balance the equation. a.Ni(s) + Pb(NO 3 ) 2 (aq) → b.MgCl 2 (aq) + Zn(s) → c.Br 2 (l) + KI(aq) → d.Cu(s) + HCl(aq) → Activity Series Sample Problem Chapter 8 – Section 3: Activity Series of the Elements Yes No reaction Pb(s) + Ni(NO 3 ) 2 (aq) I 2 (l) + KBr(aq) 2 2

27. Chemical Kinetics  Area of chemistry concerned with speed of reactions.  The collision theory relates particle collisions to reaction rate.

28. Reaction mechanisms  Molecules, atoms or ions of reactants must collide with each other with sufficient KE to form products.  The molecules must have sufficient energy to initiate the reaction.  In some cases, the orientation of the molecules during the collision must also be considered.

29. Rates of Chemical Reactions  In all chemical reactions, there is a change in the energy within the system.  Activation energy - the energy needed to start a chemical reaction.  Reaction rate - the rate at which products form or reactants are used up. –Δconcentration/Δtime = Molarity/second  Dynamic Equilibrium- when no net change occurs in the amount of reactants and products.

30. LeChatelier’s Principle  Disturbing an equilibrium will make a system readjust to reduce disturbance and regain equilibrium.

31. Examples that favor the forward reaction: 1.Remove product as it forms. 2.Form a precipitate. 3.Add more reactant. 4.Keep an exothermic rxn cool. 5.Add heat to an endothermic rxn.

32. Factors affecting reaction rates:  CONCENTRATION: the greater the concentration of a reactant, the greater the rate of the reaction.  PRESSURE: increases the rates of reactions which take place with the reactants in the gaseous phase.  TEMPERATURE: The higher the temperature, the higher the reaction rate.

33. Factors affecting reaction rates:  CATALYSTS are substances which speed up reaction rates without undergoing a permanent change in the process. The process is called catalysis, and it may be homogeneous, when both the reactants and the catalyst are in the same phase (e.g., both are in solution), or heterogeneous, when the reactants and catalyst are in different phases (e.g., gaseous reactants and a solid catalyst.)

34. Factors affecting reaction rates:  BOND TYPES: Reactions b/w ionic cmpds are usually much faster than those involving covalent bonds.  SURFACE AREA: This is very important when solids are involved. The more finely divided the solid is, the faster the reaction will take place.