1. Chemical Equations and Reactions

2. Describing Chemical Reactions  A process by which one or more substances are changed into one or more different substances  The original substances are reactants, the resulting substances are the products The mass of reactants and products must be equal

3. Indications of a Chemical Reaction  Certain easily observed changes usually indicate that a chemical change has occurred Evolution of heat and light Production of a gas Formation of a precipitate – a solid that separates from solution Color change

5. Characteristics of Chemical Equations  The equation must represent known facts  The equation must contain the correct formulas for the reactants and products  The law of conservation of mass must be satisfied  Can be either a word equation or a formula equation

7.  There are several symbols used in chemical equations

8. Significance of a Chemical Equation  The coefficients of a chemical reaction indicate relative, not absolute, amounts of reactants and products  The relative masses of the reactants and products of a chemical reaction can be determined from the reaction’s coefficients

9.  The reverse reaction for a chemical equation has the same relative amounts of substances as the forward reaction

10. Balancing Chemical Equations  1. Identify the names of the reactants and the products, and write a word equation  2. Write a formula equation by substituting correct formulas for the names of the reactants and the products

11.  3. Balance the formula equation according to the law of conservation of mass Balance atoms one at a time Balance the atoms that are combined and appear only once on each side Balance polyatomics that appear on both sides as single units Balance H and O atoms last  4. Count atoms to be sure that the equation is balanced

12. Types of Chemical Reactions  Synthesis Reactions A + X AX

13. Reactions of Elements with Oxygen and Sulfur  One simple type of synthesis reaction is the combination of an element with oxygen to produce an oxide of that element  2Mg(s) + O 2 (g)  2MgO(s)  Sulfur behaves in a similar way  8Ba(s) + S 8 (s)  8BaS(s)

14. Reactions of Metals with Halogens  Most metals form ionic or covalent compounds with the Group 17 elements  2Na(s) + Cl 2 (g)  2NaCl(s)

15. Synthesis Reactions with Oxides  Oxides of active metals react with water to produce metal hydroxides  CaO(s) + H 2 O(l)  Ca(OH) 2 (s)  Many oxides of nonmetals in the upper right react with water to produce oxyacids  SO 2 (g) + H 2 O(l)  H 2 SO 3 (aq)

16.  Decomposition Reactions AX A + X A single compound undergoes a reaction that produces two or more simpler substances (the reverse of synthesis reactions)

18.  Single- Replacement Reactions One element replaces a similar element in a compound A + BX AX + B Y + BX BY + X

20. Activity Series  A single replacement reaction will occur if the single element is more active that the one in the compound

21. Double-Replacement Reactions  The ions of two compounds exchange places in an aqueous solution to form two new compounds AX + BY AY + BX One of the products is usually a precipitate, an insoluble gas, or a molecular compound (usually water)

23. Combustion Reactions  A substance combines with oxygen, releasing a large amount of energy in the form of light and heat 2H 2 + O 2 2H 2 O C 3 H 8 + 5O 2 3CO 2 + 4H 2 O

24. Energy & Enthalpy  The heat of reaction is the quantity of energy released or absorbed as heat during a chemical reaction  Thermochemical equations includes the quantity of energy released or absorbed as heat during the chemical reactions 2H 2 (g) + O 2 (g)  2H 2 O(g) + 483.6 kJ

25.  The quantity of energy released corresponds to the molar quantities of the equation  The enthalpy change is the amount of energy absorbed or lost by a system as heat during a process at constant pressure  H = H products – H reactants  A postive H corresponds to an endothermic reaction, A negative H is exothermic

27.  The coefficients refer to the number of moles and can be written as fractions  Physical states must be included in these reactions  If the equation is multiplied by a factor, the energy change is multiplied by that same factor

28.  The heat of formation is the energy released or absorbed as heat when one mole of a compound is formed by combination of its elements Given for the standard states of the elements (state at STP)

29.  Elements in their standard states are defined as have a standard heat of formation of 0 a negative heat of formation means that a compound is more stable than its component elements Postitive or only slightly negative heats mean that it is more stable in its elemental form and will tend to decompose

30.  The heat of combustion is the energy released as one mole of material is combusted  Thermochemical equations can be added together to give changes for other reactions

31.  Hess’s law states that the overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process If a reaction is reversed the sign of is also reversed If the reaction is multiplied by a factor, so is the H

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33. The Reaction Process  Reaction mechanisms are the step-by-step sequence of reactions by which the overall chemical change occurs Species that appear in some steps but not in the net equation are known as intermediates Homogeneous reactions have reactants and products in the same phase

34.  Collision theory describes the conditions required for the effective collision of reactants to produce a product Must have a minimum amount of energy Must have a correct orientation

35.  The energy required for reactants to come together and form the activated complex is know as the activation energy

36.  The activated complex is a transitional structure that results from an effective collision that persists while old bonds are breaking and new bonds are forming

38. Reaction Rate  The change in concentration of reactants per unit time as a reaction proceeds is called the reaction rate  Chemical kinetics is the study of reaction rates and reaction mechanisms  Reaction rate depends on both the frequency of collision and collision efficiency

39.  5 factors that influence reaction rate 1. Nature of reactants 2. Surface area – the greater the surface area, the faster the reaction 3. Temperature – higher T, more and faster collisions, higher rate 4. Concentration – more reactants, more collisions, faster rate

40.  5. Presence of catalysts – substances that change the rate of a chemical reaction without actually reacting Typically provides and alternative pathway for the reaction that has a lower activation energy Homogeneous is in the same phase, heterogeneous is in a different phase

42.  An equation that relates reaction rate and concentrations of reactants is called the rate law for the reaction  For the reaction A + B  2C The forward rate law is R=k[A][B] The reverse rate law is R=k[C] 2 Note the concentration is raised to its molar coefficient

43.  If a reaction occurs in a sequence of steps, the rate law is determined by the slowest step This is the rate determining step

44. Estimated Enthalpy from Bond Energies

45.  Think of bond energy as the energy required to pull a bond apart. Keeping the sign convention in mind, all bond energies are positive!  In order to use bond energies to calculate heat of reaction (enthalpy of reaction), you substitute bond energies into this equation:

46.  You need to be very careful when applying this formula because the order of terms is the opposite of the order used in the heats of summation calculations.heats of summation

48. The molar enthalpy of combustion of ethanol based on bond energies is -1398 kJ/mol. The accepted value is -1368 kJ/mol. There is a difference due to the use of average bond energies.

49. Ex: Combustion of methane