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Understanding Chemical Changes and Reactions: Key Concepts and Examples

This chapter delves into chemical changes, defining the processes that transform reactants into products through observable signs. It explains the difference between physical changes and chemical reactions, focusing on energy transfer during exothermic and endothermic processes. The chapter also emphasizes the Law of Conservation of Mass and provides guidance on balancing chemical equations. Various reaction types—combustion, synthesis, decomposition, and displacement—are examined, alongside the activity series for predicting reactions.

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Understanding Chemical Changes and Reactions: Key Concepts and Examples

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  1. Chemical Equations Chemistry Chapter 9

  2. Chemical Change • How do you know a chemical change has taken place? • What are some common examples of chemical changes?

  3. Chemical Reactions • The process by which one or more substances (reactants) are changed into one or more different substances (products) • Observations that a chemical reaction has taken place: • The evolution of energy as heat, light or sound • The production of gas • The formation of a precipitate • Change in color

  4. Physical Changes vs. Chemical Reactions • A physical change does NOT change chemical composition or molecular structure of the reactant • Condensation, melting, and crystallization are physical changes

  5. Law of Conservation of Mass • Mass cannot be created or destroyed, so: • The products of a reaction are made up of the same number and kinds of atoms as were present in the reactants • The bonding patterns are rearranged • HCl + NaOH  NaCl + H2O

  6. Energy in a Chemical Reaction • Exothermic reaction: release energy (normally noticeable by heat), normally spontaneous at room temperature • Endothermic reaction: require energy so do not normally occur at room temp • Reactions are spontaneous if products are continually generated as long as reactants are supplied

  7. Spontaneous or Not?

  8. Nonspontaneous Reactions • Can occur spontaneously when linked to an energy source • The electrolysis of water: • Does water decompose into its component parts at room temp? • What energy source must be applied?

  9. Chemical Equations • Describe the type and number of atoms that are rearranged during a reaction • Word equations • Formula equations • Correctly written chemical equations must be balanced to satisfy the law of conservation of mass

  10. Chemical Equations • Conditions under which a reaction occurs found above or below the arrow • Physical state of the reactants and products abbreviated and put in parenthesis after the compound

  11. Balancing Chemical Equations • Inserting coefficients so that there are equal numbers of atoms for each element on each side of the equation

  12. Tips for Balancing Equations • Delay the balancing of elements (often hydrogen and oxygen) that occur in several reactants or products. • If the same polyatomic ions appear on both sides of the equation treat them as single units, like monatomic ions. • Balance the elements left to right. • Remember, balancing one element may unbalance others. • For ionic equations, be sure charges are balanced.

  13. Practice • Page 316 a-d

  14. Quantitative Relationships • A balanced chemical equation can tell you • Moles of reactants and products • Molecules of reactants and products • Molar ratios can be determined

  15. Energy Changes in Equations • Endothermic Reactions: require energy and energy needed (in kJ) is written on the reactant side of the equation • Exothermic Reactions: release energy and is written on the product side of the equation

  16. Enthalpy • The total energy content in a system • Endothermic reactions: delta H is positive because the energy needed to break the bonds increases the total energy of the system • Exothermic: delta H is negative because the energy is released when the stronger bonds of the products are created

  17. Energy expressed as a Mole Ratio • Find the amount of energy released when 100 g of CaCl2 is formed from the free elements that compose it.

  18. Types of Reactions • Combustion Reactions • Oxidation Reactions • Synthesis Reactions • Polymerization Reactions • Decomposition Reactions • Displacement Reactions • Double-displacement Reactions

  19. Combustion Reactions • Normally exothermic and require a “push” to get started • Ex: the reaction between organic compounds and oxygen • Bunsen Burner CH4(g) + 2O2(g)CO2(g)+2H2O(g) + 803kJ • Oxidation Reactions also include oxygen, but are not as dramatic • Rusting of Iron

  20. Synthesis Reactions • Complex substances are made from simpler substances • Ex: synthesis of glucose 6CO2(g) + 6H2O(l)  C6H12O6(aq) + 6O2(g) • Polymerization Reactions: a series of synthesis reactions that take place to produce a very large molecule

  21. Decomposition Reactions • A compound is broken down into smaller substances • Ex: CH3OH(g) CO(g) + 2H2(g)

  22. Displacement Reactions • A chemical reaction in which one element replaces another element in a compounds that is in solution • Ex: 2Al(s) + 3CuCl2(aq) 2AlCl3(aq)+3Cu(s)

  23. Activity Series • Used to predict how elements will react in displacement reactions (also sometimes in synthesis and decomposition reactions) • Listed in a table with the most active element at the top • In a reaction elements will replace less active elements in a compound (those below it on the table) • The farther apart two elements are on the activity series, the more likely it is that the higher one will quickly displace the lower one in compounds

  24. 2K + 2HOH  2KOH + H2

  25. Double Displacement Reactions • A chemical reaction in which ions from two compounds interact in solution to form a product (two cations displace each other) • Ex: 2KI(aq)+Pb(NO3)2(aq)PbI2(s)+2KNO3(aq)

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