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Chemical Equations

Chemical Equations. Chemistry Chapter 9. Chemical Change. How do you know a chemical change has taken place? What are some common examples of chemical changes?. Chemical Reactions.

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Chemical Equations

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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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