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Chapter 6

Chapter 6. Chemical Reactions and Equations. Warm-up 2/4. Who can recall what the 4 pictures were on the intro slide for this chapter?. Fire. Two flasks (one with a blue liquid and one with a yellow liquid). Test tube with a blue solid substance. Smoke trail.

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Chapter 6

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  1. Chapter 6 Chemical Reactions and Equations

  2. Warm-up 2/4 Who can recall what the 4 pictures were on the intro slide for this chapter? Fire Two flasks (one with a blue liquid and one with a yellow liquid) Test tube with a blue solid substance Smoke trail What do all of these things have to do with chapter 6? They all show signs of a chemical change

  3. Indications of a Chemical Change Chemical Change The change of one or more substances into new substances by the rearrangement of atoms. Change in energy Change in color Formation of a Precipitate (two liquids combine and a solid is produced) Formation of a gas

  4. Parts of a Chemical Equation Reactants Products Coefficients Physical States (solid, liquid, gas or aqueous)

  5. Writing Equations Word Equations Using the names of chemicals to describe chemical reactions Chemical Equations Using chemical formulas and numbers to describe chemical reactions

  6. Practice ProblemsUse your ion tables Re-write the following chemical equation in words Re-write the following word equation as a chemical equation Solid copper + fluorine gas = solid copper (II) fluoride

  7. Energy When chemical reactions take place, energy is either absorbed or released in the process An endothermic reaction + energy An exothermic reaction

  8. Warm up 2/5 Watch the following video to complete the warm up http://www.youtube.com/watch?v=R6bBs2D0cpA Write the following chemical equation for the reaction that occurred in the video: Aluminum + iodine = aluminum iodide You should be able to answer the following questions: What are the reactants? What are the products? What was the indication that a chemical change took place? Was this an endotherm or exothermic reaction?

  9. Physical States of Elements • Metals • Metals are always represented by their symbols from the periodic table, with no subscripts • All metals, with the exception of mercury are solid and will have an (s) following their symbols in an equation Cu(s) • Non-metals • Are often diatomic (Br,I,N,Cl,H,O,F) so they will have a subscript of 2. • Most of these diatomic elements will be a gas except iodine (s) and bromine (l) Cl2(g)

  10. Physical States of Ionic Compounds • Ionic Compounds • Their physical state will be • (aq) = aqueous if they can dissolve in water • (s) = solid if they can’t Use this table to determine the solubility of ionic compounds:

  11. More Practice Problems What would be the physical states of the following substances if written in a chemical equation? PbI2 __________ KNO3 ____________ KI __________ Pb(NO3)2 ____________ Write the following equation with physical states: Lead (II) nitrate combines with potassium iodide to form potassium nitrate and lead (II) iodide.

  12. Warm-up 2/8 Write the chemical equation for the reaction: (Use a solubility table to predict the physical states of each compound) potassium iodide + lead (II) nitrate = potassium nitrate and lead (II) iodide Watch the following video to see what happens http://www.youtube.com/watch?v=BhO8kGrc4dg Did you write the correct physical state for each compound in the reaction?

  13. Balancing Chemical Equations Why? To demonstrate the Law of Conservation of Mass which states: In a chemical reaction mass is neither created or destroyed Antoine Lavoisier What must be conserved? The number and type of atoms 2 sodium atom + 2 chlorine atoms = 2 sodium atoms + 2 chlorine atoms

  14. How to Balance a Chemical Equation • Make sure that both sides of an equation have an equal number of each type of atom represented. • Balance equations by adding coefficients. • Never mess with subscripts Na Cl Mg Br

  15. Strategies for Balancing Chemical Equations 1. The even/odd combination Make the subscript of one, the coefficient of the other and vice versa. Iron (III) oxide Iron + Oxygen Practice:

  16. Strategies Continued 2. Balance groups as a whole if the elements of the group are only represented in the group on each side of the equation. Mg OH Al Practice:

  17. Strategies Continued 3. Balancing Hydrocarbon Combustion Reactions The first step is to add a coefficient to the hydrocarbon so that the total number of hydrogens in the molecule is divisible by four. Then, balance the product side of the equation before finishing up with oxygen. C H O Practice:

  18. Strategies Continued • Splitting Water When hydroxide shows up on one side of an equation and water is on the other side, split water into H(OH) to balance out hydroxide. Practice:

  19. Things to Keep in Mind • Again, do not mess with subscripts • In the final solution, the coefficients have to be in the lowest whole number ratio. • If at first you don’t succeed, try, try again! Most of the solutions to these problems come from trial and error.

  20. Comprehensive Practice Problem • Write the equation • Note the physical state of each substance • Balance the equation Potassium chromate combines with lead (II) nitrate to form potassium nitrate and lead (II) chromate

  21. Types of Chemical Reactions • 5 General Types • Synthesis • Decomposition • Single displacement • Double displacement • Combustion You will need to know how to identify these reactions

  22. Types of Chemical Reactions – 5 General Formulas Combustion CxHy + O2 CO2 + H20 Synthesis A + B AB (Two elements combine to form a compound) Decomposition AB A + B (One compound splits to form two compounds) Always the products of these combustion reactions

  23. Reactions Cont. Single Replacement (Single Displacement) A + BC AC + B Metals replace metals Non-metals replace non-metals Double Replacement (Double Displacement) AB + XY AY + XB Compounds switch partners

  24. Practice Problems Combustion Decomposition Single replacememt Synthesis Double Replacement Identify the type of reaction

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