CHEMICAL REACTIONS I. A CHEMICAL REACTION is the process by which a chemical change is brought about. A. The Law of Conservation of Matter states that matter cannot be created nor destroyed during a chemical change. Matter is only rearranged. The same atoms that were present before the reaction are present after the reaction. B. Word equations tell the chemicals used but does not tell how much. 1. Sodium and chlorine produce sodium chloride. 2. Hydrogen and oxygen yield water. C. Formula equations tell the chemicals used and also tell how much: qualitative –chemicals that are present, quantitative how much (mass) of a chemical is present. 1. EXAMPLES: a. 2 Na + Cl2 2 NaCl b. 2 H2 + O2 2 H2O c. 2Ag + S Ag2S 2. An arrow stands for yield or forms or produces. 3. Reactants are chemicals on the left side of the arrow. 4. Products are chemicals on the right side of the arrow. 5. Plus signs ( + ) means and. Understanding Chemical Equations Let’s use the following reaction as an example… HCl + NaOH → H2O + NaCl The reactants are your starting chemicals… HCl & NaOH The products are what you end up with after the reactants interact. Here the H2O & NaCl If you were going to read the above reaction out loud you would say… HCl plus NaOH yields H2O plus NaCl products are… D. Rules for writing and balancing an equation. 1. The total number of atoms must be the same on both sides of the equation. a. H2 + O2 H2O this equation is not balanced 2H 2O 2H and 1O the number of atoms on each side are not the same therefore, use coefficients to balance the equation. balanced 2 H2 + O2 2 H2O b. Zn + HCl ZnCl2 + H2 not balanced c. Zn + 2HCl ZnCl2 + H2 coefficients are added to balance the equation AgNO3 + MgCl2 AgCl + Mg(NO3)2 not balanced 2AgNO3 + MgCl2 2AgCl + Mg(NO3)2 balanced 2. Read the statement. Write out symbols of elements with correct formulas. a. 7 diatomic elements have to be written with the subscript 2. H2 N2 O2 F2 Cl2 Br2 I2 BrINClHOF b. Formulas are written and balanced before being put in the equation. c. Once the formula is balanced with subscripts, DO NOT CHANGE IT. 3. Make the number of atoms equal on each side of the arrow, by adding a coefficient in front of each formula that needs one. DO NOT EVER CHANGE THE FORMULA! 4. Count the number of atoms on each side of the equation. If they are equal the equation is balanced. If they are not equal, recheck the charges of the elements, the subscripts of the formulas and the coefficients of the compounds. Understanding Balancing Chemical Equations Definition – An equation is balanced when there are equal numbers of all atoms on each side of the Example… Na + Cl2 → NaCl This reaction shows sodium reacting with chlorine… Notice that there are 2Cl atoms on the left of the arrow and only one on the right. Na + Cl2 → 2NaCl So the equation is not balanced! But what if we add some coefficients? 2Na + Cl2 → 2NaCl Now we have 2 Na’s on each side.. And 2 Cl’s on each side. This equation is now balanced! equation. E. EXAMPLES: 1. Magnesium and hydrogen chloride produce magnesium chloride and hydrogen. 2. Sodium and oxygen yield sodium oxide. 3. Hydrogen and iodine forms hydrogen iodide. 4. Potassium and chlorine make potassium chloride. 3 POINTS TO REMEMBER WHEN WRITING & BALANCING EQUATIONS: You cannot balance an equation using subscripts, this would change the formulas of the compounds. You cannot insert a coefficient between elements in a chemical formula. Write the chemical formula for H 2O as HOH. 3 SUGGESTIONS FOR BALANCING: Balance metal and nonmetal atoms first. Balance hydrogen and oxygen atoms last. Balance polyatomic ions as a group if they appear unchanged on both sides of the equation. II. TYPES OF REACTIONS: A. COMPOSITION (a.k.a. synthesis) reactions occur when two or more reactants combine to create one product. 1. General form: A + B AB 2. EXAMPLES: (Balance them) a. C + O2 CO2 B. b. Mg + O2 MgO c. K KCl + Cl2 DECOMPOSITION reactions occur when one reactant breaks down to form two or more products. There are three types. You will need to memorize the specific products they always form. 1. General form: AB A + B 2. Metal Carbonates: Metal carbonates decomposes into metal oxide and carbon dioxide. a. K2CO3 K2O + CO2 b. CaCO3 CaO + CO2 3. Metal hydroxides: Metal hydroxides decompose into metal oxide and water. a. Mg(OH)2 MgO + HOH b. 2NaOH Na2O + HOH 4. Metal chlorates: Metal chlorates decompose into metal chlorides and oxygen. a. 2 KClO3 2 KCl + 3 O2 b. Ca(ClO3)2 CaCl2 + 3 O2 5. Reversible reactions-some reactions can go either way depending on the conditions. C. SINGLE REPLACEMENT reactions occur when the reactant are one element and a compound. The element replaces one of the elements of the compound. A cation replaces a cation and an anion replaces an anion. 1. General form: cation replaces cation AB + X XB + A anion replaces anion AB + Y AY + B 2. EXAMPLES: (Balance them) a. Zn + HCl ZnCl2 b. Ca + HOH Ca(OH)2 + H2 c. Zn + H2SO4 d. D. KBr + Cl2 + H2 ZnSO4 + H2 KCl + Br2 DOUBLE REPLACEMENT reactions occur when the reactants are two compounds. The compounds “switch partners”. The cation of each compound switch places. The cation is always written first. 1. General form: AB + CD AD + CB 2. EXAMPLES: (balance them) a. H2SO4 b. Al2(CO3)3 c. HCl + CaCl2 CaSO4 + Ca(OH)2 + Ca(OH)2 + HCl CaCO3 + CaCl2 + Al(OH)3 HOH III. Energy Transfers in Chemical Reactions: A. Law of Conservation of Energy states that energy is not created nor destroyed. It is only changed from one form to another. 1. Energy can be stored in bonds and released when those bonds are broken. 2. Energy can be absorbed into bonds from outside. 3. Energy can be in different forms a. Light b. Heat-measured in calories or joules. B. 2 Different Types of Energy Reactions 1. Exothermic reactions involve heat being released (exits). The reaction vessel gets hot. 2. Endothermic reactions involve heat being absorbed (pulled in from the environment). The reaction vessel gets cold. IV. Reaction Rate A. Kinetics is the study of rates of chemical reactions or how fast a reaction occurs. Some factors that affect the reaction rate: 1. Concentration-defines how much of the substance is present in a given space. In general, the more concentrated, the faster the rate. 2. Surface area-describes the contact area the two reacting substances have. Example: stomach vs. small intestine 3. Temperature-the measure of the average kinetic energy of molecules. In general, the more kinetic energy, the faster the rate. 4. Catalysts – substances that speed up a reaction without becoming part of the reaction or being changed. Name_________________________________________Date_________ BALANCING CHEMICAL EQUATIONS ASSIGNMENT UNIT 6 A. Fill in the appropriate coefficient to balance the following equations. ____CaCl2 + ____H2 + _____CO2 1. ____HCl + ____CaCO3 2. _____C3H8 + _____O2 _____CO2 + _____H2O 3. _____AlBr3 + _____K2SO4 _____KBr + _____ Al2(SO4 )3 4. _____H2O2 _____H2O + _____O2 5. _____N2 + _____H2 _____NH3 B. Rewrite the following word equations as balanced chemical equations. On the left write C for composition, D for decomposition, SR for single replacement, DR for double replacement. 6. sodium chlorate when heated yields sodium chloride and oxygen gas. D 2 NaClO3 2 NaCl + 3 O2 7. hydrogen and nitrogen monoxide yields water and nitrogen gas. 8. aluminum and copper (II) chloride yield aluminum chloride and copper. 9. potassium phosphate and calcium chloride yield calcium phosphate and potassium chloride. 10. iron (III) oxide and carbon monoxide yield iron and carbon dioxide. C. Finish the word reaction and predict the products of the reactions below. To the left write the type of reaction. 11. silver nitrate and zinc chloride yields silver chloride and zinc nitrate DR 2AgNO3 + ZnCl2 Zn(NO3)2 + 2 AgCl 12. aluminum and iron (III) oxide yields __________________________________ 13. magnesium bromide and chlorine yields ____________________________________ 14. hydrogen and chlorine yields ____________________________ 15. calcium and oxygen yields ______________________________ 16. sodium and iodine produces ______________________________ 17. iron (II) sulfate and ammonium sulfide forms ______________________________________ 18. potassium iodide and chlorine yield ___________________________________________ 19. barium chloride and sodium sulfate yield ______________________________________ 20. zinc chloride and ammonium sulfide yields ___________________________________ 21. sodium and water yields ______________________________________________ 22. silver and sulfur yields ______________________________________ 23. zinc and lead (II) acetate yields __________________________________________ 24. lead (II) acetate and hydrogen sulfide produces ______________________________ 25. sodium chlorate decomposes into ________________________________________ 26. magnesium chlorate forms _____________________________________________ 27. aluminum chlorate yields ___________________________________________ 28. iron (III) hydroxide yields _______________________________________ 29. barium hydroxide yields _____________________________________________ 30. lithium hydroxide yields ____________________________________________ 31. magnesium carbonate yields ___________________________________________ 32. sodium carbonate yields ________________________________________________ 33. barium carbonate yields ________________________________________________ 34. hydrogen carbonate yields ______________________________________________
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