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Chapter 11 Chemical Reactions 11.1 Describing Chemical Reactions 11.2 Types of Chemical Reactions 11.3 Reactions in Aque

Chapter 11 Chemical Reactions 11.1 Describing Chemical Reactions 11.2 Types of Chemical Reactions 11.3 Reactions in Aqueous Solutions The objective of this chapter is for you to identify a type of equation, predict the product from the reactants, and balance the final equation.

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Chapter 11 Chemical Reactions 11.1 Describing Chemical Reactions 11.2 Types of Chemical Reactions 11.3 Reactions in Aque

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  1. Chapter 11 Chemical Reactions 11.1 Describing Chemical Reactions 11.2 Types of Chemical Reactions 11.3 Reactions in Aqueous Solutions The objective of this chapter is for you to identify a type of equation, predict the product from the reactants, and balance the final equation.

  2. Writing Chemical Equations There are two parts to a chemical equation a. Reactants – those elements or compounds that will combine together to form new compounds or molecules. Always on the left side of the equation. b. Products – those new elements or compounds that form in a chemical reaction. Always on the right side of the equation. Reactants Product The arrow means “yields”

  3. Word Equations: Consist of writing the names of the reactants on the left side of the yield sign and the products on the right side of the yield sign Iron + Oxygen Iron (III) oxide

  4. Chemical Equations: Is a representation of a chemical reaction using the formulas of the reactants and products. Fe + O2 Fe2 O3 + Used to separate reactants and products Used for reversible reactions (s) Designates a solid (l) Designates a liquid (g) Designates a gas (aq) Designates an aqueous solution; the substance is dissolved in water D Indicates heat is supplied to the reaction MnO2 A formula written above or below the yield sign indicates its use as a catalyst – ( a substance that speeds a reaction)

  5. Skeleton Equation: Is a chemical equation that does not indicate the relative amounts of the reactants and products Balanced Equation: Is a chemical equation in which each side of the equation has the same number of atoms of each element and mass is conserved. To balance the equation, whole numbers called coefficients are used on both sides of the equations to help balance the number of atoms in the reactants and product. The coefficient actually represents how many moles are present in the reactant and product that are necessary to run the reaction. Link

  6. Steps for writing Chemical Equations: 1. Write a skeleton equation 2. Use subscripts to balance charges in the product 3. Use coefficients to balance atoms on each side of the equation Try Potassium + Oxygen Potassium Oxide K + O2 K+1O+2 K + O2 K2O 4K + O2 2K2O Now Try: P + O2 P4O10 AgNO3 + Cu Ag + Cu(NO3)2

  7. Law of the Conservation of Mass: The mass of the atoms present in the reactants must equal the mass of the atoms present in the product

  8. 11.2 Types of Chemical Reactions • The following is a list of the four major types of reactions • By knowing the type of reaction, the products can be predicted. • Composition Reaction • Decomposition Reaction • Replacement Reaction • Combustion

  9. Combination Reaction: Also called a Synthesis Reaction, occurs when two or more substances combine to form a more complex substance. Composition reactions have the general form; A + X AX Examples: Iron and Sulfur combine to form Iron(II) Sulfide Fe + S FeS Magnesium and Oxygen gas form Magnesium Oxide 2Mg + O2 2MgO Water and Sulfur Trioxide H2O + SO3 H2SO4

  10. Two Special Combination Reactions: • Metal Oxide + Water Hydroxides (which are bases) • Na2O + H2O 2 NaOH (Sodium Hydroxide) • Try, CaO + H2O ? • 2. Nonmetal oxide + Water Acids • SO3 + H2O H2SO3 (Sulfurous Acid) • Try, Cl2O5 + H2O ? • Don’t remember your acids and bases then review chapter 9!

  11. Decomposition Reaction: Reactions that are in reverse to decomposition reactions. Here one substance breaks down to form two or more simpler substances. Decomposition reactions have the general form; AX A + X Examples: Water decomposes, yielding hydrogen and oxygen 2H2O 2H2(g) + O2(g) Potassium Chlorate decomposes, yielding potassium chloride and oxygen 2KClO3 2KCl + 3O2(g) Mercury(II) Oxide decomposes to form metallic mercury and oxygen 2HgO 2Hg(l) + O2(g)

  12. There are six types of decomposition reactions: • Metallic carbonates, when heated, form metallic oxides and carbon dioxide CaCO3 CaO + CO2(g) • Many metallic hydroxides, when heated, decompose into metallic oxides and water. Ca(OH)2 CaO + H2O(g) • Metallic chlorates, when heated, decompose into metallic chlorides and oxygen. 2KClO3 2KCl + 3O2(g) • Some acids, when heated, decompose into nonmetallic oxides and water. H2CO3 H2O + CO2(g) • Some oxides, when heated, decompose though most are stable. • 2HgO 2Hg + O2(g) • 6. Some decomposition reactions are produced by an electric current • 2H2O (electricity) 2H2(g) + O2(g)

  13. Replacement Reaction: • Occur when one substance is replaced in its compound by another substance. • Replacement reactions have the general form; • Single Replacement- A + BX AX + B • OR • Y + BX BY + X • Double Replacement – AY + BX AX + BY • There are four specific types of replacement reactions: • Replacement of Hydrogen in water by metals • Replacement of a metal in a compound by a more active metal • Replacement of Hydrogen in acids by metals • Replacement of Halogens

  14. Reactivity of the elements determine if the reactions will occur. One atom must be more reactive then the element that is being replaced in the equation.

  15. Replacement of Hydrogen in water by metals: The very active metals such as potassium, calcium, and sodium, react vigorously with water. They replace half the hydrogen to form metallic hydroxides. At elevated temperatures less active metals such as magnesium, zinc, and iron react with steam to replace hydrogen. Because of the high temperature involved, oxides rather than hydroxides are formed. Metals less active than iron do not react measurably with water. Example: Ca + 2H2O Ca(OH)2 + H2(g)

  16. Replacement of a metal in a compound by a more active metal: • A more reactive metal replaces the less reactive metal in a compound • It is important to understand the periodic trends for the reactivity of metals • Example: Zn + CuSO4 ZnSO4 + Cu(s)

  17. Replacement of Hydrogen in acids by metals: Many metals react with certain acids to replace the hydrogen in the acid to form a metallic compound. Metals from Li to Na will Replace hydrogen from water and acids. Metals from Mg to Pb Will replace hydrogen from acids only. Example: Zn + H2SO4 ZnSO4 + H2(g)

  18. Replacement of Halogens: Replacement of a halogen with another halogen depends on the reactivity of the two halogens involved. A more reactive halogen always replaces a less active halogen. Cl2 + 2KBr 2KCl + Br2 How does the reactivity of the halogens progress?

  19. Double Replacement Reaction: • An exchange of positive ions between two compounds. There are generally three rules that govern this type of reaction. These reactions are generally ionic in nature and take place • in an aqueous solution. To occur, one of the products must be • a. An insoluble precipitate b. A gas c. A molecular compound • Double Replacement – AY + BX AX + BY • One of the products is only slightly soluble and precipitates from solution. Na2S + Cd(NO3)2(aq) CdS + 2NaNO3(aq) • One of the products is a gas • 2NaCN(aq) + H2SO4(aq) 2HCN(g) + Na2SO4(aq) • 3. One product is a molecular compound such as water • Ca(OH)2(aq) + 2HCl(aq) CaCl2(aq) + 2H2O(l)

  20. Combustion Reaction: Occurs when an element or compound reacts with oxygen, often producing energy in the form of light and heat. The reaction involves oxygen as a reactant while the other reactant is often a hydrocarbon. In this case the complete combustion of a hydrocarbon produces carbon dioxide and water. 2C8H18(l) + 25O2(g) 16CO2(g) + 18H2O(l) Other elements can be combusted with oxygen and look much like a combination reaction 2Mg(s) + O2(g) 2MgO(s) A hydrocarbon is a compound composed only of hydrogen and carbon Many are used as fossil fuels – methane, propane, butane, and octane

  21. 11.3 Reactions in Aqueous Solution: • Net Ionic Reactions – • The earth is 70% water • Your body is 66% water • Many important chemical reaction take place in water (an aqueous solution) causing the compounds to separate into ions • Example: When sodium chloride and silver nitrate are placed in solution, the ions dissociate. You can use these ions to write a complete ionic equation. • Complete ionic equation – an equation that shows dissolved ionic compounds as dissociated ions. • Ag+(aq)+NO3-(aq)+Na+(aq)+ Cl-(aq) AgCl(s)+ Na+(aq)+NO-3(aq)

  22. Ag+(aq)+NO3-(aq)+Na+(aq)+ Cl-(aq) AgCl(s)+ Na+(aq)+NO-3(aq) Note that the sodium and nitrate ion are unchanged, the equation can be simplified by eliminating these ions because they do not participate in the reaction Ag+(aq) + Cl-(aq) AgCl(s) This is called the net ionic reaction – an equation that shows only the particles involved in the chemical change Spectator Ions – an ion that appears on both sides of an equation that is not directly involved in the reaction.

  23. Predicting the Formation of Precipitate: You can predict the formation of a precipitate by using the general rules for solubility of ionic compounds.

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