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How Can We Describe Chemical Reactions?

How Can We Describe Chemical Reactions?. Chemistry Unit 9. Main Ideas. Chemical reactions are represented by balanced chemical equations. There are four main types of chemical reactions: synthesis, combustion, decomposition, and replacement reactions.

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How Can We Describe Chemical Reactions?

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  1. How Can We Describe Chemical Reactions? Chemistry Unit 9

  2. Main Ideas • Chemical reactions are represented by balanced chemical equations. • There are four main types of chemical reactions: synthesis, combustion, decomposition, and replacement reactions. • Double-replacement reactions occur between substances in aqueous solutions and produce precipitates, water or gases.

  3. Reactions and Equations 8.1

  4. Objectives • Recognize evidence of chemical change. • Represent chemical reactions with equations • Balance chemical equations.

  5. Chemical Reactions The process by which one or more substances are rearranged to form different substances is called a chemical reaction. • Evidence of a chemical reaction is a chemical change: a process of changing one or more substances into a new substance. • Evidence: temperature changes, color changes, odor, gas bubbles or appearance of a precipitate

  6. Representing Chemical Reactions Chemical equations – statements that show chemical reactions by the use of chemical formulas and conserved matter with the relative amounts of substances in the reaction. Parts of an equation reaction: • Reactantsare the starting substances. • Productsare the substances formed in the reaction.

  7. Common Symbols

  8. Representing Reactions Word Equations: use of words for reactants and products • aluminum(s) + bromine(l) → aluminum bromide(s) Skeleton Equations: chemical formulas used for reactants and products but not balanced • Al(s) + Br(l) → AlBr3(s) • Skeleton equations lack information about how many atoms are involved in the reaction.

  9. Representing Reactions Chemical Equation: Is a statement that uses chemical formulas to show the identities and relative amounts of the substances involved in a chemical reaction. • 2Al(s) +3Br22AlBr3(s)

  10. Balancing Chemical Equations

  11. Balancing Chemical Equations A coefficientin a chemical equation is the number written in front of a reactant or product, describing the lowest whole-number ratio of the amounts of all the reactants and products. • The most fundamental law is the law of conservation of mass; a balanced equation shows this law.

  12. Balancing Chemical Equations • Write the skeleton equation: • Make sure chemical formulas are correct. • Put in symbols and physical states. liquid sodium carbonate + aqueous calcium chloride yields solid calcium carbonate + aqueous sodium chloride

  13. Balancing Chemical Equations • Count the atoms of the elements in the reactants • Group intact polyatomic ions as a single substance. Na2CO3(l) + CaCl2(aq)  CaCO3(s) + NaCl(aq)

  14. Balancing Chemical Equations • Count the atoms of the elements in the products • Group intact polyatomic ions as a single substance. Na2CO3(l) + CaCl2(aq)  CaCO3(s) + NaCl(aq)

  15. Balancing Chemical Equations • Change the coefficients to make the number of atoms of each element equatl on both sides of the equation • Never change subscripts Na2CO3(l) + CaCl2(aq)  CaCO3(s) + NaCl(aq)

  16. Balancing Chemical Equations • Write the coefficients in their lowest possible ratios Na2CO3(l) + CaCl2(aq)  CaCO3(s) + NaCl(aq)

  17. Balancing Chemical Equations • Go back and check math. Na2CO3(l) + CaCl2(aq)  CaCO3(s) + 2 NaCl(aq)

  18. Balancing Chemical Equations

  19. Example aqueous sodium hydroxide + aqueous calcium bromide yields solid calcium hydroxide and aqueous sodium bromide 2112

  20. Question 1 Which of the following is NOT a chemical reaction? A.a piece of wood burning B.a car rusting C.an ice cube melting into water D.red litmus paper turning blue

  21. Question 1 What is the coefficient of bromine in the equation 2Al(s) + 3Br2(l) → 2AlBr3(s)? A.1 B.2 C.3 D.6

  22. Practice Problems • Page 287 #4-6; page 288 #7-13

  23. Classifying Chemical Reactions 8.2

  24. Objectives • Classify chemical reactions. • Identify the characteristics of different classes of chemical reactions.

  25. Types of Chemical Reactions Chemists classify reactions in order to organize the many types. • Synthesis • Combustion • Decomposition • Single Replacement • Double Replacement (Metathesis)

  26. Synthesis A synthesis reaction is a reaction in which two or more substances react to produce a single product. • When two elements react, the reaction is always a synthesis reaction.

  27. Synthesis A synthesis reaction is a reaction in which two or more substances react to produce a single product. • When two compounds react: • AB + CD  ABCD • AB + BC  ABC

  28. Combustion In a combustion reaction, oxygen combines with a substance and releases energy in the form of heat and light. • Example: Heated hydrogen reacts with oxygen to produce heat and water in a combustion reaction. This is also a synthesis reaction.

  29. Combustion In a combustion reaction, oxygen combines with a substance and releases energy in the form of heat and light. • Element and oxygen react: A + O2 AO • Compound and oxygen react: AB + O2  AO + B

  30. Decomposition A decomposition reactionis one in which a single compound breaks down into two or more elements or new compounds. • Decomposition reactions often require an energy source, such as heat, light, or electricity, to occur.

  31. Decomposition A decomposition reactionis one in which a single compound breaks down into two or more elements or new compounds. • Compound breaks down into two elements: AB  A + B • Compound breaks down to form new compounds: ABCD  AC + BD

  32. Replacement/Displacement A reaction in which the atoms of one element replace the atoms of another element in a compound is called a single replacement reaction. • A + BX → AX + B

  33. Activity Series • A metal will not always replace a metal in a compound dissolved in water because of differing reactivities. • An activity series can be used to predict if reactions will occur.

  34. Activity Series • Halogens frequently replace other halogens in replacement reactions. • Halogens also have different reactivities and do not always replace each other.

  35. Activity Series • Metals/Halogens are listed in order of reactivity . A less reactive metal/halogen will not replace a more reactive metal/halogen

  36. Practice Problems • Page 291 #14-17; page 292 #18-20 • Page 295 #21-24

  37. Double Replacement Double replacement reactions (also called metathesis) occur when ions exchange between two compounds.

  38. Double Replacement • Metathesis reactions often form one of three products: • The solid product produced during a chemical reaction in a solution is called a precipitate. • water – is usually formed with the combination of an acid and a base. A metal salt is also formed. • gas – formed when a gas is not a reactant.

  39. Steps to Metathesis

  40. Product Prediction

  41. Question 1 Which of the following is NOT one of the four types of reactions? A.deconstructive B.synthesis C.single replacement D.double replacement

  42. Question 1 The following equation is what type of reaction? KCN(aq) + HBr(aq) → KBr(aq) + HCN(g) A.deconstructive B.synthesis C.single replacement D.double replacement

  43. Practice Problems • Page 297 #25-28; Page 298 #29-34

  44. Solubility 8.3

  45. Objectives • Identify new possible ionic compounds in a reaction • Define the terms soluble and insoluble • Predict solids based on solubility rules.

  46. Ionic Compounds in Solutions Ionic compounds in aqueous solutions mix and exchange partners (double replacement). • example: Na2SO4(aq) + CaCl2(aq) • Some of these products are solids and some of these products remain aqueous.

  47. Solubility Solubility rules are used to determine the state of matter of products in an aqueous solution. • Soluble means that the compound dissolves in water. • Insoluble means that the compound remains intact in the solid state in water.

  48. Solubility Rules • Most nitrate (NO3-) salts are soluble • Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the ammonium ion (NH4+) are soluble. • Most chloride, bromide, and iodide salts are soluble Exceptions: Ag+, Pb2+, Hg22+

  49. Solubility Rules • Most sulfate salts are soluble . • Exceptions: Bas+, Pb2+, Hg22+, and Ca2+ • Most hydroxides are only slightly soluble (treat as insoluble). • Exceptions: Na+, K+ • Most sulfide (S2-), carbonate (CO32-), chromate (CrO42-), and phosphate (PO43-) salts are only slightly soluble (treat as insoluble). • Exceptions: any containing Alkali metals and ammonium.

  50. Solubility Summary Soluble Insoluble no exceptions no exceptions Ag+, Pb2+, Hg22+ Bas+, Pb2+, Hg22+, Ca2+ Hydroxides (OH-) S2-, CO32-, CrO42- PO43- • Nitrates (NO3-) • Group 1, NH4+ • Halogens • Sulfates (SO42-)

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