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Types of Chemical Reactions: Synthesis, Combustion, Decomposition, Double-Replacement, Single-Replacement

Learn about the different types of chemical reactions, including synthesis, combustion, decomposition, double-replacement, and single-replacement. Understand their definitions, examples, and how to write and balance chemical equations.

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Types of Chemical Reactions: Synthesis, Combustion, Decomposition, Double-Replacement, Single-Replacement

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  1. Unit Four Chemical Reactions

  2. Chemical Reactions • Process by which the atoms of one or more substances are rearranged to form different substances

  3. 5 Types of Chemical Reactions • Synthesis • Combustion • Decomposition • Double-Replacement • Single-Replacement

  4. Synthesis Reactions • What does synthesis mean? • A + B AB • 2 or more compounds or elements go together to build a more complex compound • Which example is a synthesis equation?

  5. Combustion Reactions • What is necessary for something to burn? • Oxygen combines with a substance and releases energy in the form of heat and light • OXYGEN IS ALWAYS A REACTANT! • Which example is a combustion equation?

  6. Decomposition Reactions • What does decompose mean? • A single compound breaks down into 2 or more elements or new compounds • AB A + B • Which example is a decomposition equation?

  7. Double Replacement Reactions • Exchange of ions between 2 ionic compounds • AX + BY AY + BX • Which example is a double replacement equation?

  8. Single Replacement Reactions • One element replaces another element in a compound • A + BX AX + B • A metal will not always replace another metal…it depends on how active they are…most active replaces least active • Which example is a single replacement equation?

  9. Summary of Reaction Types • Funny Analogies

  10. What do all the small letters mean? • (s) solid • (l) liquid • (g) gas • (aq) dissolved in water

  11. Terminology • Word Equations: statements that describe a reaction • Skeleton Equations: uses formulas not words to describe a reaction-NOT BALANCED • True Chemical Equation: a balanced skeleton reaction that follows the law of conservation of matter

  12. Writing Chemical Equations • Read a description of the reaction • Note what is reacted with what • Note what is yielded or produced • Write formulas for each compound REMEMBER TO CRISS-CROSS IF IONIC!

  13. Balancing Chemical Equations • Reflects the law of conservation of mass which says…

  14. Steps to Balancing an Equation • Write the skeleton equation. BE SURE THE FORMULAS ARE WRITTEN CORRECTLY. • Inventory reactants • Inventory products • Add coefficients to make atoms of each element equal on both sides of the equation • Reduce the coefficients if possible

  15. Examples • www.chemfiesta.com

  16. Write and Balance the Following • Magnesium sulfate + calcium hydroxide yields magnesium hydroxide + calcium sulfate • Iron + silver chloride yields silver + iron (III) chloride

  17. So….. In a lab with beakers and chemicals, how do I know that a reaction is happening?

  18. Evidence of Chemical Reactions • Temperature change exothermic, endothermic • Release energy in the form of light • Color change • Odor • Gas bubbles • Formation of a solid…called a precipitate.

  19. Reactions in Aqueous Solution • More than 70% of earth is covered by water • 66% of the human body is water • MANY chemical reactions occur in water • Often form solids called precipitates • Remember: Ionic compounds dissociate in water.

  20. Net Ionic Equations • Represent reactions of ionic compounds in aqueous solution by writing complete ionic equations • Remove spectator ions (those appear on both sides of the equation but aren’t in the precipitate) • Leaves a net ionic equation that can be balanced

  21. Percent Composition • The relative amounts of each element in a compound are expressed in percent composition. AKA: percent by mass of each element • % of element = grams of element X 100 grams of compound EXAMPLES: pg. 191

  22. Calculating Empirical Formula • Empirical formula: lowest whole number ratio of the atoms of the elements in a compound • Empirical formula doesn’t have to be the same as the actual molecular formula of the compound. Ex. Hydrogen peroxide

  23. Calculating Empirical Formula • Given % composition. • Assume 100 grams. Allows you to change the % to g. • Convert to moles. • Find lowest number of moles. • Divide all moles by the lowest number. • Multiply by a number if necessary to get whole numbers. • Example: pg. 193

  24. Calculating Molecular Formulas • Molecular Formula= n(empirical formula) n= molecular formula mass/molar mass of empirical Examples: pg. 194

  25. Avogadro’s Number: 1 mole = 6.02 x 1023 particles (particles= atoms, formula units, or molecules) Molar Mass: 1 mole = ________ grams of an element or compound (can be found on the periodic table) Introducing the Mole

  26. Chemical Quantities Groups • Gather 2 sheets of paper, periodic table, calculator, notes, and long WS • Instructions: • Jar Lab: Number paper 1 to 5. Calculate the # of grams in each jar. Be sure to write the given, unknown, and conversion factor • % Composition and Empirical Formula Cards: Do 3 cards. Be sure to write down the #’s. • IF DONE: Begin writing the equations for the LONG WS 26-36.

  27. Stoichiometry We cannot get rid of the mole!

  28. What is Stoichiometry? • Quantitative relationships between reactants used and products formed • RELATES ONE COMPOUND TO A DIFFERENT COMPOUND • Based on law of conservation of matter

  29. What is Stoichiometry? • CANNOT DO STOICHIOMETRY WITHOUT A BALANCED EQUATION! • Why? The coefficients in a balanced equation tell us how many moles of each compound are used or produced in the reaction.

  30. What is Stoichiometry? • Mole Ratio: ratio between the number of moles of any two compounds in a balanced chemical equation • COEFFICIENTS

  31. Stoichiometric Calculations • 2 relationships in stoichiometry: • Molar mass • Mole ratios

  32. Stoichiometric Calculations • Mole to mole conversions • Mole ratio • Mole to mass conversions • Mole ratio and molar mass • Mass to mass conversions • Molar mass, mole ratio, molar mass

  33. Limiting Reactants • Limiting reactant: ends the reaction because it runs out • Excess reactant: left over at the end of the reaction • AMOUNT OF PRODUCT THAT YOU GET IN A REACTION DEPENDS ON THE LIMITING REACTANT…BECAUSE IT STOPS THE REACTION!

  34. Limiting Reactants • Finding which is limiting: • Will be given two things • Solve for unknown twice (one time using each given) • Which is the smallest answer? That is the answer to the problem. • The limiting reactant is the given found at the front of that solution!

  35. Limiting Reactants • Since the limiting reactant stops the reaction, USE THE LIMITING REACTANT TO SOLVE FOR THE AMOUNTS OF PRODUCTS PRODUCED.

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