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Today-- Welcome Back!

Today-- Welcome Back!

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Today-- Welcome Back!

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  1. Today-- Welcome Back! • Turn in “Chapter 6 Extra Credit” • Unexpected Changes Lab • Introduce Chapter 9 (LAST CHAPTER) • Outlining – NOPE!

  2. Chapter 9 – Chemical Reactions and Equations

  3. Learning Objectives: • Be able to balance chemical equations by applying the law of conservation of mass. • Be able to recognize synthesis, decomposition, single replacement, double replacement, combustion and neutralization reactions.

  4. Law of Conservation of Matter • Conservation of Matter: in all chemical and physical changes, matter is neither created or destroyed • The total mass in a chemical reaction remains constant • Antoine Lavoisier: • Made accurate and precise measurements during chemical reactions

  5. Reactants ―› Products • Reactants: the substances that enter into the reaction • Products: the substances that are produced by the reaction

  6. Why Do Reactions Occur • Think back to what we know about atoms and specifically their electrons? • Through chemical reactions, atoms have the opportunity to obtain complete sets of valence electrons and thus become more stable.

  7. Five General Types of Chemical Reactions • Direct Combination (Synthesis) • Decomposition • Single-Replacement • Double-Replacement • Combustion By knowing the type of reaction that is occurring, you can predict the products that will be formed.

  8. Chemical Equations • A method of describing chemical reactions • Word Equations • Calcium + Oxygen ―› Calcium Oxide • Formula Equations • 2Ca + O2 ―› 2 CaO the arrow → separates the reactants from the products

  9. Completing the Chemical Equation • Complete the chemical equation by describing the physical state of each substance: • Solid (s) • Liquid (l) • Gas (g) • Aqueous (aq) means dissolved in water

  10. Symbols used in equations • Double arrow indicates a reversible reaction • shows that heat is supplied to the reaction • is used to indicate a catalyst is supplied, in this case, platinum. ∆ Pt

  11. What is a catalyst? • A substance that speeds up a reaction, without being changed or used up by the reaction. • Enzymes are biological catalysts. • How can you physically speed up a reaction?

  12. I. Direct Combination Reactions (also called synthesis reactions). General form: A + B → AB (two reactants make a single product) A, B = elements or compounds AB = compound consisting of A and B • This is the only type of chemical reaction in which there is a single product formed. This single product is always more complex than the reactants.

  13. Examples of Synthesis Reactions • calcium + oxygen yields calcium oxide • 2Ca + O2 → 2CaO • Notice: All equations show two (or more) reactants, but only one product. • http://www.ric.edu/ptiskus/reactions/Index.htm

  14. II. Decomposition Reactions General form: AB → A + B (one reactant makes two or more products) AB = compound A, B = elements or simpler compounds • This is the only type of chemical reaction in which there is a single reactant. This single reactant is always more complex than the products.

  15. Decomposition Reactions: Examples • water yields hydrogen and oxygen 2H2O → 2H2 + O2 • marble (calcium carbonate) yields calcium oxide and carbon dioxide CaCO3 → CaO + CO2 • Notice: all equations show a single reactant decomposing into two (or more) products. • http://www.ric.edu/ptiskus/reactions/Index.htm

  16. Balancing Chemical Equations • The Law of Conservation of Matter states that: • Matter is neither created nor destroyed! • For mass to remain constant both before and after a reaction, the number of atoms must remain constant

  17. Step 1: Balancing Equations • Write the word equation that describes the reaction. iron + oxygen ―› iron oxide

  18. Step 2: Balancing Equations • Replace the words in the equation with symbols and formulas. Fe + O2 ―› Fe2O3 Do we have the same numbers of each atoms on both sides of arrow? Does this follow the law of conservation of matter?

  19. Step 3: Balancing Equations • Count the # of atoms of each element on both sides of the equation. Fe + O2 ―› Fe2O3

  20. Step 4: Balancing Equations • Starting with elements that only occur in one substance on each side of the equation, make sure that each side of the equation has an equal # of that element. Proceed with all elements. Remember that changing the # of one element may alter elements that have already been balanced.

  21. Fe + O2 ―› Fe2O3

  22. Let’s try: CH4+O2 ―› CO2 + H2O

  23. Never • Never change a subscript to balance an equation. • If you change the formula you are describing a different reaction. • H2O is a different compound than H2O2 • Never put a coefficient in the middle of a formula • 2 NaCl is okay, Na2Cl is not.

  24. Balancing Equations: Examples • H2 + O2 → H2O • Co + O2 → Co2O3 • Pb(NO3)2 + K2S → PbS + KNO3 • C2H6 + O2→ H2O + CO2

  25. Balance the following iron(II) chloride + sodium phosphate → sodium chloride + iron (II) phosphate FeCl2 + Na3PO4→ NaCl + Fe3(PO4)2

  26. Today • Look at Single-Replacement Reactions. • Begin “Single-Replacement Lab” set-up.

  27. Single-Replacement Reactions • Copper metal and silver nitrate: • Cu(s) + AgNO3(aq) → Ag(s) + CuNO3(aq) • What do you observe about the reaction? • What do you notice about the chemical equation? • Cu must be more reactive than Ag in order for the reaction to take place.

  28. Single-Replacement Reactions General Form: A + BX → AX + B One element and one compound recombine (switch partners) AX, BX = ionic compounds A, B = Metals X = ion that switches partners *Metal ‘A’ must be more reactive than ‘B’ for this to occur

  29. Single-Replacement Lab • Today you will do the following: • 1. Formulate a question for the lab • 2. Formulate a hypothesis • 3. Design procedures • 4. Create a data table.

  30. IV. Double-Replacement Reactions General form: AX + BY → AY + BX (Positive ions in two compounds are exchanged) A,B = positive ions X,Y = negative ions • This is the only type of chemical reaction with two compounds as reactants and two compounds as products.

  31. Double Replacement Examples • calcium carbonate and hydrochloric acid yield calcium chloride and carbonic acid CaCO3 + 2HCl → CaCl2 + H2CO3 • Notice: in this reaction, two ionic compounds exchange ions to form two new ionic compounds www.ric.edu/ptiskus/reactions/Index.htm

  32. IV. Double-Replacement Reactions General form: AX + BY → AY + BX (Positive ions in two compounds are exchanged) A,B = positive ions X,Y = negative ions • This is the only type of chemical reaction with two compounds as reactants and two compounds as products.

  33. Double Replacement Examples • calcium carbonate and hydrochloric acid yield calcium chloride and carbonic acid CaCO3 + 2HCl → CaCl2 + H2CO3 • Notice: in this reaction, two ionic compounds exchange ions to form two new ionic compounds www.ric.edu/ptiskus/reactions/Index.htm

  34. Rules of Double-Replacement Reactions • Reactants must be dissolved in water (releasing the ions). • Will occur if one of the products : • is a molecule (covalent), • a precipitate (solid comes out of solution), or • an insoluble gas.

  35. V. Combustion Reactions General Form: CxHy + O2 → H2O + CO2 (hydrocarbon and oxygen react to form carbon dioxide and water) • This is the only type of chemical reaction where something reacts with oxygen and forms carbon dioxide and water

  36. Combustion Examples • Methane reacts with oxygen: CH4 (methane) + O2 → H2O + CO2 • Gasohol reacts with oxygen: C2H5OH (ethanol) + O2 → H2O + CO2 • Notice: in both cases, water and carbon dioxide are the products. www.ric.edu/ptiskus/reactions/Index.htm

  37. 1. Write the word equation2. Write the balanced formula equation • Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas.

  38. 1. Write the word equation2. Write the balanced formula equation • Nitric acid reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate.