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Today is Thursday, April 17 th , 2014

In This Lesson: Chemical Reactions (Lesson 3 of 4 ). Today is Thursday, April 17 th , 2014. Pre-Class: What’s that? . Today’s Agenda. Where we are and where we’ve been. Chemical Reactions Balancing Chemical Reactions Types of Chemical Reactions Where is this in my book?

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Today is Thursday, April 17 th , 2014

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  1. In This Lesson: Chemical Reactions (Lesson 3 of 4) Today is Thursday,April 17th, 2014 Pre-Class: What’s that? 

  2. Today’s Agenda • Where we are and where we’ve been. • Chemical Reactions • Balancing Chemical Reactions • Types of Chemical Reactions • Where is this in my book? • P. 321 and following…

  3. By the end of this lesson… • You should be able to write, balance, classify, and predict chemical reactions.

  4. A Wide-Range Review • Way back, in a time called the beginning of the semester and a place called here, we talked about matter. • We looked at its forms and properties. • We learned how to measure it and describe it. • We talked about what it’s made of (atoms) and what atoms are made of too.

  5. A Wide-Range Review • Then we talked about electrons. • We talked about where we might find them at any given time and how they’re arranged in elements. • We talked about how chemists arrange elements in the periodic table. • We learned how elements bond with one another.

  6. A Wide-Range Review • We also learned how to name the combinations they form. • We learned another scale for measurement: the mole. • We learned how to measure proportions and write formulas. • Now, we’re going to learn how to write formulas for entire chemical reactions.

  7. Chemical Reactions • A chemical reaction is when a set of chemicals is changed into another set of chemicals. • Reactions can be endergonic or exergonic. • Endergonic: Energy absorbed. • Exergonic: Energy released. • Exergonic reactions can happen spontaneously. • More on this to come in the next unit.

  8. Chemical Reactions • Previously, we discussed how reactions are shown in basic form: • Reactants (starting stuff) are shown on the left of the equation. • Products (ending stuff) are shown on the right of the equation. • The arrow means “yields.” • Example: • Reactant(s) Product(s) • 4Fe + 3O2 2Fe2O3

  9. Just checking… • Identify the reactant(s): • Na + ClNaCl • Na, Cl • Identify the reactant(s): • 2H2O  2H2 + O2 • H2O • Identify the product(s): • Na + ClNaCl • NaCl • Identify the product(s): • 2H2O  2H2 + O2 • H2 + O2

  10. One other thing… • Don’t forget the symbols we covered back in the beginning of the year: • (s) means a chemical is in solid form. • (l) means a chemical is in liquid form. • Most liquids will be H2O for us. • (g) means a chemical is in gaseous form. • Watch for BrINClHOF elements! • (aq) means a chemical is in aqueous form – dissolved in water. • Acids are always (aq).

  11. Lastly… • Don’t forget the BrINClHOF (diatomic) elements: • Bromine (Br2) • Iodine (I2) • Nitrogen (N2) • Chlorine (Cl2) • Hydrogen (H2) • Oxygen (O2) • Fluorine (F2) • When these elements are on their own, they bond to themselves. YOU MUST REMEMBER THIS!!!!!!!!!!!!!!!!!!!!!1

  12. Types of Reactions • Now let’s talk about the types of chemical reactions. • For this class, you’ll need to know five of them: • Combination Reactions (also known as Synthesis) • Decomposition Reactions • Single Replacement Reactions • Double Replacement Reactions • Combustion Reactions • Let’s take a look at them in chemistry terms as well as…prom terms.

  13. 1: Synthesis +  “I told you they were together!”

  14. 1. Combination (Synthesis) Reactions • In combination/synthesis reactions, two or more chemicals combine to make a new compound. • A + X  AX • Examples include: • Reactions with oxygen and sulfur. • Reactions of metals with halogens. • Reactions with oxides.

  15. 2: Decomposition +  “Well we all saw that coming.”

  16. 2. Decomposition Reactions • In decomposition reactions, a single compound breaks down into two or more simpler substances. • AX  A + X • Examples include (KNOW THESE): • [metal]CO3 [metal]oxide + CO2 • [metal]OH [metal]oxide + H2O • [metal]ClO3 [metal]chloride + O2 • Acids  [acid]oxide + H2O

  17. 3: Single Replacement + +  “Scandalous!”

  18. 3. Single Replacement Reactions • In single replacement reactions, a lone element takes the place of an element in a compound. • A + BX  AX + B • BX + Y  BY + X • Examples include: • Metals replacing metals. • Hydrogen in water being replaced by a metal. • Hydrogen in acid being replaced by a metal. • Halogens being replaced by more reactive halogens.

  19. Activity Series • When we talked about Single Replacement Reactions, I mentioned “more reactive halogens.” • As it turns out, elements (not just halogens) have varying degrees of reactivity. • Chemists have created lists called Activity Series that detail how (in our class’s case) metals and halogens react with one another, assuming they even react at all. • Also known as Reactivity Series. Makes sense…

  20. Activity Series of Metals • Lithium • Potassium • Strontium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • HYDROGEN • Bismuth • Copper • Mercury • Silver • Platinum • Gold • Metals can replace other metals if they’re above the metal they’re replacing. • Metals can replace Hydrogen in acids if they’re above Hydrogen. • Metals can replace Hydrogen in water if they’re above Magnesium.

  21. Activity Series of Halogens • A halogen can replace another halogen in a compound if it is above the one it’s replacing. • Example: • 2NaCl (s) + F2 (g)  2NaF (s) + Cl2 (g) • MgCl2 (s) + Br2 (g) NO REACTION • Note that the halogen activity series is the same as the group order of halogens on the table. • Fluorine • Chlorine • Bromine • Iodine

  22. Activity Series: Reaction or No? • Cu + MgSO4 Mg + CuSO4 • No reaction (Magnesium is above Copper) • Pb + ZnSO4 Zn + PbSO4 • No reaction (Zinc is above Lead) • Fe + 2AgNO3 2Ag + Fe(NO3)2 • Reaction (Iron is above Silver) • 2Al + 3H2O  Al2O3 + 3H2 • No reaction (Aluminum is not above Magnesium)

  23. Single Replacement Technicalities**Technically, this is important. • When a Group I or Group II metal (alkali/alkaline earth) reacts with water, they only replace one of the hydrogens. • Examples: • K + H2O  KOH + O2 • Mg + H2O  Mg(OH)2 + O2

  24. Aside: Coins • There’s an interesting phenomenon with pocket change relating to the Activity Series: • Pennies tend to become very dull relatively quickly, yet quarters and other “silvery” coins tend not to. What’s the deal? • As it turns out, pennies are plated in copper, while other coins are plated in nickel. • Because nickel is higher on the list, it’s less likely to be replaced and thus tarnish. • Copper, on the other hand, is a bit of a chemical “weakling.” http://www.usmint.gov/about_the_mint/index.cfm?action=coin_specifications http://www.harpercollege.edu/tm-ps/chm/100/dgodambe/thedisk/series/3postlab.htm

  25. 4: Double Replacement + +  “Gross!”

  26. 4. Double Replacement Reactions • In double replacement reactions, ions of two compounds flip places in aqueous solutions, forming two new compounds. • AX + BY  AY + BX • Typically, one of the compounds formed is: • A precipitate (a solid or a bubblin’ gas). • A molecular compound (usually water). • One of the compounds must be insoluble!

  27. Reminder:Dissociation when Dissolved Bound ions in… …component ions out. Cl Pb Cl Cl- Pb2+ Cl-

  28. Predicting States of Matter • Remember that double replacement reactions occur in solutions. • To predict the states of matter resulting from a double replacement reaction, first write the equation. • Then, use your solubility table. • FYI, when they say “salts involving,” just think of it as saying “ionic compounds involving…” • FYI, when they say “halides,” just think of it as saying “halogens…”

  29. Solubility Example • NiNO3 (aq) + KBr (aq) → KNO3 (?) + NiBr(?) • In what states are potassium nitrate and nickel (I) bromide? • According to your solubility table: • 1./2. All salts of Group IA and nitrates are soluble, so potassium nitrate is. • 3. All salts of halides (halogens) are soluble, except… so nickel (I) bromide is. • NiNO3 (aq) + KBr (aq) → KNO3 (aq) + NiBr(aq) • So no reaction, since both of them are soluble. • Remember, one must be insoluble!

  30. Solubility Example • Ba(NO3)2 (aq) + (NH4)3PO4 (aq) → NH4NO3 (?) + Ba3(PO4)2 (?) • In what states are ammonium nitrate and barium phosphate? • According to the solubility table, • 1./2. All salts of ammonium and nitrates are soluble so ammonium nitrate is. • 5. All salts of …phosphate… are insoluble except… so barium phosphate is not. • Ba(NO3)2 (aq) + (NH4)3PO4 (aq) → NH4NO3 (aq) + Ba3(PO4)2 (s) • One compound is insoluble, so there will be a reaction.

  31. Precipitates • In the previous example, Ba3(PO4)2 “fell out” of solution. • In other words, it took on a solid form and was no longer dissolved. • We would expect it to collect at the bottom of the container. • This is an example of precipitation, or the formation of a precipitate. • A precipitate is a solid or gas substance that “falls out” of an aqueous solution. • A precipitate could be water, but this is less common.

  32. Precipitate Video/Demo • KI (aq) + Pb(NO3)2 (aq)  KNO3 (?) + PbI2 (?) • Check your solubility tables for the phase of the two products. • 1./2. All salts of Group IA and nitrates are soluble, so potassium nitrate is. • 3. All salts of halides are soluble except those of lead (II), so lead (II) iodide is insoluble. • KI (aq) + Pb(NO3)2 (aq)  KNO3 (aq) + PbI2 (s) • [Video]

  33. Remembering Solubility Rules? • The Solubility Song! • Link available in Chemistry Links. • Lyrics available in Worksheets and Keys.

  34. 5: Combustion +  This one’s hard to picture. Basically, oxygen reacts with something, usually releasing a lot of light and/or heat.

  35. 5. Combustion Reactions • In combustion reactions, a substance reacts with oxygen, releasing a large amount of energy in the form of heat and light. • When the reactants are only oxygen and a hydrocarbon, carbon dioxide and water are the products. • Examples include: • C3H8 (g) + 5O2 (g)  3CO2 (g) + 4H2O (g) • P4 (s) + 5O2 (g)  P4O10 (s) • This is a combustion and synthesis reaction!

  36. Combustion Reaction Demo • C2H5OH + 3O2 2CO2 + 3H2O

  37. Aside: Great Moments in Science • Meet Pilatre de Rozier: • Mr. Rozier wanted to test the flammability of hydrogen, so he inhaled some, then exhaled over an open flame. • Result? • Singed eyebrows. http://www.sciencephoto.com/image/228189/350wm/H4180269-Pilatre_de_Rozier,_French_balloonist-SPL.jpg

  38. Identifying Chemical Reactions • Let’s practice identifying chemical reactions: • Chemical Reactions Packet, Page 2, Upper Section • Don’t worry about balancing them yet.

  39. Balancing Chemical Reactions • In addition to identifying chemical reactions, they also need to be balanced. • According to the Law of Conservation of Mass/Matter, the mass of the reactants must equal the mass of the products. • Atoms are conserved. • So, all chemical formulas must show the same AMOUNTS OF ATOMS on both sides of the arrow. • No elements can appear or disappear, either.

  40. PhET • Balancing Chemical Equations

  41. Skeleton Equations • Up till this point in the year, sometimes we’ve been writing equations that are not balanced, just to describe which elements are reacting. • These unbalanced equations are called skeleton equations. • Think “bare bones” equations. • From now on, we’ll need to balance our equations, so here are some directions.

  42. How to Balance Chemical Equations • Under the arrow, vertically list each element. • Don’t use any additional subscripts. • Put a box around each term in the equation. • Use coefficients to balance each side. • NOT subscripts. • Balance hydrogen second-to-last and oxygen last. • How to remember this?

  43. Balancing Chemical Equations Example 4 2 3 2 • ___Al + ___O2 ___Al2O3 Al O 1 4 2 2 4 6 2 6 3

  44. Important Note • Keep in mind that, like in empirical formulas, the coefficients in a balanced equation should not be able to be reduced. • In other words: • 2Mg + 2O  2MgO should really be • Mg + O MgO • Even if it’s balanced, it has to be reduced to the lowest ratio.

  45. Balancing Synthesis Reactions1 of 3 • ___CaO + ___H2O  ___Ca(OH)2 • CaO + H2O Ca(OH)2 • ___P4 + ___O2 ___P2O5 • P4 + 5O2 2P2O5 • ___Ca + ___O2 ___CaO • 2Ca + O2 2CaO • ___Cu + ___S8  ___ CuS • 8Cu + S8  8CuS

  46. Balancing Synthesis Reactions2 of 3 • ___S8 + ___O2 ___SO2 • S8 + 8O2 8SO2 • ___H2 + ___N2 ___NH3 • 3H2 + N2 2NH3 • ___H2 + ___Cl2 ___HCl • H2 + Cl2 2HCl • ___Ag + ___S8 ___Ag2S • 16Ag + S8 8Ag2S

  47. Balancing Synthesis Reactions3 of 3 • ___Cr + ___O2 ___Cr2O3 • 4Cr + 3O2 2Cr2O3 • ___Al + ___Br2 ___AlBr3 • 2Al + 3Br2 2AlBr3 • ___Na + ___I2 ___NaI • 2Na + I2 2NaI • ___H2 + ___O2 ___H2O • 2H2 + O2 2H2O

  48. Balancing Decomposition Reactions1 of 3 • ___BaCO3 ___BaO + ___CO2 • BaCO3BaO + CO2 • ___MgCO3 ___MgO + ___CO2 • MgCO3MgO + CO2 • ___K2CO3 ___K2O + ___CO2 • K2CO3 K2O + CO2 • ___Zn(OH)2 ___ZnO + ___H2O • Zn(OH)2ZnO + H2O

  49. Balancing Decomposition Reactions2 of 3 • ___Fe(OH)2 ___FeO + ___H2O • Fe(OH)2FeO + H2O • ___Ni(ClO3)2 ___NiCl2 + __O2 • Ni(ClO3)2 NiCl2 + 3O2 • ___NaClO3 ___NaCl + ___O2 • 2NaClO3 2NaCl + 3O2 • ___KClO3 ___KCl + ___O2 • 2KClO3 2KCl + 3O2

  50. Balancing Decomposition Reactions3 of 3 • ___H2SO4 ___H2O + ___SO3 • H2SO4 H2O + SO3 • ___H2CO3 ___H2O + ___CO2 • H2CO3 H2O + CO2 • ___Al2O3 ___Al + ___O2 • 2Al2O3 4Al + 3O2 • ___Ag2O  ___Ag + ___O2 • Ag2O  Ag + O2

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