Chemical Reactions reactants  products

1. Chemical Reactionsreactants products Balancing chemical equations Types of chemical reactions

2. How reactants are transformed into products? • Reactants are transformed during chemical reactions • Energy is required (absorbed) to break a chemical bond • Energy is released when a chemical bond forms

3. Balancing equations • Obey law of conservation of matter • Chemical equations have two parts: reactants and products

4. How to balance? • The total number of atoms of each element should be the same on both sides of equation • Use coefficients to balance equations Example: 2H2 + O2 2 H2O

5. Types of chemical reactions • Single replacement • Double replacement • Decomposition • Combination • Combustion

6. Synthesis (Combination) Reactions Two or more substances combine to form a new compound. A + X  AX • Reaction of elements with oxygen and sulfur • Reactions of metals with Halogens • Synthesis Reactions with Oxides • There are others not covered here!

7. Decomposition Reactions A single compound undergoes a reaction that produces two or more simpler substances AX  A + X • Decomposition of: • Binary compounds 2H2O(l )  2H2(g) + O2(g) • Metal carbonates CaCO3(s)  CaO(s) + CO2(g) • Metal hydroxides Ca(OH)2(s)  CaO(s) + H2O(g) • Metal chlorates 2KClO3(s)  2KCl(s) + 3O2(g) • Oxyacids H2CO3(aq)  CO2(g) + H2O(l )

8. Decomposition Reactions Sulfates With the exception of alkali metals and alkaline sulfates, all other metals are decomposed by heat to form a metal oxide Nitrates Alkali metals decompose on heating to yield the nitrites and oxygen. All other metal nitrates are decomposed to nitrogen dioxide, oxygen, and the metal oxide on heating.

9. Single Replacement Reactions A + BX  AX + B BX + Y  BY + X Replacement of: • Metals by another metal • Hydrogen in water by a metal • Hydrogen in an acid by a metal • Halogens by more active halogens

10. The Activity Series of the Metals Metals can replace other metals provided that they are above the metal that they are trying to replace. • Lithium • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver • Platinum • Gold Metals above hydrogen can replace hydrogen in acids. Metals from sodium upward can replace hydrogen in water

11. Predict if these reactions will occur Al + MgCl2 3 2 2 3 Mg + AlCl3 Can magnesium replace aluminum? YES, magnesium is more reactive than aluminum. Activity Series No reaction Al + MgCl2 Can aluminum replace magnesium? NO, aluminum is less reactive than magnesium. Therefore, no reaction will occur. Activity Series Order of reactants DOES NOT determine how they react. No reaction MgCl2 + Al The question we must ask is can the single element replace its counterpart? metal replaces metal or nonmetal replaces nonmetal.

12. The Activity Series of the Halogens • Fluorine • Chlorine • Bromine • Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl(s) + F2(g)  2NaF(s) + Cl2(g) ??? MgCl2(s) + Br2(g)  ??? No Reaction

13. 0 Single Replacement Reactions 2 2 • Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g)  NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq) Cu(s) + Al(NO3)3(aq) 2 3 3 2

14. Double Replacement Reactions The ions of two compounds exchange places in an aqueous solution to form two new compounds. AX + BY  AY + BX One of the compounds formed is usually a precipitate, an insoluble gas that bubbles out of solution, or a molecular compound, usually water.

15. 0 Double Replacement Reactions • Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together • Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) • Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

16. Combustion Reactions A substance combines with oxygen, releasing a large amount of energy in the form of light and heat. • Reactive elements combine with oxygen P4(s) + 5O2(g)  P4O10(s) (This is also a synthesis reaction) • The burning of natural gas, wood, gasoline C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

17. Solubility and precipitation reactions Pb(NO3)2 (aq) + 2NaI(aq)  PbI2(s) + 2NaNO3(aq)

18. Na2CO3 FeCO3 FeCl2 NaCl + + (aq) (ppt) Predict if a reaction will occur when you combine aqueous solutions of iron (II) chloride with aqueous sodium carbonate solution. If the reaction does occur, write a balanced chemical equation showing it. (be sure to include phase notation) + iron (II) carbonate iron (II) chloride + sodium carbonate sodium chloride CO32- CO32- Na1+ Fe2+ Fe2+ Cl1- Cl1- Na1+ CO3 Na2 Cl2 Fe NaCl FeCO3 (aq) (ppt) Using a SOLUBILITY TABLE: sodium chloride is soluble iron (II) carbonate is insoluble Balanced chemical equation 2 (aq) (aq) Complete Ionic Equation Fe2+(aq) + 2Cl1-(aq) + 2Na1+(aq) + CO32-(aq) 2Na1+(aq) + 2Cl1-(aq) + FeCO3(s)

19. Solubility rules Soluble in water: • sodium, potassium, and ammonium salts; acetates and nitrates • Halides with the exception of halides of lead (II), silver(I), and mercury(I). • Sulfates with the exception of sulfates of calcium, barium, lead (II) and strontium

20. Insoluble in water • Phosphates, carbonates and sulfides except sodium, potassium, ammonium salts, and calcium sulfide • Hydroxides except sodium, potassium, calcium, and barium hydroxides

21. Acid-Base reactions(neutralization reactions) • Acid: any compound that produces hydrogen ions (H+), when added to water. • Base: any substance that produces hydroxide ions (OH-), when added to water. HCl(aq) + Na(OH)(aq)  NaCl(aq) + H2O(l)