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Advanced Physical Science B

Advanced Physical Science B. Chemistry in Review: The Ions Chemical Formulas Chemical Equations Stoichiometry. The Ions. Ions. In general ions are formed from that lose or gain enough electrons to gain a full octet in their valence shell.

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Advanced Physical Science B

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  1. Advanced Physical Science B Chemistry in Review: The Ions Chemical Formulas Chemical Equations Stoichiometry

  2. The Ions

  3. Ions • In general ions are formed from that lose or gain enough electrons to gain a full octet in their valence shell. • Elements that lose electrons become a positive CATION. • Elements that gain electrons become a negative ANION.

  4. Oxidation Numbers

  5. Monatomic Cations • Monatomic- one type of atom. • Most metals make monatomic cations, with a positive charge. • Usually the group number indicates the oxidation number of the elements in that group. • The cation simply has the same name as the element. • Transition Metals have multiple oxidation numbers.

  6. Cu+1, Cu+2 Fe+2, Fe+3 Pb+2, Pb+4 Sn+2, Sn+4 Hg2+2, Hg+2 Cuprous, cupric Ferrous, ferric Plumbous, plumbic Stannous, stannic Mercurous, mercurric High -ic and Low -ous! Elements with Multiple Charge

  7. Monatomic Anion • Monatomic- single type of atom. • Anions are usually made from Nonmetals, groups 15, 16, and 17. They gain electrons in their Valence. • All Anions end with a suffix. • Most monatomic anions end with a “-ide”.

  8. Polyatomic Cations • Polyatomic- more than one atom. • There just a few polyatomic cations. • NH4+, ammonium • Hg2+2, mercury(I)

  9. Polyatomic Anions • Polyatomic anions have more than one atom. • A nonmetal plus oxygen or oxygen and hydrogen. • Sometimes called an “oxyanion.” • Anions end with a suffix. • Most end with “-ate” • Polyatomic anions with less oxygens end with “-ite” • “ite” anions usually have one less oxygen then “ate” anions. • “ate” ate the “ite”!

  10. Chemical Formulas The building blocks.

  11. Symbols and Formulas • Names of Elements - 109 elements, >10 million known compounds • Compounds are represented by formulas combining chemical symbols and numeric subscripts. • Some elements are named for their properties. • Nitrogen-“niter forming” • Plumbic (lead)-shiny • Some elements are named for their place of origin.

  12. Symbols and Formulas (cont.2) • Some elements are named for the minerals they are found in. • Tungsten-Swedish name for “heavy stone” • Some elements are named in honor of a person. • Symbols for the elements • One or two letters, the first letter is capitalized • In 1813, JJ Berzelius, a Swedish chemist developed the modern symbols for designated elements.

  13. Chemical Formulas • Are a combination of symbols that represent the composition of a compound. • Molecular Compounds and Ionic Compounds.

  14. Ionic Compounds • Are compounds composed of charged particles. • In general: the electrons are shared between the ions. Metals tend to give up their electrons to an incomplete nonmetal. • All Ionic compounds are represented by their empirical formulas. They are always in the smallest whole number ratios.

  15. Other Types of Molecules • Diatomic Molecules:these 7 elements must exist in nature paired with itself or other elements. • Br2, Bromine - O2, Oxygen • I2, Iodine - H2, Hydrogen • Cl2, Chlorine - N2, Nitrogen • F2, Fluorine • “BrIClFOHN”

  16. Other Types of Molecules (cont.2) • Hydrates:Ionic Molecules attached to water molecules. • Organic Molecules:contains carbon as it’s central element. • Alloys: metals form these molecules where atoms are held together by a “sea” of electrons.

  17. Predicting Formulas of Ionic Compounds • Write the symbols for the elements in the compound • Always write the CATION first. • Determine the charge on each ion. • Na+1=+1, O-2=-2 • From the charge on each ion use subscripts to indicate the multiplier for the ions. • The total positive must equal the total negative. • The “total” charge of the compound must be zero. • Ex. Na2O

  18. Predicting Formulas of Ionic Compounds (cont.2) • When using subscripts for polyatomic ions, the ion is placed in parentheses, and the subscript is placed on the outside to indicate “x” ion units. • The subscript applies to all the elements in the parentheses. • If the subscript is “1”, it is understood and not written. • For monatomic ions no parentheses is used.

  19. Naming Ionic Compounds • Naming Binary Ionic Compounds: • The cation is listed first, then the monatomic anion. • For stock names include the oxidation number of the cation in parantheses. • For traditional names use the “-ous” or “-ic” name for the cation.

  20. Naming Ionic Compounds (cont.2) • Naming Ternary Ionic Compounds: • Made up with a cation and a polyatomic anion. • The suffix tells which anion. • “-ate” for more oxygen's • “-ite” for less oxygen's.

  21. Chemical Equations A chemical recipe

  22. Types of Chemical Reactions • There are 5 fundamental types of Chemical Reactions. • Synthesis (Direct Combination) • Decomposition (Analysis) • Single Replacement • Double Replacement • Combustion

  23. Synthesis(Direct Combination) • “joining together” • The general form of reaction: • A + B AB • element+element compound • Two reactants One product

  24. Synthesis (cont.2) O2 + 2NO2NO2

  25. Decomposition (Analysis) • “breaking down” • The general form of reaction: • AB A + B • compound element + element • One reactant Two products

  26. Decomposition (cont.2) 2NI3 N2 + 3I2

  27. Single Replacement • “Like ions must displace like ions” • The general form of reaction: • A + BC AC + B • element + compound compound + element • Two reactants Two products

  28. Single Replacement (cont.2) Fe2O3 + 2Al Al2O3 + 2Fe

  29. Double Replacement • “Exchanging ions” • The general form of reaction: • AC + BD AD + BC • compound+compound compound+compound • Two reactants Two products

  30. Double Replacement (cont.2) AgNO3 + NaCl AgCl + NaNO3

  31. Combustion • Special form of a decomposition rxn. • Burning hydrocarbons. • Metabolism • The general form of reaction: • hydrocarbon + oxygen CO2 + H2O • Presence of oxygen in the form, O2 • Products are always CO2 and H2O

  32. Combustion (cont.2) 2C8H18 + 17O2 18H2O + 8CO2

  33. Special Considerations for Replacement Reactions • Single Replacement Reactions: follow the “Activity Series” of elements. • Cations displace cations. • Anions displace anions. • Li+1 is the most reactive cation. • F-1 is the most reactive anion. • Double Replacement Reactions: must show evidence of a chemical reaction. • “God Punishes Chemistry Teachers” • Gas, Precipitate, Color change, Temperature change.

  34. Atom Accounting • Reactants-a starting substance in a chemical reaction. • Products-a substance produced in a chemical reaction. • Atoms in the reactants must equal the atoms in the products.

  35. Balancing Chemical Equations • Do an “Atom Accounting” • H2 + N2 NH3 H=2 H=3 N=2 N=1 • Li + Al2(SO4)3 Li2SO4 + Al Li=1 Li=2 Al=2 Al=1 S=3 S=1 O=12 O=4

  36. Balancing Basics • Rules for Balancing Chemical Equations: • Law of Conservations of Matter: “What goes IN must come OUT” • Be sure the elements in the products are in the reactants. • Make sure COMPOUNDS are good chemical formulas. • Use subscripts to make formulas.

  37. Balancing Basics (cont.2) • Balance the atoms on each side of the equation using COEFFICIENTS. • Do NOT Touch the Subscripts! • Keep the coefficients in the lowest whole numbered ratios. • Ex. 4H2 + 2O2 2H2O • Will be: • 2H2 + O2 H2O

  38. Balancing Basics (cont.3) • Balance the equation: • 3H2 + N22NH3 H=2 6 H=3 6 N=2 N=1 2

  39. Balancing Basics (cont.4) • Li + Al2(SO4)3 Li2SO4 + Al Li=1 Li=2 Al=2 Al=1 S=3 S=1 O=12 O=4 • 6Li + Al2(SO4)33Li2SO4 + 2Al Li=1 6 Li=2 6 Al=2 Al=1 2 S=3 S=1 3 O=12 O=4 12

  40. Stoichiometry Mathematical with chemical equations.

  41. Stoichiometry • A Chemical Equation gives information about the relative relationship (ratio) between reactants and products in a chemical reaction. • Coefficients of a balanced chemical equation gives three pieces of quantitative information about the reactants and the products. • The relative number of particles. • The relative number of moles. • The relative volume of a gas, at the same temperature and pressure.

  42. Stoichiometry (cont.2) • When a chemical equation is balanced, the total mass of the reactants equals the total mass of the products. (Law of Conservation of Mass) • The coefficients DO NOT give relative ratios of reactants to products by mass. • Must convert to MOLE or particles the compare coefficients.

  43. Stoichiometry (cont.3) • Organization is critical. • Balance the chemical equation, FIRST! • Determine the element/compounds that is given and the element/compound that is sought. Make a chart. • Place the information given in the problem under the correct element/compound.

  44. Extended Mole Map

  45. Mixed Stoichiometric Relationship • In general, this relationship holds: • All mixed relationship problems take 3 steps. • First, always balance the chemical equation and organize the problem. Determine what is Given and what if Sought. • Convert Given to moles, Change Given to Sought, Convert from Sought moles to whatever units asked for. • Remember: • If changing to/from mass: 1mol=Molar Mass • If changing to/from particles: 1mol=6.02x1023parts • If changing to/from volume(gases only): 1mol=22.4dm3 at STP

  46. Mixed Stoichiometric Relationship (cont.2) • MOLES RULE!!! • In One DA Table: Xunit,Sought = Given,units CF1 CF2 CF3 • Where: • CF1=converts units Given to moles • CF2=converts moles Given to moles Sought. The Mole Bridge. • CF3=converts moles Sought to units Sought.

  47. Try this mass-mass problem: • Calculate the mass of oxygen produced if 2.50g of potassium chlorate is completely decomposed to give potassium chloride and oxygen. • Balance the Chemical Equation: • 2KClO3 2KCl + 3O2 • Determine the Given and the Sought: • Given: 2.50g KClO3 • Sought: mass of O2 produced • Organize the appropriate information: • 2.50g KClO3 Xg O2

  48. In One Step • Calculate the mass of oxygen produced if 2.50g of potassium chlorate is completely decomposed to give potassium chloride and oxygen.. • The balanced chemical equation: • 2KClO3 2KCl + 3O2 • XgSought =Givenmass 1molGiven Mole MolarMassSought MolarMassGiven Bridge1mol Sought • XgO2 = 2.50gKClO3 1molKClO3 3O2 32gO2 123gKClO3 2KClO3 1molO2 = 9.76x10-1gO2

  49. #2 Try this: • How much silver phosphate is produced if 10.0g of silver acetate reacts with sodium phosphate? • Balance the Chemical Equation. • Organize the problem. • Use Three or One Step to solve the problem.

  50. How much silver phosphate is produced if 10.0g of silver acetate reacts with sodium phosphate? 3AgC2H3O2+ Na3PO4 Ag3PO4+ 3 NaC2H3O2 10.0gAgC2H3O2 XgAg3PO4 XgAg3PO4=10.0gAgAce 1AgAce 1Ag3PO4 418.58gAg3PO4 166.92gAgAce 3AgAce 1Ag3PO4 X= 8.36gAg3PO4

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