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Chemistry in Review: The Ions Chemical Formulas Chemical Equations Stoichiometry

Chemistry in Review: The Ions Chemical Formulas Chemical Equations Stoichiometry. The Ions. Ions. Ions are formed when elements lose or gain enough electrons to gain a full octet in their valence shell. Elements that lose electrons become a positive CATION.

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Chemistry in Review: The Ions Chemical Formulas Chemical Equations Stoichiometry

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  1. Chemistry in Review: The Ions Chemical Formulas Chemical Equations Stoichiometry

  2. The Ions

  3. Ions • Ions are formed when elements 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 (Positive) • Monatomic- one type of atom. • Most metals make monatomic cations, with a + charge • Usually the group number indicates the oxidation number of the elements in that group. • The cation 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 Copper (I), Copper (II) (Cuprous, cupric) Iron (II), Iron (III) (Ferrous, ferric) Lead (II), Lead (IV) (Plumbous, plumbic) Tin (II), Tin (IV) (Stannous, stannic) Mercury (I), Mercury (II) (Mercurous, mercurric) High -ic and Low -ous! Elements with Multiple Charges

  7. Monatomic Anion (negative) • 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. • Few polyatomic cations. • NH4+, ammonium • Hg2+2, mercury(I) • Hg+2, mercury(II) • H3O+, hydronium

  9. Polyatomic Anions • Polyatomic anions have more than one atom. • Anions end with a suffix. • Most end with “-ate” • Polyatomic anions with less oxygens end with “ite” • “ite” anions = 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: 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. • Germanium - Germany • Europium- Europe

  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. • Alfred NOBEL-ium • 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 • Compounds composed of charged particles. • 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. • H2, Hydrogen • N2, Nitrogen • O2, Oxygen • F2, Fluorine • Cl2, Chlorine • Br2, Bromine • I2, Iodine • “Horses need oxygen for clear bright eyes ( I’s)” • Mr. HOFBrINCl

  16. Other Types of Molecules (cont.2) • Hydrates:Ionic Molecules attached to water molecules. MgSO4. 9 H2O • Organic Molecules:contains carbon as it’s central element. C6H12O6 • Alloys: metals form these molecules where atoms are held together by a “sea” of electrons.

  17. Predicting Formulas of Ionic Compounds • Write symbols for elements in the compound • Always write the CATION first. • Determine the charge on each ion. Na+1 = +1, O-2 = -2 • From 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. Al2S3 N2O3 Naming Ionic Compounds • Naming Binary Ionic Compounds: • The cation is listed first, then the monatomic anion. aluminum sulfide dinitrogen trioxide • For stock names include the oxidation number of the cation in parantheses.

  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. sulfate, SO4-2 sulfite, SO3-2 Aluminum sulfate

  21. Chemical Equations A chemical recipe

  22. Types of Chemical Reactions Fundamental types of Chemical Reactions. • Synthesis (Direct Combination) • Decomposition • 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 O2 + 2NO2NO2

  24. Decomposition (Analysis) • “breaking down” • The general form of reaction: • AB A + B • compound element + element • One reactant Two products 2NI3 N2 + 3I2

  25. 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 Fe2O3 + 2Al Al2O3 + 2Fe

  26. Double Replacement • “Exchanging ions” • The general form of reaction: • AC + BD AD + BC • compound+compound compound+compound • Two reactants Two products AgNO3 + NaCl AgCl + NaNO3

  27. 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 2C8H18 + 17O2 18H2O + 8CO2

  28. 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 Prefers Chemistry Teachers” • Gas, Precipitate, Color change, Temperature change.

  29. 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.

  30. 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

  31. 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.

  32. Balancing Basics (cont.2) • Balance the atoms on each side of the equation using COEFFICIENTS. • Do NOTTouch the Subscripts! • Keep the coefficients in the lowest whole numbered ratios. • Ex. 4H2 + 2O2 2H2O • Will be: • 2H2 + O2 H2O

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

  34. 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

  35. Stoichiometry Mathematics with chemical equations.

  36. Stoichiometry • A Chemical Equation gives information about the relative relationship (ratio) between reactants and products in a chemical reaction. • Coefficientsof 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.

  37. 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.

  38. 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.

  39. Extended Mole Map

  40. Mixed Stoichiometric Relationship • In general, this relationship holds: • All mixed relationships problems take 3 steps. • First, always balance the chemical equation and organize the problem. Determine what is Givenand 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=22.4 L at STP

  41. Mixed Stoichiometric Relationship(cont.2) • MOLES RULE!!! • In One DA Table: X unit, 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.

  42. 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

  43. 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 = Given mass 1molGivenMole MolarMassSought • MolarMassGiven Bridge1mol Sought Xg O2 = 2.50gKClO3 1molKClO3 3O2 32gO2 123gKClO32KClO3 1molO2 = 9.76x10-1gO2

  44. #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.

  45. How much silver phosphate is produced if 10.0g of silver acetate reacts with sodium phosphate? 3AgC2H3O2+ Na3PO4 Ag3PO4+ 3 NaC2H3O2 10.0g AgC2H3O2Xg Ag3PO4 Xg Ag3PO4= 10.0 g AgC2H3O2 1 AgC2H3O2 1 Ag3PO4 418.58gAg3PO4 166.92g AgC2H3O2 3Ag C2H3O2 1 Ag3PO4 X= 8.36gAg3PO4

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