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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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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.

  • Elements that lose electrons become a positive CATION.

  • Elements that gain electrons become a negative ANION.


Oxidation Numbers


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.


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


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”.


Polyatomic Cations

  • Polyatomic- more than one atom.

  • There just a few polyatomic cations.

    • NH4+, ammonium

    • Hg2+2, mercury(I)


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”!


Chemical Formulas

The building blocks.


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.


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.


Chemical Formulas

  • Are a combination of symbols that represent the composition of a compound.

  • Molecular Compounds and Ionic Compounds.


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.


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”


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.


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


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.


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.


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.


Chemical Equations

A chemical recipe


Types of Chemical Reactions

  • There are 5 fundamental types of Chemical Reactions.

    • Synthesis (Direct Combination)

    • Decomposition (Analysis)

    • Single Replacement

    • Double Replacement

    • Combustion


Synthesis(Direct Combination)

  • “joining together”

    • The general form of reaction:

      • A + B AB

      • element+element compound

      • Two reactants One product


Synthesis (cont.2)

O2 + 2NO2NO2


Decomposition (Analysis)

  • “breaking down”

    • The general form of reaction:

      • AB A + B

      • compound element + element

      • One reactant Two products


Decomposition (cont.2)

2NI3 N2 + 3I2


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


Single Replacement (cont.2)

Fe2O3 + 2Al Al2O3 + 2Fe


Double Replacement

  • “Exchanging ions”

    • The general form of reaction:

      • AC + BD AD + BC

      • compound+compound compound+compound

      • Two reactants Two products


Double Replacement (cont.2)

AgNO3 + NaCl AgCl + NaNO3


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


Combustion (cont.2)

2C8H18 + 17O2 18H2O + 8CO2


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.


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.


Balancing Chemical Equations

  • Do an “Atom Accounting”

    • H2 + N2 NH3

      H=2H=3

      N=2N=1

    • Li + Al2(SO4)3 Li2SO4 + Al

      Li=1Li=2

      Al=2Al=1

      S=3S=1

      O=12O=4


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.


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


Balancing Basics (cont.3)

  • Balance the equation:

    • 3H2 + N22NH3

      H=2 6H=3 6

      N=2N=1 2


Balancing Basics (cont.4)

  • Li + Al2(SO4)3 Li2SO4 + Al

    Li=1Li=2

    Al=2Al=1

    S=3S=1

    O=12O=4

  • 6Li + Al2(SO4)33Li2SO4 + 2Al

    Li=1 6Li=2 6

    Al=2Al=1 2

    S=3S=1 3

    O=12O=4 12


Stoichiometry

Mathematical with chemical equations.


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.


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.


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.


Extended Mole Map


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


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.


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


  • 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


    #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.


    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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