Chapter 2 - Atoms, Molecules and Ions

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Chapter 2 - Atoms, Molecules and Ions

• Graded HW 2 due before 10/12, Weds pm.
• End-of-Chapter Problems see next slide.
• Exam 1 (Chap. 12 Labs 12) on Friday, 10/14

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End-of-Chapter HomeworkChapter 2 pages 76 -
85(Not Turned In)

• 10 to 16, 19, 21, 23, 24, 27, 28, 35, 43, 45,
  47,
• 51, 57, 59, 61, 62, 65, 67, 69, 74 to 86,
• 91 to 97, 100, 101, 107, 109, 113, 115, 117,
  119, 121 to 124

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I. Atoms Molecules A. Atom Examples

• - Matter around us is made up of tiny particles
  called atoms. There are 117 known atoms and
  these are given in the Periodic Chart.
• - The atoms have names and they have one or two
  letter symbols.
• - The following is a Game to help learn a few
  symbols and placement on the Periodic Table - the
  answers to this quiz are the names of the correct
  elements.
• Half a dime Lone Rangers Horse
• What I do when I am hungry A frivolous prisoner
• Storage place for streetcars

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I. Atoms Molecules A. Atom Examples
Continued

• - A male member of the Ganese tribe.
  (Transition Metal)
• - What to do with an ailing man two answers.
  (Columns 182)
• - Comment made at end of box of candy oops,
  they .. (Col 18)
• - What some science classes do, but not this one.
  (Col 13, Row 2)
• - Atoms can exist as
• 1) Neutral elements He Na O Hg
• 2) Negative ions (anions) F-1 or F- I-
  O-2
• 3) Positive ions (cations) Na1 or Na H
  Al3
• - www.webelements.com good site for information
  (click on atoms)

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I. Atoms Molecules B. Molecules

• - Atoms combine together in definite ratios to
  form new units called Molecules (Molecular
  Compounds) Formula Units (Ionic Compounds) -
  We represent Molecules Formula Units with the
  symbols of the elements (left to right from
  Periodic Table) and with sub numbers to indicate
  the number of atoms other than one.
• - General Example ZxTy Where equals the
  number of molecules and x y equal the
  number of atoms in one molecule. Generally place
  element closest to group I first except for
  organic molecules with C first, H second, rest
  in alphabetical order. If a compound has more
  than 1 polyatomic ion, then use parenthesis.
• - Few Examples NaF CO2 H2O
  C2H2Br2
• Ca3(PO4)2 3 H2SO4
  C256H381N65O79S6 (Insulin)

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I. Atoms Molecules C. Size

• H 10-13 cm diam. Ho 10-8 cm diam.
  Mass H 1.67x10-24 g
• Note 6.02x1023 Avogadros Number
  (6.02x1023)x(1.67x10-24g)1.00 g
• - The following are images of 1)
  Electron Microscope
• 2) Latex Molecules 3) C in Graphene 4)
  U atoms 5) H C

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Atoms Molecules D. History

• 400 BC - Democritus suggested the existence of
  atoms.
• 66 AD - Peter wrote but the day of the Lord
  will come like a thief, in which the heavens will
  pass away with a roar and the elements will be
  destroyed with intense heat, and the earth and
  its works will be burned up. From New
  Testament, 2 Peter 310.
• 1783 - Antoine Lavoisier found that matter is
  not created nor destroyed in a chemical reaction.
  Known as father of modern chemistry.

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Atoms Molecules D. History
1803 - John Dalton proposed that matter is made
up of tiny atoms that atoms of the same element
are alike that atoms combine in definite
ratios to form compounds. This set aside false
idea promoted by Aristotle 2000 years earlier
that matter was continuous, and reaffirmed
Democrituss early atomic model.
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I. Atoms Molecules D. History

• 1879 - William Crookes developed the ray
  tube which later allowed us to view electron
  beams

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I. Atoms Molecules D. History

• 1897 Joseph Thomson used the cathode-ray tube
  and discovered the electron.

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I. Atoms Molecules D. History

• 1886 - Eugene Goldstein demonstrated existence
  of particles, protons. These particles later
  found to have a charge of 1 (1.60x10-19
  coulombs) and a mass of 1.67x10-24 g (a mass of
  1.00 AMU).
• 1909 - Robert Millikan determined mass
  (9.11x10-28 g 1800 less than proton) and -1
  charge (-1.60x10-19 coulombs) of an electron.

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I. Atoms Molecules D. History

• 1911 - Ernest Rutherford (a New Zealand
  physicist) demonstrated the nuclear nature of the
  atom in which the empty space is 10,000 to
  100,000 times larger than the size of the
  nucleus.

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I. Atoms Molecules D. History

• 1932 - James Chadwick demonstrated the
  existence of the neutron which has no charge and
  about the same mass as the proton (1.00 AMU).
• - Why do you think that it took longer to
  uncover the neutron than either the electron or
  proton?

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I. Atoms Molecules Summary

• Atoms are made up of three major parts
• Part Found Mass Charge
• Electrons Outside Nucleus 9.1x10-28g
  (small) -1
• Protons Part of Nucleus 1.7x10-24g (1.0
  AMU) 1
• Neutrons Part of Nucleus 1.7x10-24g (1.0 AMU) 0
• Notes 1) Neutral atoms contain equal of
  electrons and protons.
• 2) Atoms can loose or gain electrons to
  become charged ions
• 3) protons determines the identity of
  the atom or ion.
• 4) Mass of atom Protons Neutrons

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I. Atoms Molecules E. Atomic Structure

• - Heavier protons and neutrons in the center or
  nucleus and the smaller electrons are found
  outside of the nucleus in shells of specific
  energy.
• - The outer electrons and their arrangements in
  specific energy levels are responsible for the
  chemistry of the element.
• - The number of protons (atomic number)
  determines the identity of the atom NOT the
  number of neutrons NOT the number of electrons
  and NOT the atomic weight.
• - The number of protons plus neutrons determines
  the weight of the atom in AMU. Atomic Mass n
  p (from periodic chart). Why are the
  masses not integers?
• - The ratio of electrons to protons determines
  the charge of the atom atoms can loose or gain
  electrons to have or - charges for neutral
  atoms the e p.

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I. Atoms Molecules E. Atomic Structure
Continued

• - Atoms of a given element with differing numbers
  of neutrons are called isotopes. For example (
  note Mass Atomic H )
• 11H 1P 0N 21H 1P 1N
  (deuterium) 31H 1P 2N (tritium)
• - Note that the weights of the atoms are averages
  of the masses of the naturally occurring
  isotopes.
• - For example 52.0 of Br has 44 neutrons
  mass 79.0 7935 Br
• 48.0 of Br has 46
  neutrons mass 81.0 8135 Br
• Average mass (0.520 x 79.0) (0.480 x 81.0)
  79.9
• - Note that the masses of the atoms and their
  isotopes are determined with a useful instrument
  called a mass spectrometer ( MS ) .

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I. Atoms Molecules F. Mass Spectrometer
Used as a GC LC Detector for qualitative
quantitative analysis of atoms and compounds.
See Pgs 98 99 for information a mass spectrum
of CH2Cl2
Sample Introduced to the MS frequently by GC or LC
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I. Atoms Molecules F. Mass Spectrometer
Continued

• Result mass spectrum a plot of intensity
  versus m/e m.
• Qual. from masses and Quant. from intensities of
  peaks.
• Example MS of a mixture of NaCl and NaBr plot
  of intensity (y axis) vs mass (x axis).
• 23Na 35Cl 37Cl
  79Br
  80Br

Intensity
Mass
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Poisoned Pop ICP-MS Results
Element ppb Element ppb Element ppb Li 1.8 Be
2.3 B 3.0 Na gt1000 Mg gt1000 Al 10 Si gt1000 P
28 S gt1000 Cl 70 K gt1000 Ca gt1000 Sc 8.2 46Ti
35 48Ti 73 V 32 Cr 21 Mg 33 Fe gt1000 Co 21 N
i 38 63Cu 13 65Cu 17 66Zn 37 68Zn 40 Ga 7.6 Ge
2.6 As 10 Se 13 Br 18 Se 8.9 Rb 1.6 Sr 200 Zr
0.6 Nb 0.9 Mo 120 Ru 1.3 Rh 0.4 Pd 1.6 Ag 13
Cd 27 In 0.7 Sn 0.6 Sb 1.1 I 0.5 Te 0.4 Ba gt
1000 La 0.4 Ce 0.2 Pr 0.5 Nd 1.7 Sm 0.8 Eu 0.3
Gd 0.9 Tb 0.3 Dy 0.6 Ho 0.1 Er 0.6 Tm 0.4 Yb
0.7 Lu 0.2 Hf 1.0 Re 0.3 Os 0.4 Ir 0.5 Pt 0.6
Au 0.2 Hg 3.1 Tl 6.5 Pb 22 Bi 8.0 Ur 0.4
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II. Periodic Table A. Introduction

• 1869 - Dmitri Mendeleev (Russian chemist)
  Julius Lothar Meyer (German chemist)
  independently arranged the known atoms by atomic
  weight (now by atomic number).

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II. Periodic Table A. Introduction

• - Dmitri Mendeleev Lothar Meyer independently
  arranged the known atoms by atomic weight in
  1869.
• - In early 1900s we rearranged the Periodic
  Chart by atomic number and found that numerous
  physical and chemical properties had a regular
  repetition.
• - Rows are called periods are 7 of these named
  1 to 7 properties change going across a row.
• - Columns are called groups are three names
  used (IA--VIIIA, 1--18, and common names)
  all elements in a group have similar chemical
  properties.
• - Following are some periodic properties.

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II. Periodic Table B. Periodic Properties 1.
Metals Nonmetals

• - Metals on Left Side Note dividing line (H
  nonmetal) Good conductors of heat electricity
  Poor insulators Lustrous Solids Malleable
  Lose Electrons to form Cations like Na Mg2
  Fe3 Al3
• - Nonmetals on right side Poor conductors of
  heat electricity Not lustrous Brittle if
  solid Many are gases Can gain electrons to form
  anions, can also share electrons to yield
  neutral molecules
• F- O-2 F2 F F
  H2 H H O2 O O
• B C Nonmetals
• Al Si
• Metals Ga Ge As H is a
  non-metal
• In Sn Sb Te
• Tl Pb Bi Po At

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II. Periodic Table B. Periodic Properties 2.
Size

• - Neutral atoms get larger as go down periodic
  chart.
• - Neutral atoms get larger as go left on
  periodic chart.

Fr
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II. Periodic Table B. Periodic Properties

• 3. Ionization Energy
• - Energy needed to remove outer electron of a
  neutral atom.
• - Generally increases as go up and as go to right
  on PC.
• 4. Electronegativity (Ignore group 18 helium
  group)
• - A measure of the desire of atom for electrons.
• - Gets larger as go up and as go right on the
  periodic chart F is most and Fr is least
  electronegative. Important

F
Fr
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II. Periodic Table B. Periodic Properties

• 5. Acidity/Basicity (Will define later)
• - Elements on left form bases like NaOH, Ba(OH)2
• - Elements on right form acids like HCl, HNO3,
  H2SO4
• 6. Charges Some atoms lose or gain electrons to
  mimic nearest group VIIIA/18 element. Know
  these charges.
• IA/1 IIA/2 IIIA/13 VA/15
  VIA/16 VIIA/17 VIIIA/18
• 1 2 3 -3
  -2 -1 0

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II. Periodic Table B. Periodic Properties

• Note Under normal conditions, elements in Black
  are solids (K, S, U), elements in Blue are
  liquids (Hg Br2), elements in Red are gases
  (H2, F2, Ar), and outlined elements are man-made
  (Tc, Pu).
• IMPORTANT Memorize Hydrogen, Nitrogen, Oxygen,
  Fluorine, Chlorine, Bromine Iodine normally
  exist as diatomic molecules, and have the same
  name as the element
• H2 O2 F2 Cl2
  Br2 I2

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III. Compounds A. Introduction

• - Atoms combine in definite ratios to form
  compounds.
• - Are two classes of compounds
• 1) Ionic Compounds (Compound containing a
  METAL)
• Metals lose electrons in chemical reactions to
  form ions called cations. This will occur when
  the metal reacts with a nonmetal. Nonmetals gain
  electrons to become - negative ions called
  anions. (Note H classified as a nonmetal)
• - The cations and anions attract each other in
  such a ratio so that the compound is neutral.
• - Ionic compounds form aggregates and we give
  the simplest ratio of the atoms, called the
  formula unit.
• - Examples Na F- -----) NaF (Dont
  write as NaF-)
• K K O2- -----) K2O

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III. Compounds A. Introduction

• Predicting formulas for ionic compounds
• - Need to memorize general charges for the
  groups
• IA/1 IIA/2 IIIA/13 VA/15
  VIA/16 VIIA/17
• 1 2 3 -3
  -2 -1
• - Some metals have variable charges
• Cu1 Cu2 Fe2 Fe3 Cr3 Cr6
  Hg2 Hg22
• Pb2 Pb4 Sn2 Sn4 Need to know
  these.
• - Add the anions and cations in the simplest
  ratio to get a neutral compound. Do not show
  charges in formula.
• - Examples K and F Be and O Ba and
  Cl
• Fe3 and S Mg and N Al and F
• (KF BeO BaCl2 Fe2S3
  Mg3N2 AlF3)

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III. Compounds A. Introduction

• 2) Molecular (Compound with TWO NONMETALS)
• - H and other nonmetals can chemically bind with
  other nonmetals by sharing electrons to produce
  new units called molecules, and the molecular
  formula gives the exact number of atoms in the
  molecule.
• - Ionic compounds existed as aggregates that we
  show as formula units however, nonmetals combine
  to produce distinct molecules. Molecules are
  represented with molecular formulas or with
  structural formulas. Structural formulas show
  atoms attached with covalent bonds each bond is
  a line and represents two shared electrons.
• - Examples
• H2O2 or H-O-O-H CO2 or OCO

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III. Compounds A. Introduction

• Organic Compounds
• - One class of molecular compounds is organic
  compounds.
• - Carbon containing compounds plus H and
  frequently O, N Cl are called organic
  compounds. Write C, then H, then rest of
  elements in alphabetical order C3H9NO
• - The vast majority of known chemicals are
  organic there are an infinite number of possible
  organic compounds.
• - Organic compounds are organized by functional
  group (that portion of the organic compound that
  governs its chemistry).
• Examples
• Ether R O R such as
  CH3-CH2-O-CH2-CH3 (C4H10O)
• Alcohol R O H such as CH3-OH (CH4O)
• Amine R NH2 such as NH2-CH2-CH2-CH2-CH2-NH2 (C
  4H12N2)

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III. Compounds B. Examples

• - Classify the following as either ionic or
  molecular
• NaF CO2 BaF2 H2O SF2 H2
  FeI2 CH4
• - Predict the formula when the following form
  ionic substances
• Na I Ba I K O
  Ca S
• Be F Be Se Na N
  Mg N
• Fe3 I Fe2 I Fe3 O Fe2
  O

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III. Compounds C. Polyatomic Ions

• We frequently encounter ions in which several
  atoms have combined together to form a polyatomic
  ion.
• Polyatomic Ion An ion consisting of several
  atoms bonded together and carrying a charge.
• Some of these polyatomic ions occur so frequently
  that we need to know the names, formulas and
  charges on the ions.
• When have more than one polyatomic ion, then use
  ( ).
• Examples
• NaOH OH-1 Hydroxide
• KNO2 NO2-1 Nitrite
• Ca(NO3)2 NO3-1 Nitrate

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III. Compounds C. Polyatomic Ions

• Formula Name Formula Name
• C2H3O2-1 Acetate NH4 Ammonium
• CO3-2 Carbonate CN- Cyanide
• CrO4-2 Chromate Cr2O7-2 Dichromate
• OH- Hydroxide Hg22 Mercury (I)
• NO3- Nitrate NO2- Nitrite
• MnO4- Permanganate O2-2 Peroxide
• PO4-3 Phosphate SeO3-2 Selenite
• SO3-2 Sulfite SO4-2 Sulfate
• ClO4- Perchlorate ClO3- Chlorate
• ClO2- Chlorite ClO- Hypochlorite
• HCO3- Hydrogen Carbonate HSO4-1 Hydrogen
  Sulfate
• Note Memorize the ones in Blue for exam 1
  rest know before exam 2. Need to memorize names,
  formulas and charges.

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IV. Nomenclature A. Introduction

• Use rule 1 for Ionic Compounds Use rule 2 for
  Molecular Compounds.
• Nomenclature Name more element then name more
  element change the ending to ide (for
  binary ionic compounds)
• Examples of the ide ending
• Atom Anion Name
• Chlorine Cl1- Chloride
• Oxygen O2- Oxide
• Fluorine F1- Fluoride
• Sulfur S2- Sulfide
• Nitrogen N3- Nitride
• Iodine I1- Iodide
• Bromine Br1- Bromide
• Phosphorus P3- Phosphide

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IV. Nomenclature B. Ionic Compounds, Rule
1

• - Rule 1 for ionic compounds Name the
  element, then the element change the ending
  to ide.
• Exceptions a) If have a polyatomic ion, then
  name it.
• b) If have a variable charged metal, then
  give its charge in the middle with a Roman
  Numeral in parenthesis.
• c) two, three,... x polyatomic ions use (
  )x
• - Examples
• NaCl sodium chloride FeF3 iron (III)
  fluoride
• Li2O lithium oxide LiNO2 lithium
  nitrite
• BaI2 barium iodide Cu(CN)2 copper (II)
  cyanide
• AlF3 aluminum fluoride Ca(NO3)2 calcium
  nitrate
• PbSO4 lead(II)sulfate NH4I ammonium
  iodide

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IV. Nomenclature B. Ionic Compounds, Rule
1

• NaCl Sodium Chloride
• BaI2 Barium Iodide
• Al2O3 Aluminum Oxide
• Ca(NO2)2 Calcium Nitrite
• Cu2S Copper(I) Sulfide
• Sodium Hydroxide NaOH
• Ammonium Fluoride NH4F
• Iron(III)Sulfide Fe2S3
• Calcium Phosphate Ca3(PO4)2
• Beryllium Oxide BeO

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IV. Nomenclature C. Molecular Compounds,
Rule 2

• When nonmetals combine with each other through
  sharing electrons (covalent bonds), they form
  molecules - there are no ions (H is included).
• Rule 2 When both elements are nonmetals
  (molecular compounds), then Name the the -
  change ending to ide and use prefixes of di,
  tri, tetra, penta, hexa, hepta, octa, nona, deca
  ( 2, 3, 4, 5, 6, 7, 8, 9 10 )

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IV. Nomenclature C. Molecular Compounds,
Rule 2

• CO2 Carbon Dioxide
• CCl4 Carbon Tetrachloride
• N2O Dinitrogen Oxide
• P2S5 Diphosphorus Pentasulfide
• PBr3 Phosporus Tribromide
• BI3 Boron Triiodide
• Notes (1) Organic compounds like C2H6O use
  their own rules which we wont cover in this
  class.
• (2) diatomic molecules named with the
  element name. Example O2 Oxygen

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IV. Nomenclature D. Examples

• Frequently have the chemical name and need to
  derive the formula. Need to know charges if
  ionic and name charge of polyatomic ions.
• Carbon Dioxide CO2
• Barium Oxide BaO
• Calcium Hydroxide Ca(OH)2
• Sulfur Trioxide SO3
• Hydrogen Cyanide HCN
• Iron(III) Fluoride FeF3
• Barium Nitrite Ba(NO2)2

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IV. Nomenclature D. Examples Continued

• NaF
• CS2
• NI3
• BaI2
• K3PO4
• Iron(II) Oxide
• Sodium Sulfate

• Sodium Fluoride
• Carbon Disulfide
• Nitrogen Triiodide
• Barium Iodide
• Potassium Phosphate
• FeO
• Na2SO4

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IV. Nomenclature E. Acids

• - Acids are an important class of compounds that
  have their own set of names. Know before exam 2.
• H2SO4 Sulfuric Acid
• HNO3 Nitric Acid
• H3PO4 Phosphoric Acid
• HC2H3O2 Acetic Acid
• HClO4 Perchloric Acid
• HClO3 Chloric Acid
• HClO2 Chlorous Acid
• HClO Hypochlorous Acid
• HCl Hydrochloric Acid
• HBr Hydrobromic Acid
• HF Hydrofluoric Acid
• HI Hydroiodic Acid

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V. Nomenclature Summary

• Have to know - Metals/Nonmetals
  Ionic/Molecular Compds
• - Charges on ions
• - Names, charges, formulas of polyatomic
  ions
• - Prefixes di, tri, tetra, penta, hexa
• - Rules for nomenclature
• Name then - change ending to ide
• Exceptions
• name polyatomic ion
• give charge on multicharged ion (RN)
• use di, tri. for two nonmetals
• - Be able to give formula from Name

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VI. Chemical Equations A. Introduction

• - A chemical equation is the representation of
  the rxn in terms of chemical formulas.
• - Example 2 Mg (s) 1 O2 (g) -----) 2 MgO
  (s)
• Notes - 1 may not be shown assume 1 if no
  given.
• - Mg O2 are reactants.
• - MgO is a product.
• - ------) or
  goes to give or yields.
• - 2, 1, 2 are balancing coefficients.
• - May add ?H at the end for heat lost or
  gained.
• - The subscripts of (g) (s) (l) (aq) indicate
  the physical state of the participants.

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VI. Chemical Equations A. Introduction
Continued

• Notes Continued
• - The driving force in a chemical reaction is to
  produce products which are more stable than the
  reactants.
• - How tell if a chemical reaction has taken
  place?
• 1) Gas produced
• 2) New solid, ppt, produced (ppt
  precipitate)
• 3) Light produced
• 4) Heat may be lost or gained (temperature
  change)
• - Can predict products with a) experience, with
  b) tables of solubility c) with eqns of rxns
  that produce gasses.
• - Have to balance by inspection factor.

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VI. Chemical Equations B. Balancing

• - Mass is neither lost or gained in a chemical
  reaction so, we need to balance the equation
  without changing the identity of the reactants or
  products - change only balancing coefficients.
• - Balancing is done by inspection, but for
  difficult equations it is best to balance the
  element which occurs the fewest times - first.
• - Examples
• ___Al ___Cl2 -----) ___AlCl3
• ___Ca ___H2O -----) ___Ca(OH)2
  ___H2
• ___HCl ___Al(OH)3 -----) ___AlCl3
  ___H2O
• ___Ba(OH)2 ___H3PO4 -----)
  ___Ba3(PO4)2 ___H2O