CHEM 120: Introduction to Inorganic Chemistry

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CHEM 120 Introduction to Inorganic Chemistry

• Instructor Upali Siriwardane (Ph.D., Ohio State
  University)
• CTH 311, Tele 257-4941, e-mail
  upali_at_chem.latech.edu
• Office hours 1000 to 1200 Tu Th 800-900
  and 1100-1200 M,W, F

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Chapters Covered and Test dates

• Tests will be given in regular class periods 
  from  930-1045 a.m. on the following days
• September 22,     2004 (Test 1) Chapters 1 2
• October 6,         2004(Test 2)  Chapters  3,
  4
• October 20,         2004 (Test 3) Chapter  5 6
  
• November 3,        2004 (Test 4) Chapter  7 8
• November 15,      2004 (Test 5) Chapter  9 10
• November 17,      2004 MAKE-UP Comprehensive
  test (Covers all chapters
• Grading
• ( Test 1 Test 2 Test3 Test4 Test5)
  x.70 Homework quiz average x 0.30 Final
  Average
•                               5

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Chapter 4 Structure and properties of ionic and
covalent compounds

• We now put atoms and ions together to form
  compounds

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Chapter 4. Structure and Properties of Ionic and
Covalent Compounds

• 1. Classify compounds as ionic, covalent, or
  polar covalent bonds.
• 2. Write the formulas of compounds when provided
  with the name of the compound.
• 3. Name common inorganic compounds using standard
  conventions and recognize the common names of
  frequently used substances.
• 4. Predict the differences in physical state,
  melting and boiling points, solid-state
  structure, and solution chemistry that result
  from differences in bonding.
• 5. Draw Lewis structures for covalent compounds
  and polyatomic ions.
• 6. Describe the relationship between stability
  and bond energy.
• 7. Predict the geometry of molecules and ions
  using the octet rule and Lewis structure.
• 8. Understand the role that molecular geometry
  plays in determining the solubility and melting
  and boiling points of compounds.
• 9. Use the principles of VSEPR theory and
  molecular geometry to predict relative melting
  points, boiling points, and solubilities of
  compounds.

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Start learning the formulas and the names and
charges of the ions found in table
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• Why have we been so interested in where the
  electrons are in an atom? And what is the
  importance of valence electrons?
• Valence es are involved in_______--the no of
  valence es has an important influence on ______
  of bonds formed. The filled inner core does not
  directly affect bond formation.

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Compound

• Bonds are formed by a transfer of ________ from
  one atom to another or by a ______ _________
  between 2 atoms.

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Lewis (dot) Symbols
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Lewis (dot) symbols

• Introduced by G. N. Lewis
• Useful for representative (sp block) elements
  only
• Group no. no of valence e-s (no of dots)

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Lewis symbols for A groups

• The elements symbol represents the inner core of
  electrons. Put a dot for each valence electron
  around the symbol.
• Remember that the no. of valence electrons for
  the A groups is equal to ?
• Each unpaired electron may be used in bond
  formation

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Remember the octet rule from chapter 3

• So the ions formed by the elements in
• IA
• IIA
• IIIA
• VA
• VIA
• VIIIA

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Ionic bonding

• Extra stability has been noted for the noble gas
  configuration (8 e-s in valence shell)--(for A
  elements)
• Ionic bonding
• Each atom in the ionic bond

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• Ionic compounds are formed between
• And
• When forming an ionic bond each atom in the bond
  attains a noble gas configuration by a complete
  transfer of

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• An ionic bond is the electrostatic force that
  holds ions together in an ionic compound
• An ionic bond is a very strong bond ionic cmpds
  have high m and b pts.

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Typical ionic reactions with Lewis structures
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What about Li and S?
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What about Ca and O

• Formula is

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What about Ca and N?

• Formula is

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Covalent bonding

• Not all bonds are ionic.
• ________ bonds are bonds in which two (or more)
  electrons are ______ by two atoms.
• One shared electron pair is

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• A reminder
• Only valence electrons are involved in bonding.
  Group No. valence e-s for A elements.
• Covalent bonds are formed
• Each atom in bond attains noble gas configuration
  by sharing of e- pairs (H2 bond only has 2 e-s)

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Covalent bond formation

• Look at formation of H2 molecule.
• H. .H ----gt HH (H-H)
• 1s1 1s1 bond formed by
  overlap of
  1s orbitals

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What about F2 or Cl2?
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____ _____ - pairs of valence electrons not
involved in covalent bond formation Lewis
structure - representation of covalent bonding in
which lone pairs are shown as pairs of dots and
bonding pairs are (usually) shown as lines
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Polar covalent bonding and electronegativity

• Not all covalent bonds are formed btn the same 2
  atoms (as H2, homonuclear diatomic
  _______sharing of e-s in bond)

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Polar covalent bonds

• What about the bond in H-F?
• It is known that F is more likely to attract e-s
  to itself than H, leading to an unequal sharing
  of the e- pair.
• The covalent bond in which there is unequal
  sharing

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Polar covalent bond or polar bond is a covalent
bond with greater electron density around one of
the two atoms
electron rich region
electron poor region
e- rich
e- poor
d
d-
9.5
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Continuum of bond polarity

• (Nearly) complete e- transfer ionic bond
• Unequal sharing of e- pair polar covalent bond.
  
• e-s are polarized toward Cl
• Equal sharing of e- pair nonpolar covalent bond

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Electronegativity

• Electronegativity
• .
• Eneg is a relative concept. Elements with

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Lanthanides 1.1-1,3 Actinides 1.3-1.5
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Electronegativity differences

• 0.2 - 0.5 will be a ________________ bond
• 0.5 - 1.6 will be a ________________ bond
• gt 1.6 will be a ________________ bond

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Electronegativity differences

• In general the _______ the difference in eneg btn
  the 2 atoms in the bond, the ____ ______ the
  bond.
• If the difference is zero, bond (equal
  sharing of electron pair(s) (H2, Cl2, O2, F2, N2)

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• If the difference is gt0 and lt1.9, have a
  HCl (3.0 - 2.1) HF (4.0-2.1) OH (3.5-2.1)
• If the difference is gt 1.9, have NaCl
  (3.0-0.9) CaO (3.5-1.0)

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Classify as ionic or covalent

• NaCl
• CO
• ICl
• H2
  

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• Which bond is the most polar (most ionic), which
  the least polar (most covalent)?
• Li-F Be-F B-F C-F N-F O-F F-F

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• Classify the following bonds as ionic,
  polar covalent, or covalent. A) the CC bond
  in H3CCH3
• B) the KI bond in KI
• C) the NB bond in H3NBCl3
• D) the CF bond in CF4

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Chemical formulas

• Express composition of molecules (smallest unit
  of covalent cmpds) and ionic compounds in
  chemical symbols
• H2O, NaCl

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Writing formulas for ionic cmpds

• Compounds are neutral overall. Therefore
• NaCl is array of Na and Cl- ions
• Na2S is array of Na and S2- ions

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Predict the formulas for the cmpd formed btn

• Potassium and chlorine
• Magnesium and bromine
• Magnesium and nitrogen

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Symbol Name Symbol Name
H Hydrogen ion H- Hydride ion
Li Lithium ion F- Fluoride ion
Na Sodium ion Cl- Chloride ion
K Potassium ion Br- Bromide ion
Be2 Beryllium ion I- Iodide ion
Mg2 Magnesium ion O2- Oxide ion
Ca2 calcium ion S2- Sulfide ion
Ba2 barium ion N3- Nitride ion
Zn2 zinc ion P3- Phosphide ion
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Formula Name Formula Name
NO3- nitrate CO32- carbonate
NO2- nitrite SO42- sulfate
CN- cyanide SO32- sulfite
MnO4- permanganate PO43- phosphate
OH- hydroxide PO33- phosphite
O22- peroxide ClO4- perchlorate
HCO3- hydrogen carbonate ClO3- chlorate
HSO4- hydrogen sulfate ClO2- chlorite
HSO3- hydrogen sulfite ClO- hypochlorite
HPO42- hydrogen phosphate CrO42- chromate
H2PO4- dihydrogen phosphate C2H3O- 2 acetate
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Symbol (Stock system) Common Symbol (Stock system) Common
Cu copper(I) cuprous Hg22 mercury(I) mercurous
Cu2 copper(II) cupric Hg2 mercury(II) mercuric
Fe2 iron(II) ferrous Pb2 lead(II) plumbous
Fe3 iron(III) ferric Pb4 lead(IV) plumbic
Sn2 tin(II) stannous Co2 cobalt(II) cobaltous
Sn4 tin(IV) stannic Co3 cobalt(III) cobaltic
Cr2 chromium(II) chromous Ni2 nickel(II) nickelous
Cr3 chromium(III) chromic Ni4 nickel(IV) nickelic
Mn2 manganese(II) manganous Au gold(I) aurous
Mn3 manganese(III) manganic Au3 gold(III) auric
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Polyatomic ions Table

• Just have to memorize
• NH4 ammonium ion
• CO32- carbonate ion
• CN- cyanide ion
• HCO3- hydrogen (or bi) carbonate ion
• OH- hydroxide

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• NO3- nitrate ion
• NO2- nitrite ion
• PO43- phosphate ion
• SO42- sulfate ion
• HSO4- hydrogen sulfate ion
• SO32- sulfite ion
• CH3COO- (C2H3O2-) acetate ion

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• These polyatomic ions also form ionic cmpds when
  they are reacted with a metal or a nonmetal in
  the case of the ammonium ion (or with each other
  as ammonium sulfate). These polyatomic species
  act as a

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• So the formula for the cmpd formed btn the
  ammonium ion and sulfur would be
• and between calcium and the phosphate ion

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• Ionic cmpds do not exist in discrete pairs of
  ions. Instead, in the solid state, they exist as
  a three dimensional array--crystal lattice --of
  cations and anions--are neutral overall,

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Given name, write formula

• potassium oxide
• magnesium acetate

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Naming ionic cmpds

• Name the cation and anion but drop the word ion
  from both. This includes the polyatomic ions.
• Na2S
• Ca3N2

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Name

• Na3PO4
• NH4Cl
• K2S

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Cations with more than one charge

• Cu copper(I) Cu2 copper(II)
• So Cu2O is and
• CuO is

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Given name, write formula

• Ammonium chloride
• potassium cyanide
• silver oxide
• Magnesium chloride
• Sodium sulfate
• Iron(II) chloride

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To name covalent cmpds

• Name the parts as for ionic cmpds (CO carbon and
  oxide) but tell how many of each kind of atom by
  use of Greek prefixies. (Table 4.4)
• The mono- (for 1) may be omitted for the first
  element

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• Prefix meaning
• Mono- 1
• Di- 2
• Tri- 3
• Tetra- 4
• Penta- 5
• Hexa- 6
• Hepta- 7
• Octa- 8
• Nona- 9
• Deca- 10

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• CO
• CO2
• P4S10
• Boron trichloride
• Water H2O Ammonia NH3

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Write formula

• Diboron trichloride
• Sulfur trioxide
• Potassium sulfide

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Covalent cmpds

• Remember covalent cmpds--
• A _________ is the smallest unit of a covalent
  cmpd that retains the characteristics of the
  cmpd. Molecule - two or more atoms in a definite
  arrangement held together by chemical bonds.
  (H2O, Cl2) Cl2 is considered a molecule but not
  a cmpd
• Molecular cmpds exist as

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Comparison of properties of ionic and covalent
cmpds

• Physical state
• Ionic cmpds are
• Molecular cmpds can be

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Comparison continued

• Melting (___________) and boiling
  (_________) pts
• In general the melting and boiling temps are much
  _______for ionic cmpds than for molecular
  (covalent) cmpds. The ionic bond is very strong
  and requires a lot of (heat) energy to break the
  bond. The bond btn molecular species is not as
  strong.

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Comparison continued

• Structure in solid state
• Ionic solids--
• Covalent solids--

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Comparison continued

• In aqueous (H2O) solution
• Ionic cmpds dissociate into the
• Many covalent cmpds when dissolved in water
  retain their structure and molecular identity

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• Learn the names, formulas, charges, etc for those
  ions highlighted in table 4.3.
• HCO3- you should learn as bicarbonate

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Writing Lewis structures for covalent species

• These rules are for covalently bonded cmpds only
  (btn 2 or more nonmetals)
• Do not use them for ionic cmpds.
• 1. Count the total no. of valence electrons (the
  group no. is equal to the no. of valence
  electrons).
• if the species is an anion, increase the no. of
  valence electrons by the charge on the ion

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• if the species is a cation, subtract the charge
  of the cation from the total no. of valence
  electrons.
• 2.Count the total no. of atoms, excluding H, in
  the molecule or ion. Multiply that no. by 8.
• Exception multiply the no. of Hs by 2.
• This tells you how many electrons you would need
  if you were putting 8 electrons around all atoms
  without any sharing of electrons (and 2 around
  all Hs).

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• 3. Subtract the no. of e-s calculated in step 1
  from the no. in step 2. This gives you the no. of
  e-s that must be shared to get an octet around
  all atoms in the molecule.
• 4. no. of e-s that must be shared /2 gives you
  the no. of bonds.
• 5. subtract the no. of e-s that are shared (from
  step 3) from the total no. of valence e-s. This
  gives you the no. of unshared e-s.
• If you divide the no. of unshared e-s by 2 you
  get the no. of lone pairs.

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• Write the skeletal structure and fill in with the
  info you came up with. After youve put in the
  bonds calculated, fill in the octets.
• H (and F) form only one bond. Therefore they can
  only be terminal atoms in a structure.
• So you can not have
• C---H---C
• It has to be H---C--C

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• Examples
• CH4
• PCl3
• SO32-
• NO3-
• CN-
• COBr2 (C is bonded to O and Br atoms)
• SO2
• H3O (hydronium ion
• N3-

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Draw Lewis structure of CO2 i) Valence
electrons 4 2 x 6 16 ( 8 pairs) ii)
Central atom C O -- C -- O iii)
Give octet to carbon
-- O -- C --
O -- Try to
fill octet to O iv) Count electrons 4 bond
pairs 4 pairs 4 lone pairs 4 pairs
8 electron pairs
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Multiple bonds

• In general a triple bond (N2) is ________ than a
  double bond (O2) which is ________than a single
  bond (F2).
• Bond order BO of 1--single bond, BO of 2--
  -double bond, BO of 3 --triple bond.
• The stronger the bond,

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Terminology used in describing Lewis structures
of molecules Bond pairs An electron pair shared
by two atoms in a bond. Lone pair An electron
pair found solely on a single atom. Single
covalent bond - Bond between two atoms when
they shared 1 pair Double covalent bond Bond
between two atoms when they shared 2
pairs. Triple covalent bond Bond between two
atoms when they shared 3 pairs. Lewis Structure,
Stability, Multiple Bonds, and Bond Energies
Bond order The stability of a covalent
compound is related to the bond energy. The
magnitude of the bond energy increases and the
bond length decreases in the order single bond
gt double bond gt triple bond. Bond Energy order
single lt double lt triple Bond length order
single (1) lt double (2) lt triple (3)
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Resonance

• Resonance structure 1 of 2 or more Lewis
  structures for a molecule (ion) that cant be
  represented with a single structure
• Resonance use of

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• Each resonance structure contributes to the
  actual structure
• no single structure is a complete description
• positions of atoms must be the same in each, only
  electrons are moved around
• actual structure is an average

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• Draw resonance structures for SO3 and N3-.

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Exceptions to Octet Rule

• There are three classes of exceptions to the
  octet rule.
•  
• 1) Molecules with an odd number of electrons
• 2) Molecules in which one atom has less than an
  octet
• 3) Molecules in which one atom has more than an
  octet.

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Lets do Lewis structures for

• CO2 (CS2)
• O3 (SO2)
• I3-

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3D structure of species

• Electrostatic forces in ionic bonds is
  _____________. But species with covalent bonds
  have electron pairs concentrated btn 2 atoms and
  is ..
• We use VESPR theory to predict the shape of the
  covalently bound species.

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VSEPR theory
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VSEPR

• Most stable geometry is one in which electron
  pairs (electron clouds) are as

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Shapes of molecules (3D)

• The geometry is determined by the atoms present
  in the species. See atoms that are bonded to
  other atoms. Dont see lone pairs but they
  influence geometry
• I. Diatomics (2 atoms only) always ________
• H2, HCl, CO X----X

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• II. Polyatomic (3 or more atoms) species
  Use VSEPR model to predict shapes

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Steps in applying VSEPR

• 1. Do Lewis structure
• 2. Count total e- pairs (clouds) around central
  atom (A). Multiple bonds count as one electron
  pair (cloud). In reality multiple bonds are
  bigger than single bonds (electron clouds
  larger).

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• 3. Separate e- pairs into bonded pairs (B) and
  lone pairs (E)
• 4. Apply table that I give you.
• 5. Remember that lone pairs of e-s are
  invisible, but their presence affects the final
  molecular geometry!!!!!
• Lone e- pair-lone e-pairs are more repulsive than
  bonded pair-lone pair repulsions or bonded
  pair-bonded pair repulsions.

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VSEPR valence shell electron pair repulsion

• 2 electron clouds around a central atom (A)

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2 electron clouds
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Three electron clouds
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Three electron clouds
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Four electron clouds
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Table 4.5 (changed)

• e bonded lone pairs geom
  angle clouds pairs pairs
• 2
• 3
• 3
• 4
• 4
• 4

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Predict geometry

• H2S
• SO2
• CO2
• CF4
• H2CO
• ClO3-
• ClO2-

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Polar vs nonpolar cmpds

• A molecule is polar if its centers of positive
  and negative charges do not coincide. If a
  molecule is polar we say that it acts as a
  dipole. In an electric field nonpolar molecules
  (positive and negative centers coincide) do not
  align with the field but polar molecules do.
• Next we will see why this happens and the
  implications.

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Molecules are subjected to electric field Polar
molecules align with field Nonpolar molecules are
not affected
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Polar molecules

• I. Diatomics, A-B
• a.If A B have homonuclear diatomic has
• b. A ? B have heteronuclear diatomic

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• II. Polyatomic species are more complicated.
• Lets look at VSEPR cases considered.
  
• General rule (my rule)

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Which of these are polar?

• H2S
• SO2
• CO2
• CF4
• AlCl3
• CHCl3
• SCl2

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Properties based on electronic structure and
molecular geometry

• Intramolecular forces within a molecule--bonds
• Intermolecular forces between molecules--these
  determine important properties as melting and
  boiling points and solubility

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Solubility

• Like dissolves like
• Polar cmpds dissolve in polar solvents as
  ionic and polar cmpds (HCl) in water
• Nonpolar cmpds dissolve in nonpolar solvents
  oils in CCl4

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Melting and boiling points

• Stronger the intermolecular forces the higher the
  melting and boiling points
• In general for cmpds of similar weight polar
  moleculaes have stonger forces than nonpolar
  cmpds
• In general for similar structure the greater the
  mass the stronger the forces

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Which have higher melting (boiling pts)

• CO and NO
• F2 and Br2
• CH3CH2OH and CH3CH3