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Chem 1151: Ch. 4
Forces between particles
Noble Gas ConfigurationNoble Gas Configuration• The first widely-accepted theory for chemical bonding was based on noble
gas configurations.• Noble gases have a filled valence shell (2e- for He and 8 e- for everything
else).• Because noble gases are very unreactive, and chemical reactivity depends on
electronic structure, two scientists (Lewis and Kossel) concluded that this represented stable (i.e., low energy) configuration.
• Formed the basis for the “octet rule” (1916): 8 electrons in valence shell results in greater stability.
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Noble gases have filled s and p subshells.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
Valence Shell e-Valence Shell e-• Valence shell is the outermost shell where electrons reside.• From the electronic configuration (or other means) you can determine
that the valence shell will have the highest number n.• This means that for transition metals, even though the last electrons may
be added to the d subshell for n, the valence shell e- will be found in the s subshell at n+1.– Ex. Yttrium (Y) 1s22s22p63s23p64s23d104p65s24d1
1s 1s 2s 1s 2s 3s
2p
Valence Shell e- in Lewis StructuresValence Shell e- in Lewis Structures• Lewis Structure (electron dot formula): Simplified way to represent
valence shell e-.• Element symbol represents nucleus, with valence shell e- represented by
dots surrounding it.• The number of valence e- can be found by looking at the group number,
except for the transition metals, which all have filled s orbitals at n+1.
O
C
Sr
In
Po
4
Atom # Val Shell e-
2
5
6
C
Sr
In
Po
O 6
Lewis Structure
IonsIons Ion: Atom or molecule that has either lost or gained electrons from
valence shell resulting in a net charge (positive or negative) compared to the number of protons.
Ionization Energy: Energy required to remove an e- from an atom.
Common atomic ions you should know: H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, N3-, P3-, O2-, S2-, F-, Cl-, Br-
For Group A elements, the number of e- gained or lost results in an electronic configuration like that of the noble gases (valence shell octet). If Na lose e-, has electronic configuration of Ne. If Cl gains e-, has electronic configuration of Ar. If Ca loses 2 e-, has electronic configuration of Ar.
Ionic CompoundsIonic Compounds Ionic compounds are formed when valence electrons lost by metal are
gained by non-metal with which it is reacting. Electron(s) cannot be lost from one atom unless there is another atom
available to accept the electron(s). Common atomic ions you should know: H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, N3-, P3-, O2-, S2-, F-, Cl-, Br-
Common molecular/polyatomic ions you should memorize: OH- (hydroxide), NH4
+ (ammonium), SO42- (sulfate), SO3
2- (sulfite), PO4
3- (phosphate), NO2- (nitrite), NO3
- (nitrate), CO32- (carbonate)
Ionic compounds formed by these ions will have neutral charges.
Pb(OH)2
NaOH
(NH4+)2SO4
NH3
MgCl2
H3PO4
HBr
HCl
AgNO3
Chemical BondingChemical Bonding1. Ionic bond: Attractive force that holds ions of opposite charge together.
– Involves transfer of e- from one component to the other.– Occurs between positively-charged metal (loses 1 or more e-) and non-metal
atom or molecule (gains 1 or more e-).– Usually satisfies octet rule– Common to inorganic chemistry
2. Covalent bond: Formed by sharing of electrons.– Occurs between:
• Two non-metals• Nonmetal and metalloid• Two metalloids
– Usually satisfies octet rule– Common to organic chemistry
Chemical BondingChemical Bonding1. Ionic bond: Attractive force that holds ions of opposite charge together.
2. Covalent bond: Formed by sharing of electrons.
Na Cl+ ClNa
1+ 1-
EN: 0.9 EN: 3.0
0.79 Å 0.91 Å
2.23 Å 0.97 Å
Na Cl
H C
Na Cl
H C
4H C+ CHEN: 2.1 EN: 2.5
H
HH
0.79 Å 0.91 Å
2.23 Å 0.97 Å
Na Cl
H C
Na Cl
H C
Electron-half-equations (for redox rxns)Electron-half-equations (for redox rxns)
Na Cl+ ClNa
1+ 1-
EN: 0.9 EN: 3.0
0.79 Å 0.91 Å
2.23 Å 0.97 Å
Na Cl
H C
Na Cl
H C
Na Na+ + 1e-
Cl + 1e- Cl-
oxidation
reduction
Each of these represents ½ of the rxn
Naming Ionic Binary CompoundsNaming Ionic Binary Compounds
potassium (K+) + chlorine (Cl-)
[Name of Metal] + [nonmetal stem + ide] =
potassium chloride (KCl)
strontium (Sr2+) + oxygen (O2-) strontium oxide (SrO)
3 calcium (Ca2+) + 2 nitrogen (N3-) calcium nitride (Ca3N2)
Some metals may form more than 1 type of charged ion. Exs: Cu+ and Cu2+; Fe2+ and Fe3+
Compounds with these ions would be named by adding a roman numeral equivalent to charge in parentheses after metal name:
copper (Cu+) + chlorine (Cl-)
[Name of Metal] + [nonmetal stem + ide] =
copper(I) chloride (CuCl)
iron (Fe2+) + 2 chlorine (Cl-) iron(II) chloride (FeCl2)
iron (Fe3+) + 3 chlorine (Cl-) iron(III) chloride (FeCl2)
Units of Ionic CompoundsUnits of Ionic Compounds
• Stable form of ionic compound is not a molecule, but a crystal lattice where ions occupy lattice sites.
• Molecular compounds have molecular weight
• Ionic compounds have formula weight
• Although ionic compounds form crystal lattices, we still represent them the same as molecular compounds when discussing formula weight
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
Lewis Structures for Covalent CompoundsLewis Structures for Covalent Compounds1. Use molecular formula to determine how many atoms of each type.2. Draw a structure with the elements relative to each other.3. Determine total number of valence e- for all atoms.4. Put one pair of e- between each bonded pair of atoms, subtract this number of e-
from total. 5. Use remaining e- to form octets (except for H and He) for all atoms.6. If all octets cannot be satisfied with available e-, move nonbonding e- pairs between
bonded atoms to complete octets (will form double or triple bonds).
Cl + Cl Cl2 7 + 7 = 14 e-
Cl ClUse 2 e- to form bond
Cl ClNow fill in remaining
12 e-
Lone pairs or nonbonding e-
More Lewis Structures for Covalent CompoundsMore Lewis Structures for Covalent CompoundsN + N N2
N + N N N
N + 3H NH3
N + H N H3H
HO + 2H H2O
H + O + H OH
H
Still More Lewis Structures for Covalent CompoundsStill More Lewis Structures for Covalent Compounds
S + 3O SO3
SO
O
6 + 6 + 6 + 6 = 24 e-
O24 - 6 = 18 e-
SO
O
O SO
O
O
1. Confirm all octets satisfied.2. Confirm all e- accounted for.
……And One More Lewis Structures for Covalent And One More Lewis Structures for Covalent CompoundsCompounds
C2H2
C
4 + 4 + 1 + 1 = 10 e-
C
1. Confirm all octets satisfied.2. Confirm all e- accounted for.
HH10 - 6 = 4 e-
C C HH
C C HH C C HH
Lewis Structures for Polyatomic IonsLewis Structures for Polyatomic Ions1. Follow the same instructions indicated for covalent
compounds, but add or subtract the number of e- indicated by the ion charge.
2. Use molecular formula to determine how many atoms of each type.
3. Draw a structure with the elements relative to each other.4. Determine total number of valence e- for all atoms.5. Put one pair of e- between each bonded pair of atoms,
subtract this number of e- from total. 6. Use remaining e- to form octets (except for H and He) for all
atoms.7. If all octets cannot be satisfied with available e-, move
nonbonding e- pairs between bonded atoms to complete octets (will form double or triple bonds).
Lewis Structures for Polyatomic IonsLewis Structures for Polyatomic Ions
SO42-
SO
O
6 + 6 + 6 + 6 +6 + 2 = 32 e-
O
32 - 8 = 24 e-
SO
O
O
1. Confirm all octets satisfied.2. Confirm all e- accounted for.
O O
2-
Lewis Structures for Polyatomic IonsLewis Structures for Polyatomic Ions
PO43-
PO
O
5 + 6 + 6 + 6 +6 + 3 = 32 e-
O
32 - 8 = 24 e-
PO
O
O
1. Confirm all octets satisfied.2. Confirm all e- accounted for.
O O
3-
Shapes of molecules and Polyatomic IonsShapes of molecules and Polyatomic Ions
Molecules have 3-D shapes.•VSEPR (valence-shell electron-pair repulsion) Theory: Electron pairs in valence shell of atoms are repelled by other pairs and try to get as far away from each as possible. •Shapes around central atom (any atom bonded to 2 or more other atoms) can be predicted by VSEPR.
2 Rules for determining shape:1.All valence shell electron pairs around central atom are counted equally (both bonding and non-bonding).2.Double or triple bonds are treated like a single bond when predicting shapes.
Electron Pair ArrangementsElectron Pair Arrangements
• According to the VSEPR theory, the arrangement of electron pairs around the central atom (represented by E) depends on the number of electron pairs.– Two pairs locate opposite each other. – Three pairs arrange themselves in a flat triangle around the
central atom. – Four pairs become located at the four corners of a pyramid-
like shape called a tetrahedron.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
Shapes of molecules and Polyatomic IonsShapes of molecules and Polyatomic Ions
LinearLinear
TriangularTriangular
TetrahedralTetrahedral
OO N+
OO C CH CH ClCl C C
ClCl
CC
CH3
CH3F
F
F
B
Cl
ClCl
Cl
C H
HOH
Cl
CH3
CC
CH3HH
O
H
H
HH N+
One Molecule, Different GeometriesOne Molecule, Different Geometries
CH2OH
O CH3
C
C
CH3
C
C
CH3
Polarity of BondPolarity of Bond1. Nonpolar Covalent Bond: Electrons forming bond between 2
atoms spend nearly equal time between both atoms. 2. Polar Covalent Bond: Covalent bond where electrons
polarized and remain closer to atom with higher EN (bond polarization).– Electronegativity (EN): Ability of an atom to attract shared e- of
covalent bond.
3. Ionic Bond: Electrons are transferred (highly polar).
Cl Cl H Cl
nonpolar polar
Na Cl
ionic
1.2EN1.21.0 EN0EN
Polarity indicated by arrow pointed towards destination of e-
Polarity of MoleculesPolarity of Molecules1. Nonpolar Molecules: Charge distribution from bond
polarizations is symmetric. 2. Polar Molecules (dipoles): Charge distribution from bond
polarizations is not symmetric. 3. Need to consider all charges to determine direction of
polarity.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
Binary Covalent CompoundsBinary Covalent Compounds
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
• Hazards of dihydrogen monoxide• http://www.dhmo.org/facts.html
Naming Binary Covalent CompoundsNaming Binary Covalent Compounds
1. Give the name of the less EN element first.
2. Give the stem of the name of the more EN element next, and add suffix –ide.
3. Indicate the number of each type of atom in molecule using numeric prefix.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
CO
CO2
H2O
N2O5carbon monoxide
carbon dioxide
dihydrogen monoxide
dinitrogen pentoxide
CCl4 carbon tetrachloride
S2O7 disulfur heptoxide
Naming ionic compounds containing polyatomic ionsNaming ionic compounds containing polyatomic ions1. Give the name of the metal first.2. Make sure that charges add up to zero.3. Put parentheses around polyatomic ions if more than 1 used.
potassium phosphateK and PO43-
Mg and PO43-
K3PO4
sodium nitrateNa and NO3- NaNO3
Mg3(PO43-)2 magnesium phosphate
NO3- and NH4
+ NH4NO3 Ammonium nitrate
Write formulas for the following and name them:
Crystal Lattices of non-ionic compoundsCrystal Lattices of non-ionic compounds1. Most pure substances (elements or compounds) form crystal lattices in solid state.2. These lattice sites may be occupied by neutral atoms or molecules instead of ions.3. Components of lattice sites held together by covalent bonds.4. Network Solids: Lattice formed by atoms bound in covalent bonds (ex. Si and O).5. Metallic Bond: Lattice of metal ions
– In this structure, valence e- can move about more freely, which is how metals can conduct heat and electricity.
Graphite: Each Carbon is covalently bonded to 3 other carbons in ring
Diamond: Each carbon is bonded to 4 other carbons http://www.eduys.com/Copper-
Molecular-Structure-Model-303.html
Copper
Interparticle ForcesInterparticle Forces
1. Dipolar forces: Attraction between positive end of one polar molecule and negative end of another (usually weak, ~0.5 -2.0 kcal)
2. Hydrogen Bonding: Attractive interaction of a hydrogen atom with an electronegative atom (e.g. N, O, F) from another molecule or chemical group.– The H must be covalently bonded to
another electronegative atom. – Stronger than dipolar and dispersion (12-
16 kcal)3. Dispersion Forces: Momentary
nonsymmetric electron distributions in molecules (very weak)
O
R R
O
R R
δ-
δ-
δ+
δ+
Intermoleculardipolar
H HO
Hydrogen bond δ+
δ-
H HO
δ+
δ-
Relative Strengths of Interparticle ForcesRelative Strengths of Interparticle Forces
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011
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