1. Forces between particles
Chem 1151: Ch. 4
Forces between particles

2. Noble 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 (G. Lewis and W. 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.
Noble gases have filled s and p subshells.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

3. 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 2p 1s 1s 2s 1s 2s 3s

4. Valence 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.

5. Chemical Bonding
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 (gains 1 or more e-).
Usually satisfies octet rule
Common to inorganic chemistry
Ion: Atom or molecule that has either lost or gained electrons resulting in a net charge (positive or negative).
Ionization Energy: Energy required to remove an e- from an atom.
Covalent bond: Formed by sharing of electrons.
Occurs between:
Two non-metals
Nonmetal and metalloid
Two metalloids
Common to of organic chemistry

6. Chemical Bonding
Ionic bond: Attractive force that holds ions of opposite charge together.
Covalent bond: Formed by sharing of electrons.
Na  Na+ + 1e-
Cl + 1e-  Cl-

7. Ions
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.
Common atomic ions you should know:
H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+,
N3-, P3-, O2-, S2-, F-, Cl-, Br-
Trends here?
In many cases, 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.
How might this relate to formation of compounds?

8. Ionic Compounds
Ionic bond: Attractive force that holds ions of opposite charge together.
Involves transfer of e- from one component to the other.
Electron(s) cannot be lost from one atom unless there is another atom available to accept the electron(s).
Ionic compounds are formed when valence electrons lost by metal are gained by non-metal with which it is reacting.
Common atomic ions you should know:
H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+,
N3-, P3-, O2-, S2-, F-, Cl-, Br-
Common molecular ions you should know:
OH- (hydroxide), NH4+ (ammonium), SO42- (sulfate), SO32- (sulfite), PO43- (phosphate), NO2- (nitrite), NO3- (nitrate), CO32- (carbonate)
Ionic compounds formed by these ions will have neutral charges.
NaOH
NH3
HBr
Pb(OH)2
MgCl2
HCl
(NH4+)2SO4
H3PO4
AgNO2

9. Naming Ionic Binary Compounds
[Name of Metal] [nonmetal stem + ide] =
potassium (K+) chlorine (Cl-) 
potassium chloride (KCl)
strontium (Sr2+) oxygen (O2-) 
strontium oxide (SrO)
3 calcium (Ca2+) 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:
[Name of Metal] [nonmetal stem + ide] =
copper (Cu+) chlorine (Cl-) 
copper(I) chloride (CuCl)
iron (Fe2+) chlorine (Cl-) 
iron(II) chloride (FeCl2)
iron (Fe3+) chlorine (Cl-) 
iron(III) chloride (FeCl2)

10. Units 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

11. Lewis Structures for Covalent Compounds
Use molecular formula to determine how many atoms of each type.
Draw a structure with the elements relative to each other.
Determine total number of valence e- for all atoms.
Put one pair of e- between each bonded pair of atoms, subtract this number of e- from total.
Use remaining e- to form octets (except for H and He) for all atoms.
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-
Lone pairs or nonbonding e- Cl Cl Cl Cl
Use 2 e- to form bond
Now fill in remaining 12 e-

12. More Lewis Structures for Covalent Compounds
N + N  N2 N + N N N H
N + 3H  NH3 N + 3 H N H H
O + 2H  H2O H H + O + H H O

13. Still More Lewis Structures for Covalent Compounds
S + 3O  SO3
= 24 e- O O O O S O O S O O S O
= 18 e-
Confirm all octets satisfied.
Confirm all e- accounted for.

14. …And One More Lewis Structures for Covalent Compounds
C2H2
= 10 e- H C C H H C C H
= 4 e- H C C H H C C H
Confirm all octets satisfied.
Confirm all e- accounted for.

15. Lewis Structures for Polyatomic Ions
Follow the same instructions indicated for covalent compounds, but add or subtract the number of e- indicated by the ion charge.
Use molecular formula to determine how many atoms of each type.
Draw a structure with the elements relative to each other.
Determine total number of valence e- for all atoms.
Put one pair of e- between each bonded pair of atoms, subtract this number of e- from total.
Use remaining e- to form octets (except for H and He) for all atoms.
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).

16. Lewis Structures for Polyatomic Ions
SO42-
= 32 e- 2- O O O S O O S O O O
= 24 e-
Confirm all octets satisfied.
Confirm all e- accounted for.

17. Shapes 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:
All valence shell electron pairs around central atom are counted equally (both bonding and non-bonding).
Double or triple bonds are treated like a single bond when predicting shapes.

18. Shapes of molecules and Polyatomic Ions
Linear
Triangular
Tetrahedral

19. One Molecule can have multiple central atoms

20. Polarity of Bond Cl Cl H Cl Na Cl
Nonpolar Covalent Bond: Electrons forming bond between 2 atoms spend nearly equal time between both atoms.
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.
Ionic Bond: Electrons are transferred (highly polar).
Polarity indicated by arrow pointed towards destination of e- Cl Cl H Cl Na Cl
nonpolar
polar
ionic

21. Polarity of Molecules
Nonpolar Molecules: Charge distribution from bond polarizations is symmetric.
Polar Molecules (dipoles): Charge distribution from bond polarizations is not symmetric.
Need to consider all charges to determine direction of polarity.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

22. Naming Binary Covalent Compounds
Give the name of the less EN element first.
Give the stem of the name of the more EN element next, and add suffix –ide.
Indicate the number of each type of atom in molecule using numeric prefix. CO
carbon monoxide
N2O5
dinitrogen pentoxide
CO2
carbon dioxide
CCl4
carbon tetrachloride
H2O
dihydrogen monoxide
S2O7
disulfur heptoxide
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

23. Naming ionic compounds containing polyatomic ions
Give the name of the metal first.
Make sure that charges add up to zero.
Put parentheses around polyatomic ions if more than 1 used.
Write formulas for the following and name them:
K and PO43-
K3PO4
potassium phosphate
Na and NO3-
NaNO3
sodium nitrate
Mg and PO43-
Mg3(PO43-)2
magnesium phosphate
NO3- and NH4+
NH4NO3
Ammonium nitrate

24. Crystal Lattices of non-ionic compounds
Most pure substances (elements or compounds) form crystal lattices in solid state.
These lattice sites may be occupied by neutral atoms or molecules instead of ions.
Components of lattice sites held together by covalent bonds.
Network Solids: Lattice formed by atoms bound in covalent bonds (ex. Si and O).
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.
Copper
Graphite: Each Carbon is covalently bonded to 3 other carbons in ring
Diamond: Each carbon is bonded to 4 other carbons

25. Interparticle Forces δ-
Dipolar forces: Attraction between positive end of one polar molecule and negative end of another (usually weak, ~ kcal)
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)
Dispersion Forces: Momentary nonsymmetric electron distributions in molecules (very weak) δ+ δ-
Hydrogen bond δ+ δ- δ+
Intermolecular
dipolar δ- δ+