# High School Chemistry/Intra & Intermolecular Forces

Two major types of attractive forces exist between atoms:
1. Intermolecutive forces (exist between adjacent molecules) -- weaker
2. Intramolecutive forces (exist between atoms within a molecule or compound) - stronger
When water boils, what kind of forces are being broken? Why? Intermolecular Forces; because they're weaker than intramolecular forces

## Bonds

Before we divulge into the types of bonds, a few rules need to be made clear:
* Formation of a bond = Exothermic process where energy is released.
* Breaking of a bond = Endothermic process where energy is absorbed. It takes 566J to break a bond.
The type of bond atoms form is determined by the difference in electronegativity. This is where memorizing your periodic trends come in handy.
See w:Bond-dissociation energy.

### Chemical Bond

Chemical bond - An attraction between atoms that allow for the formation of a chemical substance. Valence electrons are the part of the atom that is involved in bonding. Bonding occurs so that each atom can attain an octet/get stable (8 valence electrons). There are three types of chemical bonding patterns/intramolecular forces: Metallic, Ionic, and Covalent.

#### Metallic bonding

    See also: Metallic Bond Properties (talk page)
Metallic bonding - Metal atoms lose electrons to form positively charged cations due to their low electronegativity and ionization energy values. This causes a sea of mobile and flowing valence electrons.
We can tell if it is:
* Just one type of metal (Zn, Cu, Au, etc.)
* Metals have very high melting and boiling points due to their strong bonding.
* Fluid "sea" of electrons results in malleability, ductility, conductivity, and low solubility.
* Alloys: Homogenous mixtures (solutions) of two or more elements, at least one of which is a metal (increases strength) -- steel, brass, bronze, gold jewelry, pewter, etc.

#### Ionic bonding

    See also: Ionic Bond Properties (talk page)
Ionic bonding - Involves the transfer of electrons between a metal or one or more nonmetal (ex: CaCO$_3$) due to the large differences in electronegativity between the elements involved (nonmetals have higher electronegativity values than metals). The force is so powerful that one atom removes the electron of the other atom (transfer). There is a redistribution of the outer electron clouds.
* Cations (+) and anions (-) are formed and are attracted to each other due to the electrostatic force. Examples are in NaCl (table salt), Na$_2$SO$_4$ (sodium sulfate), etc.
* Also known as salts, compounds, and formula units.
* All ionic bonds are solid brittle crystals at room temperature, forming crystal lattices with repeating unit cells.
* When molten (liquid NaCl) or dissolved in water (NaCl$_{(Aq)}$), ionic compounds conduct electricity and are called electrolytes because free ions are present.

#### Covalent bonding

    See also: Covalent Bond Properties (talk page)
Covalent Bonding - involves sharing of electrons due to relatively small differences in electronegativity between elements involved. Covalent bonds ALWAYS occur between two nonmetals, like sulfur dioxide (SO$_2$) and water (H$_2$O).
Electrons can be either shared equally or shared unequally depending on their electronegativity difference:
* If the EN difference is from 0-0.4, it is a nonpolar covalent bond.
* If the EN difference is from 0.5-1.7, it is a polar covalent bond.
* If the EN difference is above 1.8+, it is a ionic bond.

##### Examples

* H$_2$: 2.2-2.2 = 0 → NP (Nonpolar Covalent Bond)
* NaCl: 3.2 - 0.9 = 2.3 → Ionic Bond
* SO$_3$: 3.4 - 2.6 = 0.8 → P (Polar Covalent Bond)

## Review #1
## Drawing Lewis Structures

* Steps to Drawing Lewis Structures
* Using the periodic table, figure out the total number of valence electrons for your molecules. For ions, add it is an anion and subtract if it is a cation. The total number should be the final structure.
* Divide by 2 to get the number of total electron pairs you will be working with on your final structure.
* Draw a skeleton structure by choosing a central atom (Hydrogen and Fluorine are always terminal, and can never be in the middle; this must be the molecule with the lowest EN values), and hanging all the other atoms off with a sigma bond.
* Subtract the pairs used to make these bonds: this is the remaining lone pairs you've to work with.
* Place these remaining lone pairs around each terminal atom, except hydrogen, to satisfy the octet rule. If they're remaining lone pairs, add them to the central atom [if the central atom is from the 3rd or higher period, it can be assigned more than four electrons pairs-six four electrons pair]. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs (atoms attached to the central atom) to a double/triple (pi bond) air. NOTE that not all elements can form pi bonds! These elements are carbon, nitrogen, oxygen, phosphorus, and sulfur (CNOPS).
* See also
* How to Draw Lewis Structures: Five Easy Steps - YouTube
* Lewis Diagrams Made Easy: How to Draw Lewis Dot Structures - YouTube

### Exceptions (Octet Rule)
#### Fewer than eight

* Hydrogen: Stable with 2 electrons (only one bond)
* Beryllium: Stable with 4 electrons (only 2 bonds)
* Boron: Stable with 6 electrons (only 3 bonds)

#### Expanded Valence

Can only happen if the central element has d-orbitals, which means it is from the 3rd period or greater. Therefore, it can be surrounded by more than 4 valence pairs in certain compounds. The number of bonds depends on the balance between the ability of the nucleus to attract electrons and the repulsion between the pairs. Look at SF$_6$.

#### Odd Electron Compounds

A few stable compounds contain an odd number of valence electrons and thus cannot obey the octet rule, NO, NO$_2$, and ClO$_2$.

### Bond Terminology

* Single Bond - 1 pair of electrons shared (meaning 1 sigma bond)
* Double Bond - 2 pairs of electrons shared (meaning 1 sigma, 1 pi bond)
* Triple Bond - 3 pairs of electrons shared (meaning 1 sigma, 2 pi bonds)
Multiple bonds increase electron density between two nuclei, therefore enhancing the nucleus to electron density attractions.

## Determining the Geometrical Shape based on the VSPER Theory

* Only true if the element is from the 3rd energy levels + d levels (s and p).
The shapes of covalently bonded molecules and ions can be determined by considering the number of lone vs. bonding pairs around the central atom. The electron pairs repel one another and try to get as far apart as possible, causing unique shapes. This theory is known as the Valence Shell Electron Pair Repulsion Theory, or VSEPR Theory.

### VSPER Shapes

    See VSPER Shapes
* Linear - 180 degrees (angle)
* Bent - 90°<θ<120°
* Trigonal planar - 120 degrees (angle)
* Trigonal pyramidal - 90°<θ<109.5° degrees (angle)
* Trigonal bipyramidal - 90°, 120° degrees (angle)
* Tetrahedral - 109.5 degrees (angle)
* Seesaw - 90°, 90°<θ<120° degrees (angle)
* Octahedral - 90 degrees (angle)

## Polar Bonds vs. Polar Molecules
### Bond Polarity

Bond Polarity is based on differences in electronegativity.
A dipole is created by equal and opposite charges separated by a short distance. The direction of the dipole is from the positive (less electronegative atom) to negative (more electronegative atom) and is represented by an arrow pointing towards the negative pole and crossed tail situated at the positive pole. The larger the difference in electronegativity between the two bonded atoms, the greater the dipole.

### Molecular Polarity

    See also: Molecular Polarity - PennState
Molecular Polarity is based on whether the molecule has a net dipole movement (a negative end and a positive end). A molecule must have polar bonds in order to be a polar molecule, though if there are lone pairs on the central atom, it will make it a polar molecule.
However it is also possible for a molecule to have polar bonds, however be nonpolar overall. In this case, the dipoles present cancel each other out because they are pulling in opposite directions. CO$_2$ is an example of a nonpolar molecule since it is linear and oxygen is more electronegative than carbon.
All the diatomics and tetrahedral structures (such as C$_3$H$_8$) are nonpolar.

## Molecular vs. Electronic Geometry, Hybridization, and Connections to Bond Angles

Molecular Geometry (VSPER Shape): determined by the number of bonding vs. nonbonding (or lone) pairs around the central atom.
Electronic Geometry: determined by the number of regions of electron density around central atom (these count as just one region of electron density: a lone pair, a single/double/triple bond).
Hybridization is used to explain how the orbitals of an atom become rearranged when atoms form their covalent bonds.
| Regions of electrons density | Type of hybridization |
| 2 | sp |
| 3 | sp$^2$ |
| 4 | sp$^3$ |
| 5 | sp$^3$d |
| 6 | sp$^3$d$^2$ |

## Intermolecular Forces

Remember that intermolecular forces are forces that exist between adjacent molecules.
London-dispersion forces (LD forces) are present in ALL molecules and atoms.
A dipole-dipole force is an intermolecular force (force between adjacent molecules) which occurs only between polar molecules (molecules with partial positive and negative poles).
In acetone (C$_3$H$_6$O), the negatively charged oxygen is attracted to the positively charged carbon (electronegativity values). The dipole-dipole force pulls these molecules together and keeps acetone a liquid.
Hydrogen bonds also occur between two polar molecules, but only when the hydrogen is directly attatched to a oxygen, flourine or nitrogen atom. It is a very strong intermolecular force.

### Boiling Points

Boiling points increase with stronger force + more electrons (makes the force/bond stronger).