1. Into to chemical bonding

2. Intro to chemical bonding
A mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

3. Two types of chemical bonds
Ionic bonding
Bonding that results from the electrical attraction between large numbers of cations and anions
Covalent bonding
Results from the sharing of electron pair between two atoms

4. Wait it’s not that simple
Bonding isn’t usually purely ionic or covalent it usually falls somewhere in between.
The electronegativity is a measure of an atom’s ability to electrons
Therefore more time is spent by the electron around the more electronegative atom.

5. Two types of covalent bonds
Nonpolar-covalent bond
This bond is a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. NO POLES
Polar-covalent bond is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons. HAS POLES LIKE A MAGNET

6. Non polar polar H H H H + O + -
End 6-1

7. Molecular compounds with covalent bonding
A molecule is a neutral group of atoms that are held together by covalent bonds.
A molecular compound is a chemical compound whose simplest units are molecules.
A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts.

8. More bonding
A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound
A diatomic molecule is a molecule containing only two atoms.

9. Why?
Chemicals bond because they have a lower potential energy when bonded then when they are alone.
Remember atoms want to lowest energy always
Different chemicals have different energies so bond length of those molecules may differ
Bond length is the distance between tow bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms.

10. Can we break the bond
Bonds can be broken but it takes energy to do this.
This energy is called bond energy.
The energy required to break a chemical bond and form neutral isolated atoms.
Reported in KJ/mol
So what we learned is that bond lengths and energies differ based on what and how atoms are bonded together.

11. The octet rule What is an octet?
What does oct mean?
So an octet is a set of how many electrons?
These valence electrons are in the outer shell of the atom.
Atoms will gain lose or share electrons to fill their outer shells with an octet.
As always there are exceptions to the octet rule.
These atoms can be happy with more or less than an octet in their outer shell.

12. Electron-dot notation
An electron-configuration notation in which only the valence electrons of an atom of a particular element are shown indicated by dots place around the element’s symbol.
Example
Fluorine F

13. Lets do these on the board
Carbon
Potassium
Sulfur
Argon
Lithium
Hydrogen
Helium

14. Lewis Structures Expands on electron-dot notation
Lewis structures are formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons
In Lewis structures the lone dots represent an unshared pair of electrons. Unshared pair is also called a lone pair.
A pair of electrons that is not involved in bonding and that belongs exclusively to one atom

15. Lets do some Lewis structures
CH3I
H2O
NH3
H2S

16. Structural formula
This is the Lewis formula without the electron pairs
Indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule.
Just dashes no dots

17. Bonds and dashes
A single bond is a covalent bond produced by the sharing of one pair of electrons between two atoms.
Represented by one dash (-)
A double bond is a covalent bond produced by the sharing of two pairs of electrons between two atoms
Represented by two dashes (=)
A triple bond is a covalent bond produced by the sharing of three pairs of electrons between two atoms.
Represented by three dashes(=)
Double and triple bonds are reffered to a multiple bonds.

18. Lets do some multiple bonding
CH2O
CO2 CO
HCN

19. Resonance structures
Refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure…….
Ozone is one of these O3
Lets do it

20. Ionic bonding
Ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal.
Ionic compounds are represented by a formula unit
The simplest collection of atoms from which an ionic compound’s formula can be established.

21. Characteristics of ionic bonding
Remember that molecules want the lowest potential energy.
In ionic molecules this means forming a crystal and then those crystals forming a lattice
Crystal lattice
Form by the attractive forces at work
Opposite charges, like charges and electrons
Table salt looks like this

23. More about Lattice structures
Lattice structures are three-dimensional arrangements of ions
Strengths and sizes of these structure differ by ionic compound
Bonds strengths in ionic compounds is known as lattice energy
The energy released when one mole of an ionic crystalline compound is formed from gaseous ions.

24. Ionic vs. Molecular compound
Greater attraction between molecules
Higher melting points
Higher boiling points
Do not vaporize at room temperature
Hard but brittle
Ions cant move making it hard
But if they do it breaks
Lower attraction between molecules
Lower melting points
Lower boiling points
Most vaporize to gas at room temperature
Malleable
Compounds can move
Bend don’t break

25. Polyatomic ions These bond covalently and ionic
Becomes a charged group of compounds
Shown with charges
Some Lewis structures of polyatomic ions

26. Metallic bonding Different than in other types of bonding
Give metals their properties
Good conductors of electricity
This property due to the highly mobile valence electrons of the atoms that make up a metal
With empty shells in the P-block and or D-block that overlap
The electrons can roam freely
Delocalized –not belonging to any one electron
The attraction between metal atoms and the surrounding sea of electrons is called metallic bonding

27. Properties of metals….again
Good conductors of heat and electricity
Due to the fact the electrons are free to move in the network of metal atoms
Shiny (luster)
The free electrons absorb energy from light then release it giving metals their shine
Malleability and Ductility
Due to the fact that metallic bonding is the same in all directions.
Atoms can slide past another without encountering any resistance or breaking any bonds

28. Metallic bond strength
Varies with nuclear charge and the number of electrons in the metal’s electron sea
Heat of vaporization
Atoms in the normal state (usually solid) are converted to individual metal atoms in the gaseous state.
The amount of heat required is a measure of the strength of the bonds that hold metal together.

29. Molecular geometry This is the shape of molecules
The three-dimensional arrangement of a molecule’s atoms in space
The polarity of each bond along with the shape determines
Molecular polarity
The uneven distribution of molecular charge

30. VSEPR Theory Stands for Valence-shell, electron-pair repulsion
This theory states
Repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be orientated as far apart as possible.
Gives us all the different shapes of molecules
Linear, bent or angular, trigonal-planar Tetrahedral, trigonal-pyramidal, octahedral, trigonal-bipyramidal

32. Hybridization
The mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies
Hybrid orbitals
Orbitals of equal energy produced by the combination of two or more orbitals on the same atom

33. Intermolecular forces
The forces of attraction between molecules
Dipole-Dipole forces
A dipole is created by equal but opposite charges that are separated by a short distance
In polar molecules the attraction between particles are dipole-dipole forces
This is what keeps water together
These forces raise boiling points

34. More Intermolecular forces
Hydrogen bonding
The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule
Represented by dotted lines
London dispersion forces
The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles.
Named for Fritz London
Only forces acting on noble gases and nonpolar molecules
Results in the low boiling points of the noble gases
Strength increases with the number of electrons
This can be seen in the boiling point of the noble gases
They increase down the group