1. Unit 10 Chemistry Langley
BONDING AND GEOMETRY
Unit 10
Chemistry
Langley
**Corresponds to Chapter 7 and 8 (pages ) in the Prentice Hall Chemistry textbook

2. PERIODIC TABLE REVIEW
Location of Metals and Nonmetals on the periodic table:
Metals are to the left of the “staircase”
Nonmetals are to the right of the “staircase”
For bonding, the 7 metalloids will treated as metals
All though hydrogen is to the left of the “staircase”, it is not, nor has it ever been a metal. IT IS A NONMETAL!

3. ATOMS AND IONS REVIEW Atoms are neutral Ions have a charge
They have the same number of protons and electrons
Number of positives = number of negatives
Example: Na 11 protons, 11 electrons  11 – 11 = 0
Ions have a charge
They have a different number of protons and electrons
Example: Na+111 protons, 10 electrons 11 – 10 = +1
If an atom GAINS an electron  becomes negatively charged  ANION
If an atom LOSES an electron  becomes positively charged  CATION

4. TYPES OF BONDS
Bonding occurs because every element is either trying to get to 0 electrons in the valence or 8 electrons in the valence (zero and 8 are both stable)
Valence is the outer electron shell—place where bonding occurs
Ionic – Bonding between a metal and a nonmetal
Metallic – Bonding between two metals
Covalent – Bonding between two nonmetals

5. IONIC BONDING Very stable and strong Strongest possible bond
Requires a large amount of energy to break an ionic bond
Forms compounds known as “ionic compounds”
All ionic compounds will dissolve in water and carry a current (electrolyte)
Generally have high melting and boiling points
Compounds are generally hard and brittle

6. IONIC BONDING Draw the dot diagram for Na AND Cl
Na has 1 valence electron, wants to give that 1 away and get to zero and be stable
Cl has 7 valence electrons, wants to get 1 electron so it can get to eight and be stable
Na give an electron to Cl and Cl takes that electron from Na

7. Give the Lewis electron-dot symbol for each of the following atoms
Sodium
Magnesium
Chlorine
Oxygen

9. Give the Lewis electron-dot symbol for each of the following IONS
Sodium ion
Magnesium ion
Chlorine ion
Oxygen ion

10. METALLIC BONDING
Metal atoms are pieces of metal that consist of closely packed cations (positively ions)
Cations are surrounded by mobile valence electrons that are free to drift from one part of the metal to another
Metal atoms are arranged in very compact and orderly (crystalline) patterns
Metallic bonding is the electrostatic attraction between conduction electrons, and the metallic ions within the metals, because it involves the sharing of free electrons among a lattice of positively-charged metal ions
Occurs between 2 or more metals
Result of the attraction of free floating valence electrons for the positive ion
These bonds hold metals together

11. METALLIC BONDING Properties of metallic bonds
Good conductors of electricity
Electrons are free flowing
Malleablehammered into sheets
Ductiledrawn into wires
Alloy-two metals are bonded together to get the benefits of each
14 karat gold

12. COVALENT BONDING Covalent: A covalent compound is called a molecule
Prefix “co” means share, together
“valent” means valence
Covalent bonds are when atoms SHARE VALENCE electrons
A covalent compound is called a molecule
Covalent bond ALWAYS occurs between 2 nonmetals

13. TYPES OF COVALENT BONDS
Single Bond
Covalent bond where one pair of electrons (2 electrons total) are shared between 2 atoms
Atoms share electrons so that each has a full octet (8 valence)
Electrons that are shared count as valence electrons for both atoms
Examples
HCl
Cl2

14. COVALENT BONDING Double Bonds
Bond in which two pairs of electrons (4 electrons total) are shared between 2 atoms
Examples O2
C2F2
Triple Bonds
Bond in which 3 pairs of electrons (6 total electrons) are shared between atoms N2
AsP

15. COVALENT BONDING Covalent Bonds with more than 2 atoms Electron Pairs
Examples
CH4
OF2
Electron Pairs
Electron pairs involved in the actual bond are called BONDING PAIR or SHARED PAIR electrons
Electrons not involved in the actual bond, those surrounding the rest of each element are called LONE PAIR electrons

16. POLAR BONDS AND MOLECULES
Covalent bonds are formed by sharing electrons between two atoms
The bonding pair of electrons is shared between both elements, but each atom is tugging on the bonding pair
When atoms in a molecule are the same (diatomic) the bonding pair is shared equallythis bond is called non polar covalent
When atoms in a molecule are different, the bonding pair of electrons are not shared equallythis is called a polar covalent bond

17. POLAR BONDS AND MOLECULES
Why is the bonding pair not shared equally?
The answer lies within electronegativity
One of the elements is more electronegative than the other and therefore has a greater desire for the shared pair
The MORE electronegative element tends to pull the electrons closer and thus has a slightly negative charge
The LESS electronegative element has a slightly positive charge since the shared pair is being pulled away

18. POLAR BONDS AND MOLECULES
Drawing/Indicating Polarity

19. POLAR BONDS AND MOLECULES
Polar Molecules
Molecule in which one end of the molecule is slightly negative and the other end is slightly positive
Just because a molecule contains a polar bond DOES NOT mean the entire molecule is polar
The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds

20. POLAR BONDS AND MOLECULES
Example: CO O = C = O
Carbon and Oxygen lie along the same axis.
Bond polarities are going to cancel out because they are in opposite directions
Carbon dioxide is a nonpolar molecule even though there are two polar bonds present
Would cancel out if the polarities moved towards each other as well
When polarities cancel out, the molecule is non-polar

21. POLAR BONDS AND MOLECULES
Example: H2O
Example: CH3Br

22. FORCES IN A MOLECULE Dipole-Dipole Forces
Dipoles are created when equal but opposite charges are separated by a short distance
Have to have a positive and a negative end so that one of the elements is pulling on the electron
Only happens in polar molecules
Dipole forces are extremely strong and lead to high melting and boiling points

23. FORCES IN A MOLECULE Hydrogen Bonding Very strong type of dipole force
Only occurs when hydrogen is covalenty bonded to a highly electronegative atom
Always involves hydrogen
Example: HF, HCl

24. FORCES IN A MOLECULE London Dispersion Forces
Electrons are in constant motion around a nucleus
At any given time there might be more electrons on one side of an atom than on the other
For a split second, the side with more electrons is negative, and the side with less electrons is positive

25. FORCES IN A MOLECULE London Dispersion Forces
Recall that Noble Gases have a full outer shell and you have been told they are unreactive BUT due to London Dispersion Forces, they COULD bond for an instant
Example: Ar2
London Forces are very weak
The smaller the mass of the atom, the smaller the London Force

26. BOND DETAILS Terminology
Bond strength-energy required to break a bond
Bond axis-imaginary line joining two bonded atoms (example: C-C)
Bond length-the distance between two bonded atoms at their minimum potential enery; the average distance between two bonded atoms
Bond energy-energy required to break a chemical bond and form neutral isolated atoms
Chemical compound tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level

27. BOND DETAILS
Comparison of Bond Length/Strength for Covalent Bond Types:
Longer bond = less bond strength
Rating 1-3 (with 3 as the largest and 1 as the smallest)
Bond
Length
Strength
Single 3 1
Double 2
Triple

28. BOND DETAILS Coordinate Covalent Bonds Very rare
Tend to form harmful molecules
Occurs when both of the bonding pair of electrons in a covalent bond come from only ONE of the atoms
Example: CO

29. BOND DETAILS Resonance
Occurs when there are more than one possible structures for a molecule
Refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure
Example: CO2
To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures
Even though all of the structures are different, the number of bonding pair of electrons and lone pair of electrons stay the same in each structure

30. VSEPR THEORY Valence Shell Electron Pair Repulsion Theory
Allows us to picture molecules in 3 dimensions
Centers around the fact that electrons have negative charges and repel one another
So electron pairs within a structure try to arrange themselves to be as far away from other pairs as possible

31. VSEPR THEORY Tetrahedral
Central atom bonds to 4 atoms and has zero lone pairs
CH4

32. VSEPR THEORY Pyramidal
The central atom bonds to 3 atoms and has 1 lone pair of electons
NH3

33. VSEPR THEORY Trigonal Planar
The central atom bonds to 3 atoms and has zero lone pairs
CO3-2

34. VSEPR THEORY Bent Triatomic
The central atom bonds to 2 atoms and has 2 lone pair of electrons
H2O

35. VSEPR THEORY Linear Triatomic
The central atom bonds to 2 atoms and has zero lone pair of electrons
CO2

36. VSEPR THEORY
Linear
One bond between 2 atoms
HCl N2