1. Unit 7: Bonding

2. Why Bond?
Elements bond in order to get a full valence shell
Elements want to have the same number of electrons as the noble gases
The Octet Rule: in forming compounds, atoms tend to achieve the electron configuration of a noble gas
**** only VALENCE electrons are involved in bonding

3. Types of Bonds
Ionic
Covalent
Metallic

4. Review of Metals Metals lose electrons to form positive ions (cations)
Ionic radius is smaller than atomic radius
Group 1: Alkali metals
Lose 1 electron
Group 2: Alkaline Earth
Lose 2 electrons

5. Review of Nonmetals Gain electrons to form negative ions (anions)
Ionic radius is larger than atomic radius
Group 17: Halogens
Gain 1 electron
Group 16: Chalcogens
Gain 2 electrons

6. Ionic Bonds
The force of attraction that holds ions of opposite charge together
This attraction is between a metal and a nonmetal
Forms from the TRANSFER of electrons
Electrons are transferred from METAL NONMETAL
Large difference in electronegativity
Overall charge of compound is neutral (+ and – cancel out)
Ex) NaCl

7. Animation
s/chang_7e_esp/bom1s2_11.swf

8. Writing Lewis Dot Structures for Ionic Compounds
Lewis Dot diagrams help us visualize what is happening to valence electrons when a bond forms
Remember- Follow the octet rule!!
Ex) Li and F
Steps:
1) Draw all elements and their individual Lewis Dot diagrams.
2) Draw an arrow indicating the transfer of electrons
3) Redraw your bonded compound with appropriate charges

9. More Practice…
Draw the Lewis Dot Diagrams for the following ionic compounds:
BaO
K2S
CaCl2

10. What if you are not given the formula?
Draw the Lewis Dot diagram for the ionic bonding between sodium and sulfur.

11. Ionic Compounds
An ionic compound exists as a collection of positively and negatively charged ions arranged in repeating patterns
A formula unit is the lowest whole-number ratio of ions in an ionic compound
NaCl (1:1 ratio) , MgCl2 (1:2 ratio)

12. Properties of Ionic Compounds (salts)
Ionic bonds are the strongest bonds, so all are solids
Hard
High melting points and boiling points
Soluble in water
Cannot conduct electricity as a solid
Can conduct electricity as a liquid or in an aqueous solution
They are electrolytes- substances that conduct electricity when dissolved in water

13. Covalent Bonds
A bond between two nonmetals that involves a sharing of electrons (“tug of war”)
Can have EQUAL or UNEQUAL sharing
Small or no difference in electronegativity
Overall charge of a covalent compound is neutral
Covalent compounds are called molecules
Ex) O2, CO, CCl4, H2

14. Types of Covalent Bonds
There are 3 types of covalent bonds
Covalent Bonds
Non-Polar Covalent
Polar covalent
Coordinate Covalent

15. Non-Polar Covalent Bond
Electrons are shared equally
Same atoms- same electronegativity
All diatomics have non-polar bonds! Remember: HOFBrINCl
Two of the same non-metal
Ex) Br2

16. Polar Covalent Bond Electrons are shared unequally
Different atoms with different electronegativity values
Ex) HF
The more electronegative atom attracts electrons more strongly and gains a slightly (-) charge
The less electronegative atom has a slightly (+) charge

17. Bond Polarity
Bond Polarity: refers to a separation of charge in a bond
Ex) HF : partial (+) charge on H and partial (-) charge on F
This separation of charge is often called a dipole
***The greater the difference in electronegativity, the greater the polarity.

18. Challenge Question
Which of the following covalent compounds has the greatest degree of polarity? Choice 1) CO Choice 2) HCl Choice 3) NO Choice 4) HBr

19. Drawing Lewis Structures for Covalent Compounds
Remember! After drawing your diagram, all atoms MUST have 8 valence electrons (except for hydrogen, which should have 2 electrons) Ex) F2

20. What if we have more than two non-metals?
Ex) H2O Ex) NH3

21. Lewis Structures Continued…
When elements from Group 14 are involved in a covalent bond, they spread their e- out
Carbon tends to form 4 covalent bonds
Ex) CH4
Ex) CCl4

22. Practice Problem
Draw the Lewis Structure for CH3Br

23. Can you have more than one bond?
Yes! So far all we have seen are single covalent bonds in which one pair of electrons is shared
Double covalent bond: a bond between two atoms where 2 pairs of e- are shared
Ex) O and O
*** Oxygen tends to form 2 bonds!
Triple covalent bond: a bond between two atoms where 3 pairs of e- are shared
Ex) N and N
*** Nitrogen tends to form 3 bonds!

24. More Practice with Lewis Diagrams..
Draw the Lewis diagrams for CO2

25. Challenge Problem
Draw the Lewis Structure for HCN

26. Let’s Summarize
C, H, O and N form how many bonds??

27. Coordinate Covalent Bond
A bond in which both electrons of the shared pair come from the same atom
Ex) NH4+

28. VSEPR Geometry
Lewis Dot diagrams fail to show the 3-dimensional shapes of molecules
VSEPR Theory (Valence-Shell Electron Pair Repulsion Theory)
Repulsion between e- pairs causes molecular shapes to adjust so that valence e- pairs stay as far apart as possible
shapes

29. Molecular Polarity
Molecular polarity is different from bond polarity!!
In a non-polar molecule, electron distribution is even (symmetrical)
In a polar molecule, electron distribution is uneven (asymmetrical)

30. Intermolecular Forces (IMFs)
Only between covalent molecules, never ionic compounds
Weak forces that act between molecules and hold
molecules to each other
****IMFs are not bonds!!!
IMFs occur BETWEEN molecules, bonding occurs WITHIN molecules

31. IMFs vs. Bonds

32. Intermolecular Forces of Attraction
Van der Waals Forces
Dipole interactions
Hydrogen Bonds

33. London Dispersion Forces
London Dispersion forces: weakest of all molecular interactions; caused by temporary shifts in charge
Between nonpolar molecules
london-forces.shtml
The bigger the atom or molecule  the greater the strength of dispersion forces  the higher the BP

34. Think about it… Which has the strongest London dispersion forces? F2
Cl2
Br2 I2

35. Dipole-Dipole Interactions
Dipole interactions: attraction between polar molecules
The positive and negative charges of different molecules attract each other
Ex. HCl

36. Hydrogen bonds: A special case of dipole-dipole interactions
Hydrogen bonds: intermolecular force between the H of one molecule and a highly electronegative atom of another molecule (must be N, O, or F)
Ex. H2O, NH3
***The high b.p. of water is due to hydrogen bonding

37. Molecule-Ion Attraction
Partial charges on a molecule are attracted to ions
This is what happens when NaCl dissolves in water
Picture:
The hydrogens of water align themselves towards the anion and the oxygen align themselves towards the cation
ssolveslike.htm

38. Molecule-Ion Attraction

39. Properties of Covalent Compounds
Soft
Low melting points and boiling points
Cannot conduct electricity in any phase
Generally insoluble in water
Except sugars! (C12H22O11)

40. Network Solids Properties: A special case of covalent bonding
Atoms held together in a very strong covalent network
Ex. Carbon (Diamond) and SiO2
Properties:
Hard
High m.p. and b.p.
Poor conductors

41. Metallic Bonding Holds metals together
Electrons are mobile and move from one atom to another, creating (+) charged metal ions
Charged metal ions are immersed in a “sea of mobile electrons”

42. Bond Energy Ex) F(g) + F(g)  F2 (g) + ENERGY
Bonds do not break and form spontaneously- an energy change is required
When a bond is broken, energy is ABSORBED (required)
Ex) F2 (g) + ENERGY  F (g) + F (g)
The greater the # of bonds between atoms, the more energy you need to break them
When a bond is formed, energy is RELEASED (given off)
Ex) F(g) + F(g)  F2 (g) + ENERGY