1. Unit 7: Bonding

2. Why Bond?
Elements bond in order to get a stable valence e- configuration (a full valence shell)
Noble gases have stable valence e- configurations and tend not to bond
Bonding is all about becoming more stable (at the lowest energy possible)

3. Breaking/Forming of Bonds
When a bond is broken, energy is ABSORBED (required)
Ex) F2 (g) + ENERGY  F (g) + F (g)
When a bond is formed, energy is RELEASED (given off)
Ex) F(g) + F(g)  F2 (g) + ENERGY

4. Types of Chemical Bonds
Chemical bonds are formed when valence electrons are:
Transferred from one atom to another (ionic)
Shared between atoms (covalent)
Mobile within a metal (metallic)

6. Don’t forget about electronegativity!
The ability of an atom to attract electrons when in a compound
F- 4.0 (highest EN) and Fr-0.7 (lowest EN)
Noble gases- no assigned value
EN values can be used to determine the type of bond that forms between two atoms
Bigger difference in EN ( difference >1.7 )  more ionic character

7. Ionic- Metal to Non-Metal
An ionic bond is the force of attraction that holds ions of opposite charge together, forming an ionic compound
***Electrons are transferred from M NM
***Metal loses electrons (becomes positive) and the non-metal gains electrons (becomes negative)
Large difference in electronegativity between the metal and non-metal
Overall charge of compound is neutral

8. Writing Lewis Dot Structures for Ionic Compounds
sodium and chlorine
barium and oxygen
calcium and fluorine
potassium and sulfur

9. 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 is neutral
Covalent compounds are called molecules

10. Types of Covalent Bonds
Non-Polar Covalent
Equal sharing of electrons
Polar covalent
Unequal sharing of electrons
Coordinate Covalent
One nonmetal provides BOTH e- to be shared

11. Non-Polar Covalent Bond
Equal sharing of e-
Look for two of the same nonmetals
All diatomics have non-polar bonds! Remember: H2 O2 F2 Br2 I2 N2 Cl2
Ex) F2

12. Polar Covalent Bond Unequal sharing of electrons
Different NMs with different electronegativity values
Ex) HF
The more electronegative atom has a slight (-) charge
The less electronegative atom has a slight (+) charge

14. Bond Polarity
Bond Polarity: refers to a separation of charge in a bond
Ex) HCl : partial (+) charge on H and partial (-) charge on Cl
This separation of charge is called a dipole
***The greater the difference in electronegativity, the greater the polarity.
Ex) HF is more polar than HBr

15. Formation of Covalent Bonds

16. Formation of Covalent Bonds
Bond length: the average distance between two bonded atoms
Bond energy: the energy required to break a chemical bond and form neutral atoms (kJ/mol)
In general, bond energies become larger as bond lengths become shorter

17. Drawing Lewis Structures for Covalent Compounds
After drawing your diagram, all atoms MUST have 8 valence e- (except for hydrogen, which should have 2 valence e-)
When elements from Group 14 are involved in a covalent bond, they spread their e- out
Carbon tends to form 4 covalent bonds
Atoms of carbon can bond to each other to form chains, rings and networks

18. Lewis Structures may contain…
Single covalent bond: one pair of e- is shared between two atoms (2 total e-)
Double covalent bond: 2 pairs of e- are shared between atoms (4 total e-)
Triple covalent bond: 3 pairs of e- are shared between atoms (6 total e-)

19. Coordinate Covalent Bond
The third type of covalent bond…
One non-metal donates BOTH electrons to be shared
Examples: NH3 and H2O

20. PROPERTIES OF IONIC AND COVALENT COMPOUNDS

21. 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)

22. Properties of Ionic Compounds (salts)
Ionic bonds are very strong bonds, so all are solids (hard with a crystalline structure)
High melting points and boiling points
Soluble (can dissolve) in water
Cannot conduct electricity as a solid (ions are not free to move)
Can conduct electricity as a liquid or in an aqueous solution (ions are free to move)
Electrolytes- substances that conduct electricity when dissolved in water

23. Properties of Covalent Compounds (Molecules)
Many are gases or liquids at room temperature
Molecular solids tend to be soft
Low melting points and boiling points
Cannot conduct electricity in any phase
Generally insoluble in water
Except sugars! (C12H22O11)

24. Network Covalent Solids
A special case of covalent bonding
All atoms are held together in a very strong covalent network
Ex. Carbon (Diamond) and SiO2
Properties:
Very hard
High m.p. and b.p.
Poor conductors

25. Metallic Bonding: The “sea of mobile electrons”

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

27. IMFs vs. Bonds

28. LD Forces

29. Dipole-Dipole Attractions

30. Hydrogen bonding

31. 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

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

33. 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