1. Chemical Bonding
Ms. Manning

2. Back to Compounds….. 2 Types: Covalent Compounds
Formed when non-metals bond with other non- metals
Ionic Compounds
Formed when metals bond with non-metals

3. Conductivity in Liquid
Classification
Compound of:
Bonding
Structure
Tm and Tb
Conductivity in Solid
Conductivity in Liquid
Metals
Metallic
Lattice
High
Metals and Non Metals
Ionic
Low
Non Metals
Covalent
Molecular

4. Properties of Metallic compounds
Relatively dense solids (exception Hg)
Good conductors of heat and electricity
Lustrous when clean/ freshly cut
Strong, malleable (can be shaped) and ductile (can be drawn into a wire)
Sonorous: Ringing sound when hit
Relatively high melting and boiling points
Usually form positive ions

5. Properties of Non-Metals
Non-lustrous
Can exist in any state - generally gases at room temperature
Brittle, non-ductile
Poor conductors of heat and electricity
Usually exist as molecules in their elemental form
Low densities, melting and boiling points.
Combine with other nonmetals to form covalent bonds
Generally form negative ions, e.g.  Cl-, SO42-, and N3-

6. Properties of Metalloids
Generally look metallic but are brittle (not malleable or ductile)
Neither good conductors or insulators; instead they are semiconductors.

7. Chemical Bonding
Chemical Bond = the force of attraction holding atoms or ions together
This is how compounds are made!

8. Classifying Compounds
Ionic Compound = a pure substance formed from a metal and a nonmetal
NaCl
CaO
Molecular Compound = a pure substance formed from two or more different nonmetals
SO2
CO2

9. Ionic versus Molecular Compound
Electrical Conductivity = the ability of a material to allow electricity to flow through it
Ionic Compounds  conduct electricity
Molecular Compounds  DO NOT

10. Electrolyte
Electrolyte = a substance that forms a solution that conducts electricity
Ionic compounds  form electrolytic solutions
Molecular compounds  form non- electrolytic solutions

11. Ionic Bonding Ions = atoms that have gained or lost electrons
Ionic Bond = the electrostatic attraction between positive and negative ions in a compound
Metals lose electrons
Non-metals gain electrons
Both form octets = MORE STABLE

12. Ionic Bonding – Bohr Diagrams

13. Lewis Dot Diagrams – Ionic Bonding
KBr
MgCl2

14. Naming Ionic Bonding Ionic Compounds Metal + Non-metal
Metal name  same as on the atom name
Non-Metal  suffix “-ide”
Example:
NaCl = Sodium Chloride
LiF = Lithium Flouride
MgO = Magnesium Oxide

15. Non-metal Suffixes Nitrogen = Nitride Oxygen = Oxide
Fluorine = Fluoride
Phosphorus = Phosphide
Sulfur = Sulfide
Chlorine = Chloride
Selenium = Selenide
Bromine = Bromide
Iodine = Iodide
Text: page 73 Q

16. How Many Atoms in a Molecule?
Diatomic Molecules = a molecule consisting of two atoms of the same or different elements CO
Polyatomic Molecules = a molecule consisting of more than two atoms of the same or different elements
NH3

17. Covalent Bonding
Covalent Bond = the attractive forces between two atoms that results when electrons are shared by the atoms
A simultaneous attraction of two nuclei for a shared pair of electrons
In Lewis Diagrams – the shared pairs of electrons are shown as lines and the lone pairs as dots

18. Octet Rule Still Applies!
The shared pair of electrons is considered to be a pair of electrons that make both atoms have an octet 8 8

19. Lewis Dot Diagrams – Covalent Bonds

20. The Lone Pair
Lone Pair = a pair of valence electrons not involved in bonding

22. Bonding Capacity
Bonding Capacity = the number of electrons lost, gained or shared by an atom when it bonds chemically
Allows us to predict how many bonds an atom can form

23. Bonding Capacity Carbon 4 Nitrogen 5 3 Oxygen 6 2 Halogens 7 1
Atom
# of Valence Electrons
Number of Bonding Electrons
Bonding Capacity
Carbon 4
Nitrogen 5 3
Oxygen 6 2
Halogens 7 1
Hydrogen

24. Choosing the Central Atom for Polyatomic Molecules
The central position…
Is usually occupied by the element with the highest bonding capacity
C and N are often in the central position
The least electronegative atom is usually the central atom
Hydrogen is NEVER the central atom
Oxygen and Halogens are usually not the central atom
Page 79 gives step by step instructions

25. Covalent Bonds = Strong
A large amount of energy is needed to separate the atoms that make up molecules
The stronger the bond the greater the amount of energy needed to break the bond
Single bond = strong
Double bond = stronger
Triple bond = strongest

26. Single, Double & Triple Bonds

27. Polar Covalent Bonds
Polar Covalent Bonds = a covalent bond formed between atoms with significantly different electronegativities; a bond with some ionic characteristics
When electrons are shared between two atoms = covalent bond
In a bond between identical atoms the electrons are shared equally
In a bond between two different atoms the sharing is unequal

28. Non-Polar versus Polar Covalent

30. Comparison…

31. Difference in Electronegativity…
If the electronegativity difference is:
less than 0.2 = bond is pure covalent
is between 0.2 and 1.6 = bond is polar covalent
is greater than 1.7 = bond is ionic

32. Polar Molecules
Polar Molecules = a molecule that is slightly positively charges at one end and slightly negatively charged at the other because of electronegativity differences

33. Types of Forces
Intramolecular Force = the attractive forces between atoms and ions within a compound
Ionic
Polar Covalent
Non-polar Covalent
Intermolecular Force = the attractive force between molecules

34. IntRA versus IntER-molecular Forces

35. Some Intermolecular Forces
3 major types of Intermolecular Forces:
Dipole-dipole forces
London dispersion forces
Hydrogen bonding
The first two are known as van der Waals Forces
London dispersion forces and dipole-dipole forces

36. van der Waals Force
Dipole-dipole force = an attractive force acting between polar molecules
Attraction between oppositely charged ends of polar molecules

37. van der Waals Force
London Dispersion force = an attractive force acting between all molecules including nonpolar molecules
A result of temporary displacements of the electron “cloud” around the atoms in a molecule resulting in a extremely short- lived dipole

38. London Dispersion Force

39. Hydrogen Bonding
Hydrogen Bonding = a relatively strong dipole-dipole force between a positive hydrogen atom of one molecule and a highly elecgtronegative atom (F, O or N) in another molecule

40. Hydrogen Bonding

41. T M