1. Unit 3: Chemical Bonding and Nomenclature Part 1

2. States of Matter Solid Liquid Gas Definite shape and volume
Cannot be compressed
Liquid
Definite volume but takes shape of container
Gas
No fixed volume or shape
Uniformly fills container

3. States of Matter

4. Phase Change
Physical state that involves moving from one state to another
Occurs when energy is added or removed

5. Phase Change

6. Phase Change Melting: phase change from solid to liquid
Freezing: phase change from liquid to solid
Temperature where liquid and solid phases coexist at equilibrium is:
Melting point of the solid
Freezing point of the liquid

7. Phase Change Vapour: gas
Vapour pressure: pressure exert by a gas on its container when it is at equilibrium with its condensed phases (solid or liquid)

8. Phase Change
Boiling: the process of molecules in the liquid phase breaking apart from neighbouring molecules to enter the gas phase
Boiling point: temperature when a liquid’s vapour pressure equals the external pressure acting on the liquid surface

9. Phase Change Evaporation: liquid  gas Condensation: gas  liquid
Sublimation: solid  gas
Deposition: gas  solid
Melting: solid  liquid
Freezing: liquid  solid
Boiling: liquid  gas

10. Phase Change and Kinetic Molecular Theory
Temperature is a measure of the average kinetic energy of the molecules of a substance
Adding energy heats up substances
Energy causes more movement of molecules
Altering molecular movement alters the state of substances

11. Phase Change

12. Phase Change
Remember, molecules do not lose their structure when they undergo a phase change
H2O is still H2O
Steam, water, ice
Molecules simply have more space between them

13. Octet Rule
Atoms bond together to obtain a stable electron configuration
Atoms gain, lose, or share electrons until they are surrounded by eight valence electrons
Some elements require 2 valence electrons (not 8)

14. Octet Rule Think of your orbital diagrams
Elements want to look like the closest noble gas
We use Lewis dot diagrams to show valence electrons and help us see how bonding occurs

15. Lewis Dot Diagrams

16. Lewis Dot Diagrams
Remember, electrons repel each other (negative charge)
They don’t want to fill the same orbital if it can be avoided

17. Common Charges

18. Remember Electrons are negative
Gaining electrons makes charge more negative
Losing electrons makes charge more positive
Main group elements are lazy and want to look like the closest noble gas
Metals want to lose electrons
Non-metals want to gain electrons

19. Ionic Vs. Covalent Bonds
In ionic bonding atoms gain and lose electrons
Charge (ions)
In covalent bonding atoms share electrons
No charge (atoms)

20. Ionic Vs. Covalent Bonds
Ionic bonds contain a metal and one or more non-metals
Covalent bonds contain only non-metals

21. Ionic Bonding Atoms form ions by gaining and losing electrons
Ionic bonds form a crystal lattice

22. Ionic Compounds Contain a metal and one or more non-metals
Contain ions (charged atoms) due to transfer of electrons
NO sharing electrons

23. Ionic Compounds
Electrons are transferred from the metal ion to the non-metal ion
Smallest unit is the formula unit
1Na and 1Cl bonded together is one formula unit of NaCl

24. Ionic Compounds Very high melting and boiling points
Crystalline and can be cleaved
Broken along smooth flat surfaces
Brittle
Conduct electricity when dissolved
Break into ions

25. Ionic Compounds Lewis Structures
We begin by drawing the individual atoms involved in the bonds
NaCl

26. Ionic Compounds Lewis Structures
Then we show the transfer of electrons

27. Ionic Compounds Lewis Structures
Our final structure has square brackets and the charge of the ions

28. Ionic Compounds Lewis Structures

29. Binary Ionic Compounds (Type 1 metals) Formula to Name
Type 1 metals only have one possible charge
Metals keep their name from the table
Non-metal is named ending in -ide

30. Binary Ionic Compounds (Type 1 metals) Formula to Name
MgCl2
Magnesium chloride
NaCl
Sodium chloride
AgBr
Silver bromide

31. Binary Ionic Compounds (Type 1 metals) Name to Formula
Put the symbols for each element
Balance charges (criss cross)

32. Binary Ionic Compounds (Type 1 metals) Name to Formula
Cesium bromide
CsBr
Cadmium fluoride
CdF2
Aluminum sulfide
Al2S3
Zinc sulfide
ZnS

33. Roman Numerals
I=1
II=2
III=3
IV=4
V=5
VI=6
VII=7
VIII=8
IX=9
X=10

34. Binary Ionic Compounds (Type 2 metals) Formula to Name
Type 2 metals can have more than one charge
We must tell other people which form we are talking about
Use roman numerals to differentiate metals
Name as you did type 1 compounds
Add roman numeral in brackets after name of the metal

35. Binary Ionic Compounds (Type 2 metals) Formula to Name
AuCl3
Gold (III) chloride
NbN
Niobium (I) nitride
VBr5
Vanadium (V) bromide

36. Binary Ionic Compounds (Type 2 metals) Name to Formula
Important to remember the roman numeral tells you CHARGE not how many atoms
Write symbols for the metal and non-metal
Put roman numeral as charge
Balance charges (criss cross)

37. Binary Ionic Compounds (Type 2 metals) Name to Formula
Iron (II) bromide
FeBr2
Nickel (III) nitride
CuN
Lead (IV) oxide
PbO2

38. Polyatomic Ions In your booklet
Charged chemical species composed of two or more atoms
Act as a unit

39. Polyatomic Ions Name to Formula
Follow the rules for the type of compound you are using
If there are multiples of the polytomic ion, use brackets
Remember they act as a unit

40. Polyatomic Ions Name to Formula
Copper (II) carbonate
CuCO3
Magnesium permanganate
Mg(MnO4)2
Silver phosphate
Ag3PO4

41. Polyatomic Ions Formula to Name
Name by following the rules for the type of compound you are using
Don’t change the name of the polyatomic ion

42. Polyatomic Ions Name to Formula
AgCN
Silver cyanide
Cu3(PO3)2
Copper (II) phosphite
Mn(HCO3)2
Manganese (II) hydrogen carbonate
Manganese (II) bicarbonate

43. Hydrates
Formed by the addition of water or its components to another substance
Substances without water are called anhydrous
Water molecules form lattice around central
compound

44. Prefixes Mono=1 Di=2 Tri=3 Tetra=4 Penta=5 Hexa=6 Hepta=7 Octa=8
Nona=9
Deca=10
Tell you the number of water molecules that are present
Example:
Hexahydrate = 6 water molecules

45. Hydrates Formula to Name
Name the base compound by following rules
Add “hydrate” with the appropriate prefix

46. Hydrates Formula to Name
LiClO4 • 3H2O
Lithium perchlorate trihydrate
NiSO4 • 6H2O
Nickel (II) sulfate hexahydrate

47. Hydrates Name to Formula
Write the formula for the base compound by following previous rules
Separate water molecules from central compound with “•”
Write H2O with appropriate coefficient

48. Hydrates Name to Formula
Copper (II) sulfate pentahydrate
CuSO4 • 5H2O
Magnesium carbonate pentahydrate
MgCO3 • 5H2O

49. Covalent Bonds Bonds made between non-metal atoms
Electron sharing due to similar affinities for electrons
No transfer of electrons
Smallest unit is the molecule
1C and 4H bonded together is one molecule

50. Covalent Compounds Low melting and boiling points
Pliable in solid form
Do not conduct electricity when dissolved
Do not ionize in solution

51. Covalent Compounds and Lewis Structures
We begin by drawing the individual atoms involved in bonding
Atom needing the most electrons goes in the middle

52. Covalent Compounds and Lewis Structures
We circle the electrons that will be shared by the atoms

53. Covalent Compounds and Lewis Structures
Where there are 2 electrons circled by 2 atoms we replace the electrons with a line
Represents a bond

54. Covalent Compounds and Lewis Structures

55. Covalent Compounds and Lewis Structures
When 2 electrons are shared it is a single bond (1 shared pair)
Can have multiple bonds
4 electrons shared (2 shared pairs) = double bond
6 electrons shared (3 shared pairs)= triple bond

56. Covalent Compounds and Lewis Structures
For more complicated compounds:
Add up total valence electrons of bonding atoms
This is the number of electrons we need in our final structure
Draw one bond between the central atom and the other bonding atoms
Each bond counts as using up 2 of the electrons we started with
Draw in the valence electrons on the atoms
Borrow electrons to give the central atom a full octet

57. Covalent Compounds and Lewis Structures
Draw the Lewis Structure of the following:
CO2
HCN
COS

58. Covalent Compounds Formula to Name
Prefixes
Mono=1
Di=2
Tri=3
Tetra=4
Penta=5
Hexa=6
Hepta=7
Octa=8
Nona=9
Deca=10

59. Covalent Compounds Formula to Name
Write the names of the elements present
Change ending to –ide for last element
Add prefixes to match the number of atoms of each element

60. Covalent Compounds Formula to Name
CH4
Carbon tetrahydride CO
Carbon monoxide
N2S3
Dinitrogen trisulfide

61. Covalent Compounds Name to Formula
Write the symbols for each element in the compound
Use the prefixes as the subscripts in the formula

62. Covalent Compounds Name to Formula
NCl3
Nitrogen trichloride
CS2
Carbon disulfide
BrCl
Bromine monochloride