1. Valence electrons
Electrons in the highest occupied energy level of an atom.
These are the electrons that determine the element’s properties.

2. Electron dot structures
Diagrams that show the valence electrons as dots.
The core electrons and the nucleus are included in the symbol of the element

3. Octet Rule
In forming compounds, atoms tend to achieve noble gas configuration.
8 electrons in the highest energy level.

4. IONS Atoms or groups of atoms that have a positive or negative charge.
Cations - positive ion resulting from loss of electrons.
Anions - negative ions resulting from gain of electrons.

5. Formula Unit Represents an ionic compound.
Lowest whole number ratio of ions in the compound.

6. Formula units
Ionic compounds form as repeating links in a crystal matrix.
Each cation is bound to each neighboring anion
The ions are “locked in place”
NaCl is the smallest ratio that indicates this matrix

7. Ionic Bond
Bonds resulting from the electrostatic attraction between oppositely charged ions. In an ionic compound the net ionic charge is 0.

8. Ionic Compounds Metal + Non-metal Polyatomic cation+Non-metal
Metal+ Polyatomic anion
Solid at room temperature
High melting point >300°C

9. Ionic compounds NaCl Na2SO4
CaCO3
Crystalline solids that have high melting points.
They are often soluble in water
They conduct electricity when in solution, or when molten.

10. Predicting the formula
Identify the charge of the cation
Identify the charge of the anion
Make a “T” table
Add ions until the positive charge equals the negative charge

11. Predicting the formula
Polyatomic ions are groups of atoms that stay together, they are treated like “super atoms” The entire group carries the charge.
Table 9.3 p.257
NH4+

12. Lets practice !!!! Potassium + Phosphorus Lithium + Selenium
Aluminum + Chlorine
Gallium + Sulfur
Magnesium + Iodine
Sodium + Carbonate
Sodium + Hydrogen Carbonate
Strontium + Phosphate
Ammonium + Chromate
Barium + Acetate

13. Lets practice !!!! Potassium + Phosphorus K3P Lithium + Selenium Li2Se
Aluminum + Chlorine
AlCl3
Gallium + Sulfur
Ga2S3
Magnesium + Iodine
MgI2
Sodium + Carbonate
Na2CO3
Sodium + Hydrogen Carbonate
NaHCO3
Strontium + Phosphate
Sr3(PO4)2
Ammonium + Chromate
(NH4)2CrO4
Barium + Acetate
Ba(C2H3O2)2

14. Lets Review!!! calcium chloride cesium oxide aluminum perchlorate
barium sulfide
sodium dichromate
aluminum phosphate
calcium carbonate
sodium carbonate

15. Lets Review!!! calcium chloride CaCl2 cesium oxide Cs2O
aluminum perchlorate
Al (ClO4)3
barium sulfide
BaS
sodium dichromate
Na2Cr2O7
aluminum phosphate
Al PO4
calcium carbonate
CaCO3
sodium carbonate
Na2CO3

16. Compounds with transition metals
Transition metals can have more than one charge.
You may have more than one possible compound: FeO, or Fe2O3
Make tables & work backwards to determine cation charge
Indicate charge with a roman numeral
REMEMBER THE TABLE MUST BE BALANCED !!!!!
Fe+?
O-2
iron(II) oxide
Fe+?
O-2
iron(III) oxide

17. Some Ions we need to just Know
Silver is always +1
Zinc is always +2
Cadmium is always +2
Do not use a roman numeral with these
Iron may be +2 or +3
Tin may be +2 or +4
Lead may be +2 or +4
More in table 9.2 p.255

18. Lets Review!!! potassium oxide strontium nitride strontium nitrate
strontium nitrite
aluminum hydroxide
magnesium sulfate
iron(III) oxide
silver oxide

19. Lets Review!!! Potassium Oxide K2O Strontium Nitride Sr3N2
Strontium Nitrate
Sr(NO3)2
Strontium Nitrite
Sr(NO2)2
Aluminum Hydroxide
Al(OH)3
Magnesium Sulfate
MgSO4
Iron(III) Oxide
Fe2O3
Silver Oxide
Ag2O

20. Lets practice!!
Na2S
Hg2S
Na2Cr2O7
Hg2Cr2O7
CuO

21. Lets Practice Answers Na2S sodium sulfide Hg2S mercury(I) sulfide
Na2Cr2O7
sodium dichromate
Hg2Cr2O7
mercury(I) dichromate
CuO
copper(II) oxide

22. Lets Practice calcium carbonate ammonium sulfate copper(I) phosphate
chromium(IV) acetate
cadmium perchlorate

23. Lets Practice calcium carbonate CaCO3 ammonium sulfate (NH4)2SO4
copper(I) phosphate
Cu3PO4
chromium(IV) acetate
Cr(C2H3O2)4
cadmium perchlorate
Cd(ClO4)2

24. Links to practice tests and games
Interactive link
Interactive link 2
Interactive link 3
Interactive link 4

25. Metallic Bonds
The force of attraction that holds metals together. The attraction of the free floating electrons for the positively charged metal ions

26. Metallic Properties
Malleable
Ductile
Conduct heat and electricity.

27. Single covalent bond
A bond in which two atoms share a pair of electrons between them in order to achieve noble gas configuration.

28. Structural formulas
Chemical formulas that show the arrangement of atoms in molecules and polyatomic ions. Each dash represents a pair of shared electrons.

29. Unshared pairs
Pairs of valence electrons that are not involved in bonding, not shared between atoms.
Also called lone pairs or non-bonding pairs

30. Double covalent bond
Two atoms share two pairs of electrons between them to attain noble gas configuration.
O2 and CO2

31. Triple covalent bond
Two atoms share three pairs of electrons between them to attain noble gas configuration. N2

32. Coordinate covalent bond
A covalent bond in which one atom contributes both bonding electrons.
CO and NH4+ and N2O

33. Exceptions to the octet rule
NO2
BF3
PCl5
SF6

34. Law of Definite Proportions
In samples of any chemical compound, the masses of the elements are always in the same proportions.

35. Law of Multiple Proportions
When two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.

36. Isoelectronic ions Ions containing the same numbers of electrons.
Generally for isoelectronic ions size decreases as nuclear charge increases.

37. Bond energy The energy required to break a bond. Table 8.4 p.365
Bond length- the distance where energy is minimum.Table 8.5 p.365

38. Coulomb’s Law E=2.31x10-19Jxnm Q1Q2 r E= energy in joules
r= distance between ion centers in nm
Q1&Q2= ion charges

39. Molecular Compounds All non-metals covalently bonded.
Solid, liquid or gas
Low melting point <300°C
Smallest representative particle is a molecule.

40. Non-polar covalent bond
A covalent bond in which the electrons are shared equally. The two atoms have nearly the same electronegativities

41. Polar covalent bond
A covalent bond in which the electrons are not shared equally. The more electronegative atom will pull more of the electrons toward itself.

42. Polar molecule
One end of the molecule has a slightly positive charge and one has a slightly negative charge.
This is called a dipole.
Depends on the shape.

43. Lattice energy Lattice Energy=k(Q1Q2) r
The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.
Lattice Energy=k(Q1Q2) r

44. Use the following to calculate DH°f of BaCl2(s).
Lattice energy= kJ/mol
1st ionization Ba= 503kJ/mol
2nd ionization Ba= 965kJ/mol
Electron affinity Cl=-348kJ/mol
Bond energy Cl2=239kJ/mol
DH sublimation Ba=178kJ/mol

45. Bond Energies & Enthalpy
DH=SD(bonds broken)-
SD(bonds formed)
S=sum of terms
D=bond energy per mol of bonds, always positive.

46. Localized Electron Bonding Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Lewis Structures
VSEPR Theory

47. Writing Lewis Structures
Sum the valence electrons from all the atoms.
Use a line to show a pair of electrons between each pair of bound atoms (Bonding Pairs)

48. Writing Lewis Structures
Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the 2nd row elements. (Lone Pairs)
Double or triple bonds may be needed.

49. Comments on Octet Rule
C,N,O,F obey octet rule.
B and Be often have less than 8 electrons. Very reactive.
2nd row never exceed rule.
3rd row and up often obey octet rule but may exceed it., due to d orbitals.

50. When writing lewis structures satisfy the octet rule for the atoms first. Place any remaining electrons on the elements that have available d orbitals.

51. Resonance structures
Occur when it is possible to have two or more valid electron dot structures for the same molecule or ion.
SO3, SO2

52. Formal Charge (FC)
A method to decide which of many possible non-equivalent Lewis structures is most likely to occur.
Atoms in molecules try to achieve FC as close to 0 as possible.

53. FC=(# valence e- on free atom) -(# valence e- assigned to the atom in the molecule).
(Valence e-)assigned = (# lone pair electrons) + 1/2(#shared electrons)

54. VSEPR theory Valence-shell-electron-pair repulsion theory.
Because electron pairs repel molecular shape adjusts so the valence electron pairs are as far apart as possible.

55. Hybrid orbitals
In hybridization several atomic orbitals mix to form the same number of equivalent hybrid orbitals

56. Sigma bonds
Formed along the axis that joins the atomic nuclei when two atomic nuclei combine to form a molecular orbital.

57. Pi bond
Electron in pi bonds are found in sausage shaped regions above and below, or in front and behind the bond axis.

58. Paramagnetic molecules
Show an attraction to an external magnetic field.
Molecules contain one or more unpaired electrons.

59. Diamagnetic Molecules
Molecule is repelled by an external magnetic force.
Associated with paired electrons

60. Electronegativity
The ability of an atom in a compound to draw electrons to itself.
Pauling electonegativity values Table 14.2 p.405
Large electronegativity differences correspond to ionic bonds