1. Unit 4: Ionic, Covalent, and Metallic Bonding
Chemistry Pre-AP
Calallen High School

2. Day 1

4. Bohr Model Drawings-Review
What did we learn about the electrons orbiting the nucleus? Is this model technically correct?
What does each ring represent?
What are the electrons on the outside ring called?

5. Valence Electrons - Review
Valence electrons: found in the outermost energy level (sometimes called a shell)
These electrons are used for bonding
Example: Nitrogen = 1s2 2s2 2p3
Add up the number of e- (superscripts) in the highest energy level
So, nitrogen has = 5 valence electrons

6. Valence Electrons - Review

7. Lewis-Dot Diagrams - Review
Lewis Dot Diagrams are a way to represent the valence electrons in an atom.
Element’s symbol represents the nucleus and inner-level electrons
Dots represent the valence electrons

8. Lewis Dot Diagram Rules - Reveiw
How to draw Lewis Dot Diagrams:
Dots are placed one at a time on the four sides of the symbol, then paired until all valence electrons are used…
Maximum of 8 e- can be around the symbol
d sublevel electrons are not valence electrons – they are in a lower energy level!

9. Lewis-Dot Diagrams - Review

10. IONS Examples of atoms: Na Ca I O
IONS are charged atoms or groups of atoms with a positive or negative charge.
Ions are used to form ionic bonds.
To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! Examples of ions: Na+ Ca+2 I- O-2
Examples of atoms: Na Ca I O

11. Forming Cations & Anions
A CATION forms when an atom loses one or more electrons.
An ANION forms when an atom gains one or more electrons
F + e- --> F-
Mg --> Mg e-

12. F + e- --> F-
Mg --> Mg e-

13. PREDICTING ION CHARGES
In general
metals (Mg) lose electrons ---> cations (+)
nonmetals (F) gain electrons ---> anions (-)
metals
non-metals
metalloids

14. Ions
Previously, we have said that the atomic # indicates the number of protons and electrons…
The word atom implies a neutral charge
Example: Cl has 17 protons and 17 electrons
Ions are charged atoms that have gained or lost electrons; Example: Cl- has 17 protons and 18 electrons

15. Ions Cation – positive ion (+) Anion – negative ion (-)
Loses electrons to become more positive
Example: Be = 1s22s2 → Be2+ = 1s2
Anion – negative ion (-)
Gains electrons to become more negative
Example: F = 1s22s22p5 → F- = 1s22s22p6

16. Practice
Determine if the atom gains or loses electrons and if it is a cation or anion.
Atom
Loose or Gain
Ion
Cation or Anion? Al O P Br
Lose Al cation
Gain O anion
Gain P anion
Gain Br anion

17. Ion Venn Diagram Cation Anion Negative Charge
Positive Charge
Generally formed by metals
Loses Electrons Al Cu Sn
Anion
Negative Charge
Generally formed by non metals
Gains Electrons Br O P
Charged Atom
Used to form Ionic Bonds

18. Ions & Lewis Dot Diagrams
Lewis dot diagrams for ions have brackets around them with the charge denoted as a superscript
Example:
O2- = 1s22s22p6
[ O ]2-

19. Start 3 types of bonds foldable

20. Day 2

21. Bonding Occurs Between:
Ionic Bonding:
Metal + Nonmetal
Covalent Bonding:
Nonmetal + Nonmetal
Metallic Bonding:
Metal + Metal

22. Properties Ionic Bonding Covalent Bonding Metallic Bonding
Exhibit a crystal structure
Exist as solids
Dissolve easily in water
Have high melting and boiling points
Conduct electricity in solutions
Have high electronegativity differences
Covalent Bonding
Exist as gases, liquids, or solids
Do not dissolve easily in water
Have low melting and boiling points
Do not conduct electricity in solutions
Have low electronegativity differences
Metallic Bonding
Malleable
Ductile
Conductors of heat
Conductors of electricity
Must be in the solid state to conduct heat/electricity.
Deform under pressure

23. Crystal Structure of Salts
In a solid ionic compound, the structure is called a crystal lattice and contains a regular, repeating, three-dimensional arrangement of ions.

24. Picture Ionic Bonding Covalent Bonding Metallic Bonding Na Cl
e- are transferred!
e- are shared!
Sea of e-!

25. Electronegativity Difference
Ionic Bonding
If EN difference is greater than 1.7
Covalent Bonding
0.5 – 1.7 = polar covalent bond
< 0.5 = nonpolar covalent bond
Metallic Bonding

26. Bonding & Electronegativity

27. Types of Compounds Ionic Bonding Covalent Bonding Metallic Bonding
Binary Salts (2 elements)
Ternary Salts (3 or more elements)
Salts w/ multiple oxidation #s
Covalent Bonding
Polar covalent compound
Nonpolar covalent compound
Metallic Bonding

28. Examples Ionic Bonding Covalent Bonding Metallic Bonding NaCl NaHCO3
(table salt)
NaHCO3
(sodium bicarbonate)
NaOH
(sodium hydroxide)
FeCl3
Fe(III) chloride
Covalent Bonding
Water
H2O
Carbon dioxide
CO2
Metallic Bonding
Sterling silver jewelry
Brass instruments
Bronze materials
Stainless steel
Metal alloys:
Brass: copper and zinc
Bronze: copper and tin
Stainless Steel: iron and chromium
Sterling Silver: copper and silver

29. Common Ionic Compounds
Sodium bicarbonate (NaHCO3) is baking soda used in food, cleaning products, and antacids
Sodium chloride (NaCl)
is table salt used in food
Potassium nitrate (KNO3) is saltpeter used in fertilizer, gunpowder, and food preservatives
Sodium hydroxide (NaOH) is lye used in soaps and drain cleaner

30. Common Covalent Compounds
Nitrous oxide (N2O) is laughing gas used as an anesthetic and to boost auto engine power
Methane (CH4) is a flammable gas used as fuel and in homes for domestic heating and cooking purposes.
Sugars, like sucrose (C12H22O11) and glucose (C6H12O6), are used in food and energy production
Hydrogen peroxide (H2O2) is used as a bleaching agent, emetic (induces vomiting), and an antiseptic (clean cuts and scrapes)

31. Common Covalent Compounds: CO2
Carbon dioxide (CO2) is a gas used by plants during photosynthesis, produced in respiration; it is a byproduct of burning fossil fuels, and it has many uses as a fire extinguisher, refrigerant (dry ice), and carbonation in drinks

32. Common Covalent Compounds: H2O
Water (H2O) is a liquid that is vital for life. It covers about 71% of earth’s surface, and makes up about 60% of the human body. It is known as the universal solvent, and has many unique properties and uses.

33. Chemical Bonding practice

34. Day 3

35. Start with candy compounds

36. Chemical Bonding
Bonds are the attractive force between atoms or ions in a compound and depend upon:
The electron configuration (involves valence electrons)
Electronegativity
Why do elements bond?
To achieve a stable electron configuration (8 electrons; noble gas configuration)
Octet rule – atoms lose, gain or share electrons to achieve a stable configuration of eight valence electrons
To achieve the lowest possible energy state (lowest potential energy)

37. Formation of Ionic Bonds
Ionic bonds occur when valence electrons are transferred from one atom to another – usually between a metal and a nonmetal
Cations donate electrons
Anions accept electrons
Opposite charges make them attracted (bonded) to each other by electrostatic forces

38. Formation of Covalent Bonds
A Covalent Bond occurs when valence electrons are shared between atoms (usually between nonmetallic elements)
Covalently bonded compound known as a molecule
These shared electrons are part of the valences of all atoms involved (satisfies octet rule)
Since electrons are shared, no charges appear
Many combinations can occur between two nonmetals
Example: carbon and oxygen can form carbon monoxide (CO) and carbon dioxide (CO2)
Carbon Dioxide

39. Single Covalent Bonds
A Single Covalent Bond occurs when atoms of some elements attain a noble gas configuration by sharing one pair of electrons between two atoms
When drawing these molecules, a line can represent a pair of shared electrons
Ex: Methane (CH4)

40. Multiple Covalent Bonds
Sometimes, atoms share more than one pair of electrons between two atoms creating multiple covalent bonds
When drawing these molecules, multiple lines can represent pairs of shared electrons
Ex: diatomic oxygen (O2) and nitrogen (N2)

41. Day 4-6: Ionic Bonding Puzzle
2-3 days

42. Day 7

43. Molecular Shapes Applies to covalent compounds only.
VSEPR steps (valence shell electron pair repulsion):
1) Identify the central atom as the element that can form the most bonds
2) Draw the Lewis dot structure for the molecule
3) Count total # of electron pairs around the central atom
4) Count # of bonding pairs of electrons around the central atom
5) Count # of lone pairs of electrons around the central atom
6) Look at summary chart, identify shape

44. # of e- pairs around central atom
Molecular Shapes
# of e- pairs around central atom
# of bonding pairs of e-
# of lone pairs of e-
Name
Shape 2
linear 3
trigonal planar 4
tetrahedral 1
trigonal pyramidal
angular (bent)

45. Exceptions to Octet Rule
Beryllium
Has 2 valence electrons
Full with 4 valence electrons
Ex: BeI2
Aluminum
3 valence electrons
Full with 6 valence electrons
Ex: AlCl3
Boron
Ex: BH3

46. Molecular Shape Example
What is the shape of water (H2O)?
1) Central atom? oxygen
2) Draw Lewis Dot…
3) # of total electron pairs around central atom? 4
4) # of bonding pairs around central atom? 2
5) # of lone pairs around central atom? 2
6) Refer to chart to identify shape…angular (bent)

47. Day 8: Covalent Bonding Project

48. Day 9 -10: PHET Project

49. End of Unit 4

51. Types of Electron Sharing
Covalent bonds share electrons equally and unequally (“tug of war” with electrons)
Which molecule appears “balanced”?
Which one does not?
H2O O2
Water (H2O)
Oxygen (O2)

52. Unequal Sharing
Unequal sharing of electrons occurs when atoms that share electrons have different electronegativities and are called polar bonds
Shared electrons spend a greater amount of time at the more electronegative atom
Example: water (H2O)

53. Unequal Sharing
Symbols, delta plus and delta minus, represent a partial positive charge and a partial negative charge.

54. Unequal Sharing Example: H2O
Water Molecule Animation

55. Equal Sharing
Equal sharing of electrons occurs in non-polar covalent bonds
No difference (or very, very small difference) in electronegativity between atoms that are sharing
Example: chlorine (Cl2)

56. Polar Molecules
The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape

57. Polar & Nonpolar Molecules
Polar molecules have a partial positive and partial negative charge – called a dipole
Molecules are considered polar if they:
Have an asymmetrical shape
Some reasons for asymmetricl shape include
Has a lone pair of electrons
Has Polar Bonds AND has different elements around the central atom

58. SYMMETRICAL MOLECULES
This molecule is non-polar due to the symmetrical shape of the molecule but it does contain polar bonds. YOU CAN HAVE A NON-POLAR MOLECULE AND POLAR BONDS!!! H
Polar bonds H C H H

59. NON-SYMMETRICAL MOLECULES
The electrons are no longer being pulled with equal force nor in equal directions. H H C Cl H
Different elements

60. WARNING!!!!!!!
Just because you have polar bonds, DOESN’T MEAN YOU ARE A POLAR MOLECULE.

61. Polar Molecule Practice
Is this Polar or Non Polar?
Non-Polar
H Cl + -

62. Polar Molecule Practice
Is this Polar or Non Polar?
Non-Polar
BF3 F B

63. Polar Molecule Practice
Is this Polar or Non Polar?
Polar
net
dipole
moment
H2O H O

64. Polar Molecule Practice
Is this Polar or Non Polar?
Polar
CHCl3 H Cl
net
dipole
moment

65. Polar & Nonpolar Molecules
*IMPORTANT: Polar & nonpolar molecules are different from polar & nonpolar bonds!!!
Non-polar molecules have charges that are evenly distributed, due to shape (ex: any diatomic molecule, gasoline)
Polar molecules have a partial positive and a partial negative charge (ex: water)

66. Be Prepared for Unit 7 Test on 11/18.
End of Unit 7
Be Prepared for Unit 7 Test on 11/18.