1. Ch. 6 The Structure of Matter
The Importance of BONDING

5. Important Terms
Element = a pure substance that cannot be separated or broken down into simpler substances by chemical means
Atom = the smallest unit of an element that maintains the chemical properties of that element
Compound = a substance made up of atoms of two or more different elements joined by chemical bonds

6. Bonding Atoms with unfilled valence shells are considered unstable.
Atoms will try to fill their outer shells by bonding with other atoms.
Chemical bond = the attractive force that holds atoms or ions together in a compound

7. Atomic Bonds Atoms form atomic bonds to become more stable.
Atoms become more stable by filling their valence shell or at least meeting the octet rule by getting 8 valence electrons.
Exception to Octet Rule of 8 valence electrons: Helium—which only has 1 energy level and holds a max. of 2 electrons

8. Atomic Bonds
There are three main types of chemical bonds used by atoms to fill their valence shell:
Covalent
Metallic
Ionic
“Bond,
Chemical Bond”

9. Chemical Formulas A chemical formula tells us:
the type of atoms present
the number of atoms present
the type of compound

10. Chemical Formulas Example: table salt: Sodium Chloride
NaCl
Count the atoms present:
1 Na atom
1 Cl atom

11. Chemical Formulas Sometimes there are subscripts present.
A subscript is a small number that is in a chemical formula. If no subscript is present assume that it is 1.
Example - water: H2O
2 H atoms
1 O atom
Subscript

12. Chemical Formulas
Sometimes there are parentheses with a subscript. The subscript only applies to the atoms within the parentheses.
Example - calcium hydroxide (kidney stones): Ca(OH)2.
1 Ca atom
2 O atoms
2 H atoms

13. Chemical Formulas
Sometimes there are subscripts in the parentheses. Multiply the subscript outside the parentheses by the subscript of each element within the parentheses.
Example - calcium nitrate: Ca(NO3)2
1 Ca atom
2 N atoms
6 O atoms (3 oxygens x 2 = 6)

14. Covalent Bonds
Covalent bonds form between two non-metals. Groups on the Periodic Table
Covalent bonds are formed when atoms SHARE electrons.
Both atoms need to gain electrons to become stable, so they share the electrons they have.
Atoms can share more than one pair of electrons to create double and triple bonds.

15. Properties of Covalent Compounds
Results in a NEUTRAL molecule
Weak bonds
Physical State usually liquids or gases
Low Melting and Boiling Points
Poor conductors of electricity (no free electrons to move around)

16. Each chlorine atom wants to gain one electron to achieve an octet.
Covalent Bonds
Use Lewis structures to draw valence electrons for each atom in the covalent pair. Cl Cl
Each chlorine atom wants to gain one electron to achieve an octet.

17. Cl Cl octet octet Covalent Bonds
The octet is achieved by each atom sharing the electron pair in the middle. Cl Cl
octet
octet
Now, each Chlorine atom has 8 valence electrons because it is sharing one pair.

18. Cl Cl Cl Cl Cl2 Chlorine Molecule
It is a single bonding pair so it is called a single covalent bond. The compound is now called a molecule. Cl Cl
Cl Cl Cl2

19. Covalent Bonds O O
How will oxygen bond?

20. O O Covalent Bonds O2 Two bonding pairs, making a double bond.
The double bond can be shown as two dashes O O O2

21. Covalent Bonds
Elements can share up to three pairs of electrons (6 total electrons).
Single Bond (2e)
Double Bond (4e)
Triple Bond (6e)

22. Covalent Bonds Atoms can share their electrons equally or unequally.
When atoms share electrons equally, it is called a non-polar covalent bond.
Non-polar covalent bonds form between atoms of the same type. Ex: H2, Cl2,
When atoms share electrons unequally it is called a polar covalent bond.
One atom pulls the electrons closer to itself.
The atom that pulls the electrons more gets a slightly negative charge.
The other atom gets a slightly positive charge.
Ex: Water molecule
Bonding Animation

23. Covalent Bonds Nomenclature
Naming binary covalent compounds:
Two nonmetals
Name each element
Change the ending of the 2nd element to –ide
Use prefixes to indicate the # of atoms of each element
Do not use “mono” with the first element
# of Atoms
Prefix 1
mono- 2
di- 3
tri- 4
tetra- 5
penta- 6
hexa- 7
hepta- 8
octa- 9
nona- 10
deca-

24. Covalent Bonds Nomenclature
carbon monoxide
CO2
carbon dioxide
PCl3
phosphorus trichloride
CCl4
carbon tetrachloride
N2O
dinitrogen monoxide
# of Atoms
Prefix 1
mono- 2
di- 3
tri- 4
tetra- 5
penta- 6
hexa- 7
hepta- 8
octa- 9
nona- 10
deca-

25. Covalent Bonds Nomenclature
Given the following covalent compounds, WRITE the correct chemical formula.
Name
Chemical Formula
Hydrogen Disulfide
Diphosphorus pentoxide
Trinitrogen hexafluoride
HS2
P2O5
N3F6

26. Practice: Drawing Covalent Bonds
We can illustrate covalent bonding using Lewis structures.
1 – Draw a Lewis structure for each element.
Ex: C H
2 - Continue adding atoms until all atoms have a full valence H
H C H
CH4
carbon tetrahydride

27. Ions Ions are formed when atoms gain or lose electrons.
Ions are charged atoms (positive or negative).
Positive ions are called cations.
Formed when the atom loses electrons.
Lose negative charge, becomes positive ION
Metals
Negative ions are call anions.
Formed when the atom gains electrons.
Gain negative charge, become negative ION
Non-metals

28. Ionic Bonds Ionic bonds are formed between metals and non-metals.
Ionic bonds are formed between oppositely charged atoms (ions).
Ionic bonds are formed by the transfer of electrons.
One atom loses (gives away) electrons.
One atom gains (receives) electrons.

29. Ionic Bonds
Use the number of valence electrons to determine the # of electrons that are lost or needing to be gained.
The transfer of electrons create a positive ion and a negative ion. The opposite charges attract one another, causing a chemical bond to form.
Bonding Animation

30. Atoms with 4 or less valence electrons want to LOSE (give away) their valence electrons.
[Groups 1, 2, 13, 14]
Atoms with 4 or more valence electrons want to GAIN (receive) more electrons to satisfy their octet. [Groups 14, 15, 16, 17]

31. Ionic Bonds
The normal charge of an ion can be quickly determined using the oxidation number of an element.
The oxidation number of an atom is the charge that atom would have if the compound was composed of ions.

32. Ionic Bonds To find the oxidation number : Look at Group #
Determine # of valence electrons
If 4 or less, atom will lose (give away) valence electrons (ion is positive)
If 4 or more, atom will gain the needed # to fill valence shell. (ion is negative)

33. Ionic Bonds Example: Beryllium is in Group 2 Be has 2 e-
Wants to achieve octet
Loses the 2 e-
Oxidation #/Ion charge of +2
Nitrogen is in Group 15
N has 5 e-
Needs 3 more for octet
Gains 3 e-
Oxidation #/Ion charge of -3

34. Practice: Determining Oxidation Numbers
Atom
Group
Valence Electrons
Oxidation Number
Oxygen
Calcium
Fluorine
Phosphorus
Sodium 16 6 -2 +2 2 2 -1 17 7 15 5 -3 1 1 +1

35. Ionic Bonding Nomenclature
To name Binary Ionic Compounds:
2 elements—one METAL and one NON-METAL
Cation is always written first [Metal]
Cation name stays the same
Anion is written second [Non-metal]
Change the non-metal’s ending to “-ide”.
NO PREFIXES ARE USED FOR IONIC COMPOUND NAMING

36. Name the nonmetal ion, changing the suffix to –ide.
Examples
NaCl
Name the metal ion
Sodium
Chloride
CaO
Name the nonmetal ion, changing the suffix to –ide.
Calcium
Oxide
Al2S3
Aluminum
Sulfide
MgI2
Magnesium
Iodide
BaNa2
You should recognize a problem with this one
The name of this is Banana (JOKE – haha)
This is two metals – not a binary ionic compound

37. Drawing Ionic Bonds 1 – Draw the Lewis structure for each element.
Ex: Na Cl
2 – Draw arrows to show the TRANSFER (gain/loss) of electrons [draw extra atoms if needed]

38. Drawing Ionic Bonds (continued)
3 – Draw ion Lewis diagrams showing the new charge for each ion.
Ex:
4- Write the chemical formula for the compound formed represents the ratio of negative ions to positive ions.
Ex: NaCl – for every 1 sodium ion, there is also 1 chlorine ion.
Chemical Formula = NaCl

39. Practice Drawing Ionic Bonds
Elements Lewis Transfer Formula
Diagram
Calcium
Fluorine
Sodium
Oxygen

40. Be F BeF2 2+ 1- “Swap & Drop” Method beryllium fluoride
Given the name of an Ionic Compound, you can determine the chemical formula using the “swap and drop” method:
Write the symbols for each ion.
Determine the oxidation number of each ion.
Swap and Drop
Reduce (if necessary).
Rewrite
Be F
beryllium fluoride
BeF2

41. Ionic vs. Covalent Bonds in Binary Compounds
Ionic Bonds
Form when electrons are transferred between atoms.
Form between a metal and a non-metal.
Covalent Bonds
Form when electrons are shared between atoms.
Form between two non-metals.
Both types of bonds result in all atoms having a full outer energy level.

42. Ionic vs. Covalent Bonds in Binary Compounds
Other comparisons between Ionic and Covalent Compounds:
Ionic Compounds
Results in a Neutral Compound
Crystalline Solid
Strong Bonds
High Melting Point
Covalent Compounds
Results in a Neutral Molecule
Mostly results in gases or liquids
Weak Bonds
Low Melting Points

43. Polyatomic Ions
A polyatomic ion is a group of covalently bonded atoms that have lost or gained an electron. (Example: Nitrate NO3- and Ammonium NH4+).
Oppositely charged polyatomic ions can form compounds. (Example: Ammonium nitrate NH4NO3).

44. Common Polyatomic Ions
Naming of these compounds follows the same rules as binary ionic compounds.
The most important part is recognizing there is a polyatomic ion present.
Common Polyatomic Ions
ammonium
NH4+
carbonate
CO32-
bicarbonate
HCO3-
hydroxide
OH-
nitrate
NO3-
nitrite
NO2-
phosphate
PO43-
sulfate
SO42-
sulfite
SO32-
acetate
C2H3O2-

45. Practice: Polyatomic Ions
To go from the formula to the name:
Name the cation.
Name the anion.
Be(NO3)2
(NH4)2S
beryllium
ammonium
nitrate
sulfide

46. O2- (NH4)2O Polyatomic Ions ammonium oxide To go from name to formula:
Write the symbols for each ion.
Determine the oxidation number of each ion.
Swap and Drop
Reduce (if necessary).
Put parentheses around the polyatomic ion if receives a subscript greater than one.
Rewrite
ammonium oxide
NH41+
O2-
(NH4)2O
** Remember charges CANCEL out each other!!

47. Practice: Polyatomic Ions
Compound Name
Oxidation #s
Chemical Formula
Calcium phosphate
Sodium hydroxide
Ammonium sulfate
Ca PO43-
Ca3(PO4)2
Na OH1-
NaOH
(NH4)2SO4
NH SO42-

48. Metallic Bonds
Metallic bonds are metal to metal bonds formed by the attraction between positively charged metal ions and the electrons around them.
Atoms are packed tightly together to the point where outermost energy levels overlap.
This allows electrons to move freely from one atom to the next making them great conductors of electricity.

49. Transition Metals--Ionic Compounds
Transition metals are cations that have variable charges that makes them hard to name.
We use Roman numerals to indicate the charge of a transition metal.
Example:
copper (II) oxide – charge of copper for this compound is +2
titanium (IV) sulfide – charge of titanium for this compound is +4

50. Transition Metal Ionic Compounds2+ 2-
To go from formula to name you need to determine the Roman numeral for your transition metal.
If there are no subscripts, simply give the transition metal the equal and opposite charge to the nonmetal.
Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.
Fe O
iron (II) oxide

51. Transition Metal Ionic Compounds
To go from formula to name you need to determine the Roman numeral for your transition metal:
If there are subscripts present use the reverse “Swap and Drop.”
Now use normal ionic bonding rules putting your new number in Roman numerals to the right of your transition metal ONLY.
Fe2O3
iron (III) oxide