1. Bonding and Naming
Chapter 6

2. Chemical Bonds & Stability
Chemical Bond: force that holds two atoms together
Valence electrons are the electrons involved in bonding
What are the most unreactive elements on the Periodic Table? How many valence electrons do they contain?
Lack of reactivity is a result of having a full set of valence electrons in the outermost orbital that creates a stable atom.
Atoms will lose, gain, or share electrons in order to obtain a full set of valence electrons and achieve stability.

3. Formation of Ions Cation: positively charged ion
Formed by loss of electrons
Typically formed by metals
Charge of ion is equal to the number of valence electrons lost by atom
Example: Ca will lose 2 valence electrons forming an ion with a +2 charge
Ca2+
Example: K will lose 1 valence electron forming an ion with a +1 charge
K1+

4. Formation of Ions Anion: negatively charged ion
Formed by gaining of electrons
Typically formed by nonmetals
Charge of ion is equal to the number of valence electrons gained by atom
Example: Cl will gain 1 electron (for a total of 8 valence electrons) forming an ion with a -1 charge
Cl-
Example: O will gain 2 electrons forming an ion with a -2 charge
O2-

5. Positive and negative ions are attracted to each other
Crystal Lattices
Positive and negative ions are attracted to each other
Charged ions will surround each other and pack together as closely as possible to form an electrically neutral lattice.
Ions cluster in ratios that result in a net charge of zero. (Remember: Matter is neutral so positive & negative charges have to be present in equal amounts.)

6. Types of bonds
3 major types:
Ionic
Covalent
Metallic

7. Ionic Bonds
Form because electrons are transferred from one atom to another.
The transfer of electrons results in formation of anions and cations that are attracted to each other.
Ionic bonds are formed between metals and nonmetals.

8. Properties of Ionic Compounds
High MP and BP
Ionic bonds relatively strong, require large amounts of energy to break apart
Color related to structure
Hard, rigid, and brittle solids
Due to strong attractive forces holding ions in place
Conductivity
Nonconductors of electricity as solids
Conduct electricity when dissolved in H2O (electrolyte)

9. Examples of Ionic Compounds
Table Salt (NaCl)
Road Salt (CaCl2)
Rust (Fe2O3)

10. Covalent bonds
Covalent bonds are formed when atoms share their valence electrons.
Covalent bonds are formed between 2 nonmetals.

11. Properties of Covalent Molecules
Low melting and boiling points
Soft and brittle solids
Do not conduct electricity, even if dissolved in water.

12. Examples of Covalent Molecules
Sugar, C6H12O6
Water, H2O
Isopropyl (Rubbing) alcohol, C3H7OH

13. Metallic Bonds Lattice formed by metal atoms
Do not share electrons or form ions
Outer energy levels of atoms overlap
Electron Sea Model: all metal atoms in metallic solid contribute valence e- to form “sea” of electrons

14. Properties of Metals Melting points vary greatly Malleable and ductile
Generally durable
Good conductors (heat and electricity)
Luster
Hardness and strength (increase w. # of delocalized e-)

15. Writing Ionic Compound Formulas
All ionic compounds have a net charge of zero because matter is neutral.
Use subscripts to indicate the quantity of each ion required to bond to result in a neutral compound.
Example:
Na and Br
Na will form +1 ion to achieve stability
Br will form -1 ion to achieve stability
Opposites attract to form a neutral ionic compound
NaBr

16. Writing Ionic Compound Formulas
Example:
K and O
K will form +1 ion to achieve stability
O will form -2 ion to achieve stability
Opposites attract to form a neutral ionic compound
K2O
Polyatomic Ions: formed from covalently bonded nonmetals by contain extra electrons or fewer than the required number of electrons giving the ion a charge (see list)
To write a formula, put the polyatomic ion in parentheses before using a subscript to tell how many you need.
Examples: Ca(NO3)2 and Na2SO4

17. Rules for Naming Ionic Compounds
There are two general types of compounds
Binary compounds
compounds formed by only two different elements.
There are three different types of binary compounds:
Binary Salts
Binary Acids
Binary Molecular Compounds

18. Rules for Naming Ionic Compounds
Ternary Compounds
Contain a polyatomic ion.
Polyatomic ion
A group of atoms covalently bonded that possess an overall charge.
These ions act as a unit and form ionic compounds.
There are three different types of ternary compounds
Bases
Ternary Acids
Ternary Salts

19. Rules for Naming Ionic Compounds
Binary Salts
Made up of a metal and a nonmetal.
To name this type of compound
Name the first element
Name the second element but change the ending to “ide”
Example:
CaS calcium sulfide
MgCl2 _______________________________

20. Rules for Naming Ionic Compounds
Binary Acids
All acids must begin with hydrogen.
If it is a binary acid then the first element is hydrogen and the second is a nonmetal.
To name a binary acid:
Hydrogen is named “hydro”
The second element is named but the ending is changed to “ic” followed by the word “acid”
Example:
HBr hydrobromic acid
HCl ________________________________

21. Rules for Naming Ionic Compounds
Bases
All bases end with the hydroxide ion (OH-).
To name a base:
name the first ion
follow with hydroxide.
Examples:
Al(OH)3 - aluminum hydroxide
Na(OH) ________________________

22. Rules for Naming Ionic Compounds
Ternary salts
Generally if the formula ends with a polyatomic ion but does not contain hydroxide (a base) and does not begin with hydrogen (an acid) then it must be a salt.
To name a ternary salt:
Name the first element
Name the polyatomic ion
Examples:
Na2(SO4) sodium sulfate
Ca3(PO4) _______________________________

23. Rules for Naming Ionic Compounds
The Stock or IUPAC Method
Transition Metals and a Few Other Metals
Have various charges
When one of these metals is found in a compound you must use a Roman numeral in the name to indicate the charge.
Examples:
FeCl3 - iron (III) chloride
Sn(SO4) - tin (II) sulfate
Cr(PO4) - ____________________________________

24. Important Terminology
Molecules are made up of atoms that are covalently bonded. Therefore, substances that have ionic bonds are not referred to as molecules.
Substances with ionic bonds are made up of ions and the particles are referred to as formula units.

26. Structural Formulas
Steps to determine Lewis structures for covalent molecules:
Electrons that are shared are written between two atoms.
There is always a central atom, unless there are only 2 atoms in the compound.
Single bonds are represented by 2 dots or a single dash; double bonds are represented by 4 dots or 2 dashes; and triple bonds are represented by 6 dots or three dashes.

27. Structural Formulas
Electrons (dots) can be moved around the symbol of element in any order but CANNOT go from one element to another ~ if shared, they must stay between atoms!
Electrons between atoms are counted as belonging to both atoms.
Dot diagram is CORRECT if all atoms have 8 electrons except atoms in Groups 1, 2, and 13.
Atoms from Groups 1, 2, 13, and 17 never have double or triple bonds.

28. Lewis Dot Diagrams for Covalent Compounds - Examples
H2O HCl
BI CCl4

29. Lewis Dot Diagrams for Covalent Compounds - Examples
PO NH4+

31. VSEPR Model
Valence Shell Electron Pair Repulsion model
All electron pairs around the central atom affect the shape of the molecule.
The shape of molecules depends upon the number of both shared and unshared (lone) pairs of electrons around the central atom.
Draw the dot diagram for water and label the shared and unshared electron pairs:

32. VSEPR Model
Unshared pairs take up more room than shared pairs and therefore they are important in predicting the shape of the molecule.
The shape of a molecule is determined by drawing lines connecting the centers of the atoms in the molecule.
There are six (6) shapes that you have to learn.

33. Molecular shapes 2 – atom linear 3 – atom linear
Trigonal planar Tetrahedral
Trigonal pyramid Bent

34. Steps for determining shape
Draw the Lewis Dot Diagram.
Count up the number of shared and unshared electron pairs around the central atom.
Memorize the following chart…

35. Example
Electron Dot Diagram
Shared Pairs
Unshared (Lone) Pairs
Molecular Shape
Structural Formula
HCl
BeI2
AlBr3
CH4
NH3
H2O

37. Rules for Naming Binary Molecular Compounds
Formed between two nonmetals.
Since nonmetals can form many different compounds with each other, Greek prefixes are used to show how many atoms of each element are present in the compound.
Greek Prefixes:
mono - one hexa - six
di - two hepta - seven
tri - three octa - eight
tetra - four nona - nine
penta - five deca – ten

38. Rules for Naming Binary Molecular Compounds
Rules similar to those for binary salts EXCEPT
If there is more than one atom in the first element, then use the appropriate prefix in front of the name of that element.
The second element always has a prefix (use mono if there is only one) and change the ending to “ide.”
Examples:
CO carbon monoxide
P2O5 ________________________________