1. 8.1 – Forming Chemical Bonds
Objectives:
Describe how and why chemical bonds form.
Use the octet rule to predict ionization.
Describe how cations and anions form and from what elements they tend to form.

3. 8.1 – Chemical Bonds Chemical bond Attraction between atoms
Based on electrostatic force (+ attracts -)
Occur so that atoms can become more stable (lose energy)
Strong Bonds (occur between atoms)
Ionic bonds – attraction between positive & negative ions
Covalent bonds – attraction between nucleus & electrons of other atoms
Metallic bonding – attraction of metallic nuclei to delocalized electrons in a metallic matrix
Weak Bonds (occur between molecules)
Hydrogen bonding
Dipole-dipole attractions
London dispersion forces

4. 6.2 – Organizing the Elements
Valence Electrons
Electrons in the outermost energy level of an atom
Determine the chemical properties of the element
Elements are grouped on the periodic table based on having the same number of valence electrons
There may be no more than 8 valence electrons for an atom.
Li 1s22s1
Na 1s22s22p63s1
K 1s22s22p63s23p64s1
Rb 1s22s22p63s23p63d104s24p65s1
Cs 1s22s22p63s23p63d104s24p64d105s25p66s1

5. 6.2 – Valence Electrons & the P.T.
Valence Electrons & Period
The period number of an element indicates the energy level of the valence electrons for that element.
Valence Electrons and Group Number
The Roman numerals on the periodic table for the main-group elements show the number of valence electrons available for bonding.
The number of valence electrons for the transition elements is technically 2, but in reality the electrons from the lower s sublevel mingle with the d sublevel electrons to create a variety of bonding possibilities.

6. 8.1 – The Octet Rule
Atoms tend to gain, lose or share electrons in order to acquire a full set of valence electrons.
If an atom has a full set of valence electrons its electron configuration is the same as that of a noble gas and is very stable.
Less stable arrangements may occur where individual sub-levels are filled, or larger sub-levels are half-filled.
Example:
Iron (Fe)
Iron III (Fe3+)
Ne He
F- Be2+ : 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d   1s 2s 2p 3s 3p 4s 3d 4p 5s 4d  

7. 8.1 – The Octet Rule
Pseudo-noble gas configurations occur when a d sub-level is filled and the higher energy level electrons are lost (as in Cu+ or Zn2+)
Examples:
Copper (Cu)
Copper I (Cu+)
Zinc (Zn)
Zinc (Zn2+) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d   1s 2s 2p 3s 3p 4s 3d 4p 5s 4d  1s 2s 2p 3s 3p 4s 3d 4p 5s 4d  1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 

8. 8.1 - Ions Positive Ions – Cations
Form when an atom loses one or more electrons to get a full octet
Occurs in metals because they have weaker nucleus-electron attractions
Reactivity depends on how easily an atom loses electrons

9. 8.1 - Ions Negative Ions – Anions
Form when an atom gains one or more electrons to get a full octet
Occurs in non-metals because they have stronger nucleus-electron attractions
To indicate an atom has gained electrons, the ending of the element’s name is changed to –ide
Reactivity depends on how easily an atom gains electrons
Some non-metals may gain or lose electrons depending on whether the other elements present are stronger or weaker than them.

10. 8.2 – Nature & Formation of Ionic Bonds
Objectives:
Describe an ionic bond.
State the properties of ionic compounds.
Relate the ionic structure to the properties of ionic compounds.

11. 8.2 – Formation of Ionic Bonds
Electrostatic attraction between oppositely charged particles (anions & cations)
An ionic compound is a substance composed of positive & negative ions. (This is not by definition a molecule!)
In an ionic compound, the charges of the cations must equal the charges of the anions.
Binary ionic compounds – 1 metallic cation & 1 non-metallic anion

12. 8.2 - Properties of Ionic Compounds
Ionic compounds form crystalline patterns in the solid state with anions and cations arranging themselves in a repeating geometric pattern.
Unlike molecules, there is no single ionic unit formed.

13. 8.2 – Properties of Ionic Bonds
Properties due to crystalline structure
High melting point (typically solid at room temperature)
High boiling point
Hard/Brittle
Conduct electricity in liquid but not solid state.
Molten ionic compounds usually conduct electricity well.
Solid ionic compounds make good insulators.
Electrolytes - soluble in water to form electrically conductive solutions

14. 8.2 – Energy in Ionic Bonds
Energy is absorbed or released in all chemical reactions
Energy absorbed – endothermic
Energy released – exothermic
When ionic bonds forms, the process is always exothermic
When the atoms exchange electrons, they become more stable
Lattice energy – energy required to separate one mole of ions of an ionic compound
Relates to size of ions (smaller are stronger)
Relates to charge (higher is stronger)

15. Magnesium oxide MgO Potassium oxide Lithium hydroxide K2O LiOH FeCl2
8.3 – Names & Formulas of Ionic Compounds
Magnesium oxide
MgO
Objectives:
Write correct formulas for ionic compounds.
Differentiate between monatomic and polyatomic ions.
Name ionic formulas correctly.
Potassium oxide
K2O
Lithium hydroxide
LiOH
FeCl2
Iron (II) chloride

16. 8.3 – Formulas for Ionic Compounds
There are no single units of an ionic compound
Formula units – the simplest ratio of ions in an ionic compound
Binary compounds – compounds made up of only 2 elements
The number of electrons gained by the anions must equal the number of electrons lost by the cations so only one ratio can exist for each combination of ions.

17. 8.3 – Types of Ions Monatomic ions – ions made up of only 1 atom.
Most cations are monatomic, very few anions are monatomic
Examples: Mn2+, Al3+, Ag+, K+, Zn2+, C4+, O2-, N3-, Cl-
Polyatomic ions – ions made up of more than 1 atom.
The atoms making up the polyatomic ion are covalently bonded to create the ion.
Most polyatomic ions are anions, very few polyatomic cations
Examples: OH-, SO42-, NH4+, NO3-, CO32-, NO2-, ClO3-

18. 8.3 – Types of Ions Multiple Oxidation States
Some elements form ions with different charges (e.g., Fe2+, Fe3+). These elements are said to have multiple oxidation states.
To indicate in the chemical name which ion is present, the charge of the ion is listed as a Roman numeral following the name of the ion.
Fe2+ = Iron(II) Fe3+ = Iron(III)

19. 8.3 – Writing Ionic Formulas
An ionic formula represents the ratio of cations and anions in an ionic compound
The number of electrons gained by anions must equal the number of electrons given by the cations
Cations are always listed first in a formula
When no subscript is listed, the number of atoms is assumed to be 1
Write the formula of the cation, then anion with their charges
Find the lowest common multiple of the charges – this will be the number of electrons exchanged
Write a subscript to the lower right of the cation to indicate the number of cations necessary to generate the electrons to be given
Write a subscript to the lower right of the anion to indicate the number of anions necessary to accept the electrons to be gained

20. 8.3 – Writing Ionic Formulas
Examples:
Calcium Oxide
Potassium nitride
Aluminum oxide
Lowest common multiple of 2 and 2 is 2, so the atoms exchange 2 e-. Only one of each ion is required.
Ca2+
O2-
CaO
Lowest common multiple of 1 and 3 is 3, so the atoms exchange 3 e- . Three K+ and one N3- are required. K+
N3-
K3N
Lowest common multiple of 3 and 2 is 6, so the atoms exchange 6 e- . Two Al3+ and three O2- are required.
Al3+
O2-
Al2O3

21. 8.3 – Writing Ionic Formulas
Polyatomic ions
When polyatomic ions are a part of the compound, they should be treated like a single unit
the charge applies to the whole unit
the subscripts on the polyatomic ion should NEVER be changed
To show multiple polyatomic ions, place the formula of the ion in parentheses and the multiple of that ion to the lower right
Example:
Strontium phosphate
Lowest common multiple of 2 and 3 is 6, so the atoms exchange 6 e- . Three Sr2+ and two PO43- are required.
Sr2+
PO43-
Sr3(PO4)2

22. 8.3 – Naming Ionic Compounds
The name has two parts
First part = name of cation
monatomic cations take the name of the metal from which they are derived.
ions from elements with multiple positive oxidation states should include the Roman numeral of the charge number in parentheses.
Examples:
Co2+ = cobalt (II) or ‘cobaltous’ in the old style
Co3+ = cobalt (III) or ‘cobaltic’ in the old style
polyatomic cations use the name of the polyatomic ion.
Second part = name of anion
monatomic anions take the name of the non-metal from which they are formed and substitute –ide into the ending.
polyatomic anions use the specific name of the polyatomic ion.

23. 8.3 – Naming Ionic Compounds
Oxyanions
A polyatomic ion composed of one element, usually a non-metal, bonded to oxygen atoms (SO42-, PO43-, MnO4-)
When several oxyanions share the same non-metal, the endings of the names are changed to represent the amount of oxygen
Per__________ate
__________ ate
__________ ite
Hypo __________ ite
ClO4- = perchlorate
ClO3- = chlorate
ClO2- = chlorite
ClO- = hypochlorite
more oxygen
PO43- = phosphate
PO33- = phosphite
ClO2- = chlorite
ClO- = hypochlorite

24. 8.4 – Metallic Bonding Objectives: Describe metallic bonds.
Describe how metallic bonds relate to the properties of metals.
Describe alloys.

25. 8.4 – Metallic Bonding Electron-sea model
Metals often form lattices similar to ionic lattice structures
The outer energy levels overlap forming a “sea of electrons” that move from atom to atom freely
free electrons are referred to as “delocalized” are characterized as “highly mobile”
A metallic bond is the attraction of a metallic cation for delocalized electrons

26. 8.4 – Metallic Bonding Properties of Metals
Melting points vary greatly based on strength of attractions
Malleable
Ductile
Durable
Highly conductive because of mobile electrons
Delocalized electrons interact with light to create luster

27. 8.4 – Metallic Bonding Metal Alloys
Alloy – a mixture of elements that has metallic properties
Properties of alloys differ from the elements they contain
Substitutional – atoms in alloy are all about the same size
Interstitial – small atoms fill in gaps between larger atoms