1. Ionic Bonds

2. Ionic Bonds & Compounds
Octet Rule- the tendency of atoms to gain atoms to gain or lose electrons so that their outer s and p orbitals are
full with 8 electrons.

3. Atoms of metals
Can achieve a stable “noble gas” configuration by losing electrons. These atoms form simple cations.
Example: Na+, Li+, Ca2+, Al3+

4. Atoms of Nonmetals
Can achieve a stable “noble gas” configuration by gaining electrons. These atoms form simple anions.
Example:
Cl-, F-, O2-, S2-, N3-

5. Ionic Compounds
Are any chemical compounds that are composed of oppositely charged ions.
Example:
Cation + Anion from a metal and a nonmetal
Na+ + Cl > NaCl

6. Binary Ionic Compounds
Are compounds that consists of cations of one element and anions of another element.
Their combined charge must cancel--equal zero.
Special Note: You must memorize all cation and anion charges for the main group elements.
Gp 1, 1+; Gp 2, 2+; Gp 13, 3+; Gp 14, 4+/-; Gp 15, 3-; Gp 16, 2-; Gp 17, 1-

7. Nonpolar Covalent Bond
Is a covalent bond in which bonding electrons are shared equally by the bonded atoms resulting in a balanced distribution of charge.
electronegativity value

8. Polar Covalent Bond
Has an uneven distribution of charge. This is where the electrons shared are more strongly attracted by the more electronegative element.
0.5 – 1.8 electronegativity

9. Chemical Bonds
Are bonds formed from the electrical attraction between the nuclei and valence electrons of atoms that binds the atoms together.

10. Valence Electrons
Electrons available to be gained or lost or shared in the formation of chemical compounds.
For main group elements, they are those electrons usually in the outer s and p orbitals subject to influence by nearby atoms or ions.

11. Ionic Bonds
Are formed from chemical bonding that results from the electrical attraction between cations and anions.
This involves the giving and taking of electrons.
electronegatibity
Metal + Nonmetal
Na+ + Cl- --> NaCl

12. Covalent Bonds
Result from sharing of electron pairs between two atoms.
Nonmetal + Nonmetal
C + O2 --> CO2

13. Electronegativities in Bonds
The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in the elements electronegativities. The difference must be in absolute value.
Example: F = 4.0 and H = 2.2
|1.8|

14. Electronegativity Values
Nonpolar Covalent Bonds
Polar Covalent Bonds
Ionic Bonds

15. Naming Binary Ionic Compounds
Binary means two
Formula a shorthand way to write
the name of a compound
using symbols and subscripts.

16. Naming Ionic Compounds
Cation- the cation is written first in the formula and its name is the same name as its element on the periodic table.
Example: Na+ = Sodium

17. Continued
Anion - the anion is written last in the formula and is given the same name as its element on the periodic table, but with an ide suffix.
Example: Cl- Chlorine --> Chloride
Compound Formula NaCl
Compound Name Sodium Chloride

18. Other Examples ZnS Zinc Sulfide Mg3N2 Magnesium Nitride
Other ide endings:
Oxide, Bromide, Iodide, Fluoride, Nitride, Selenide, Phosphide, Hydride, Silicide, Etc…

19. Balancing Charges
The charge of the cation and anion must equal zero before writing the formula and naming the compound.
It is based on the Law of Multiple Proportions.

20. The Law of Multiple Proportions
The mass for one of the elements combines with a fixed mass of another element and can be expressed in small whole numbers.

21. Subscripts
Are small whole numbers written below and to the right of the atom’s symbol to denote the number of atoms needed to complete the formula.
Example: 1). In order to determine the formula for a simple binary ionic compound you must balance the charges of the cation and anion.
Na+ + Cl and = 0

22. Formula & Name
2). Write the formula. Since you only used one Na and one Cl, your formula is NaCl.
3). Write the name per the rules.
Sodium Chloride
Complete Al S2-

23. Example
Al3+ + S2-
You use two Al3+ ions plus three S2- ions to balance the charges
The formula becomes Al2S3. The subscripts identify how ions you use of each.
The formula name is Aluminum Sulfide

24. Covalent Bonds
Covalent Bonds- a bond formed when two or more valence electrons are attracted by the positively charged nuclei of two atoms and are thus shared between both atoms.

25. Molecular Orbital
After the two atoms get together and share their valence electrons they form a common orbital. Example: H2 H:H This is the simplest form of a covalent bond- a diatomic covalent bond

26. Explanation
As the two Hydrogen atoms approach each other, the positive nucleus of each atom attracts the other’s single electrons. At the same time, the positive nucleus of each repels each other and the electrons of each repel each other.

27. Conclusion
You have a tug of war between like charges repelling each other and opposite charges attracting each other with neither nucleus having enough energy to remove an electron and so they are shared instead.

28. Types of Covalent Bonds
Polar covalent
Nonpolar covalent
Metallic

29. Metallic Bond
Metallic Bond- a chemical bond of metals in which valence electrons are shared among the atoms in a usually stable crystalline structure. Example: Cu:Cu

30. Lewis Dot Structures
Lewis Dot Structures- are structures in which symbols represent nuclei and inner electrons and dots are used to represent valence electrons. Example: O2

31. Single, Double, & Triple Bonds
In covalent bonds, electrons are shared as pairs. Either one pair, two pairs, or three pairs may be shared at one time. Single Bonds- one pair of electrons is shared by the two atoms. Group 17 atoms for single bonds. Example:

32. Single Bond
A single line joining the Chlorines represents one pair of electrons between the two atoms.

33. Double Bonds
Double Bonds- two pairs of electrons are shared by the two atoms. Group 16 atoms form double bonds.
A double line represents two pairs of electrons shared
two atoms.

34. Triple Bond
Three pairs of electrons are shared by two atoms. Group 15 atoms form triple bonds.
Triple lines represent three pairs of electrons shared by two
Atoms.

35. Conclusion
In this way each atom gets the pairs to make its octet. The electrons that are shared vibrate back-n-forth to complete the octet.

36. Naming Covalent Binary Compounds
Naming- The least electronegative element is written 1st unless carbon is present. It is written 1st. The name is written the same as it is on the periodic table with a prefix telling how many atoms are used. Example: H2S Dihydrogen Monosulfide

37. Naming Continued
The di- prefix is placed before Hydrogen to denote that two atoms of Hydrogen were used in the formula. Exception- There is one exception to this rule. The first element in the formula does NOT contain a prefix if there is only one atom contained in the formula. Example: CO Carbon Monoxide NOT Monocarbon Monoxide

38. Naming Continued
The 2nd element is named the same as it is on the periodic table. A prefix is added telling how many atoms are used. Also, an –ide suffix is added.

39. Prefixes
One- mono Eight- octa Two- di Nine- nana Three- tri Ten- deca Four- tetra Five- penta Six- hexa Seven- hepta

40. Example
CO2 Carbon Dioxide SiO2 Silicon Dioxide C4H10 Tetracarbon Decahydride

41. Rules for Making Bonds Using Lewis Dot Structures
Determine the type and number of atoms in the molecule. ie…metal/nonmetal or nonmetal/nonmetal and number of atoms.
Write the electron dot notation for type of atom.
Determine the total number of valence electrons available for bonding in the atoms to be combined.

42. Rules Arrange the atoms to form a skeleton structure for the molecule.
If carbon is present, it is the central atom. Otherwise, the least electronegative element is central. Hydrogen is NEVER central.
Connect electrons by electron pairs.

43. Rules
Add unshared pairs of electrons to each nonmetal atom except Hydrogen such that each atom is surrounded by eight electrons.
Count the electrons in the structure to be sure that the number of valence electrons equals the number available.
Be sure the central atom and other atoms besides Hydrogen (two electrons) have an octet!

44. Resonance Structures
Examples

45. Molecular Structure

46. The VSEPR Model

47. The VSEPR Model
Two Pairs of Electrons

48. The VSEPR Model
Three Pairs of Electrons

49. The VSEPR Model
Four Pairs of Electrons

50. Steps for Predicting Molecular Structure Using the VSEPR Model

51. Steps for Predicting Molecular Structure Using the VSEPR Model

54. Molecules with Double Bonds