1. Ionic Bonding

2. Ionic Bonding
Occurs when electrons are transferred from one atom to another, forming two ions
The ions stay together because of electrostatic attractions
Ionic bonds NEVER form molecules
Ionic bonds form easily between alkali metals and halogens

3. The Octet Rule
Atoms tend to gain, share, or lose electrons in order to obtain a full set of valence electrons (in most cases this equals 8)
An octet of electrons consists of full s and p sublevels on an atom.
Exceptions: transition elements and rare earth elements

4. Example - + Na Cl Na + Cl +

5. Properties of Ionic Compounds
Ionic compounds do not form molecules; they form a crystal lattice
The green spheres are Na+ and the red spheres are Cl -

6. This is a crystal of CaCl2
This is a crystal of CaCl2. Each ion is held rigidly in place by strong electrostatic forces that bond it to several oppositely charged ions

7. Other Properties Normally form between metals and nonmetals
Ionic compounds have ions that form very strong bonds, which makes them hard and brittle
They have high melting points and high boiling points
When dissolved in water, the solution will conduct electricity

8. Types of Ionic Compounds
Ionic compounds will be a combination of a metal and a nonmetal (if the cation is monatomic)
There are two types of ionic compounds
Binary Ionic Compounds: contains the ions of only two elements (NaCl, CaCl2)
Polyatomic Ionic Compounds: contain at least one polyatomic ion (CaCO3, Mg(OH)2)
The names DO NOT indicate the ratio of ions present, but the formulas do

9. Types of Ions There are two types of ions
Monatomic: cation or anion that consists of a single atom. Examples: Na+ and Cl-
Polyatomic: two or more atoms that act as a single ion (or particle). Examples: (CO3)2- and (OH)-

10. Names of Charged Particles
When electrons are lost, the ion has an overall positive charge and is called a cation Examples: Na+, Ca2+
When electrons are gained, the ion has an overall negative charge and is called an anion Examples: F-, S2-
The negative ions will attract the positive ions and form a bond
All ionic compounds are electrically neutral

11. Ionic Charges Monatomic ions Ions that consist of only one atom
Charges often can be determined by using the periodic table
Many of the transition metals have more than one charge
The charge on the ion is indicated by using a Roman numeral next to the name of the element
Cu has a 1+ and a 2+ charge. Cu 1+ is called Copper I, and Cu2+ is called Copper II
A few transition metals have only one charge
The names of these do not have to include a Roman numeral
Zn, Cd

12. Atoms that commonly form ions
Monatomic Ions
Group
Atoms that commonly form ions
Charge on ions 1
H, Li, Na, K, Rb, Cs 1+ 2
Be, Mg, Ca, Sr, Ba 2+ 13
B, Al 3+ 15
N, P, As 3- 16
O, S, Se, Te 2- 17
F, Cl, Br, I 1-

13. Ionic Compounds Writing formulas for binary ionic compounds
Compounds composed of two elements are called binary compounds
When the formula is written, the charge of the cation must be balanced by the charge of the anion
The overall charge of the ion combination must be zero
The cation is always written first in the formula

14. Writing Formulas for Ionic Compounds
To write the formula for an ionic compound:
Write the chemical symbol and overall charge of the cation or polyatomic cation: Na+, Ca2+, NH4+
Write the chemical symbol and overall charge of the anion or polyatomic anion: Cl-, O2-, NO3-
Add the charges
If equal, write the chemical symbols together, e.g. NaCl, CaO, NH4NO3
If not equal, crisscross values of the charges and make them subscripts, e.g. Ca(NO3)2, CaCl2

15. Crisscross Method (also called “Drop and Swap”
Na+ can combine with S2-
The value of the charge on Na, which is 1, becomes the subscript for S: S
The value of the charge on S, which is 2, becomes the subscript for Na: Na2
The resulting formula is Na2S

16. Naming Ionic Compounds
Naming binary ionic compounds
It is important to know the Stock naming system and the charges on cations before naming ionic compounds
The process is the reverse of writing formulas

17. Polyatomic ionic compounds
Compounds that contain atoms of three different elements
They usually contain a polyatomic ion; must be recognized first to be able to name the compound correctly
Tightly bound groups of atoms that behave as a unit and carry a charge
Have the suffix –ite and –ate
These ions will be given to you on a sheet; you will not have to memorize them
Hydroxide and cyanide are the only common polyatomic ions that end in –ide

18. Common Polyatomic Ions
Name
NH4+
Ammonium
NO2-
Nitrite
NO3-
Nitrate
OH-
Hydroxide
CO32-
Carbonate
SO42-
Sulfate
O22-
Peroxide
C2H3O2-
Acetate
SO32-
Sulfite
ClO3-
Chlorate
You will get these on a chart for tests and quizzes!!

19. How Do I Write the Formulas for Polyatomic Ions?
Put parentheses around any polyatomic ion and add subscripts to the outside of the parentheses.
Never change the subscript of a polyatomic ion (it will change the composition of the ion):
Ca(OH)2, Ca3(PO4)2
Include the subscript inside the parentheses

20. Practice Sodium combines with fluorine Chlorine combines with hydrogen
The nitrite ion combines with calcium
The ammonium ion combines with nitrogen

21. Percentage Composition
The percent by mass of any element in a compound can be found by dividing the mass of the element by the mass of the compound and multiplying by 100%
Example: 65g of element X and 45g of element Y in 110 g of compound
%X = (65g ÷ 110g) x 100% = 59%
%Y = (45g ÷ 110g) x 100% = 41%

22. Examples: What is the percentage of oxygen in H2O?
What is the percentage calcium in Ca(OH)2?

23. Representing Chemical Compounds
The Laws of Definite and Multiple Proportions must be obeyed:
Law of Definite Proportions
States that in samples of any chemical compound, the masses of the elements are always in the same proportions (e.g. H and O always combine as H2O to form water)
Law of Multiple Proportions
Whenever element combine to form a compound, they will always combine in small, whole-number ratios with each other (e.g. H2O, CO2, H2CO3, etc)

24. What is molar mass? Mass of 1 mole of a compound Measured in grams
It is the mass in grams of one mole of a given substance
It is equal to the average atomic mass of that element, written in grams
It depends on the masses of the elements that make up the substance

25. How to find molar mass Find the molar mass of Na2SO4
Find the molar mass of MgSO4 · 7H2O

26. How do I convert from moles to mass?
You must use molar mass
How many moles of water are in g of water?
What is the mass of 15.2 moles of CO2?

27. What is empirical formula?
Empirical formula: gives the simplest whole number ratio of the atoms of the elements
Gives the ratio of the atoms, but doesn’t necessarily give a correct molecular formula

28. To find empirical formulas
To calculate:
Convert the mass (g) of each element in the compound to moles
Compare the moles in a ratio
Put the ratio in lowest terms by dividing by the smallest number in the ratio
Round the ratio to whole numbers
It is easy to go from % composition to empirical formulas by assuming the percentage is the amount of element in grams (i.e., drop the % and replace with a “g”).

29. Examples What is the empirical formula of the compounds below?
59.95% O; 40.05% S
48.64% C; 8.16% H; 43.20% O

30. Molecular vs Empirical Formulas
Molecular gives actual number of atoms of each element in a molecule
Always a whole # multiple of empirical formula
Empirical gives ratios of each element in each compound
Subscripts always in lowest possible whole number

31. To Determine Molecular Mass
Divide the molar mass of the unknown compound by the molar mass of the empirical formula
This lets you know what multiple of the empirical formula the formula of the unknown is
molecular formula = (empirical formula)n

32. Example
A compound was found to contain 65.45% C, 5.45% H, and 29.09% O. The molar mass of the compound is g/mol. What is the molecular formula?

33. Example
A colorless liquid composed of 46.68% N and 53.32% O has a molar mass of g/mol. What is the molecular formula?