1. Unit 5 – Bonding & Nomenclature Ionic Bonds, Covalent Bonds, and Writing Formulas

2. In our last unit…
We talked all about the ATOM
Development of the atomic model, subatomic particles
HOWEVER, we were only talking about atoms as one individual unit
In this unit we will be looking at how atoms combine (or BOND) with other atoms to form compounds!

3. Electrons are Arranged in Shells
As we move down a family, the elements gain an electron shell.
The first shell can hold up to 2 electrons. (Row/Period 1)
The second and third shells can hold up to 8 electrons each. (Periods 2 & 3)
The fourth and fifth shells can hold up to 18 electrons each. (Periods 4 & 5)
The sixth and seventh shell can hold up to 32 electrons each. (Periods 6 & 7)

4. The Valence Electrons Give Atoms Their Properties
The outer most shell is called the valence shell.
The electrons in the valence shell are called the valence electrons.
Atoms tend to gain or lose electrons to end up with a full valence shell. (8 e-)
Except H2 and He which want a duet (2 e-)
The number of valence electrons correspond to the digit in the ones place of the group number for elements in groups 1, 2,

5. Electron Dot Diagrams Show Valence Electrons
An electron dot diagram depicts an atom as its symbol and its valence electrons.
Ex: Carbon . . . C .
Carbon has four electrons in its valence shell (carbon is in group 14), so we place four dots representing those four valence electrons around the symbol for carbon.

6. Electron Dot Diagrams Show Valence Electrons
Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions, only doubling up if there are five or more valence electrons.
Example: Chlorine (7 valence electrons b/c it is in group 17) . . . . . Cl . .

7. Paired and Unpaired Electrons
As we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired.
When it comes to bonding, atoms tend to pair up unpaired electrons.
They do this in two ways that we will discuss in the chapter.
A bond that forms when one atom gives an unpaired electron to another atom is called an ionic bond.
A bond that forms when atoms share unpaired electrons between each other is called a covalent bond.

8. Ionic Bonding Ionic Bond
Bond between metal and nonmetal due to “electrostatic interactions”
Electrons are transferred from metal to nonmetal forming a positively charged metal ion (cation) and negatively charged nonmetal ion (anion)

9. 6.1 Ions and Ion Formation
If one or more valence electrons are lost or gained, the balance between positive and negative particles in the atom is upset.
When an atom gains electrons, there are more negative charges in the atom than positive charges, so it becomes negative
Example – Oxygen gains two electrons so it gains a charge of -2.
When an atom loses electrons, there are more positive charges in the atom than negative charges, so it becomes positive.
Example – Magnesium loses two electrons so it gains a charge of +2.

10. Atoms Do What is Easiest For Them!
Atoms tend to do the least to gain a filled valence shell.
Elements in groups 1, 2 and 13 tend to lose electrons to wind up with a filled valence shell because it only involves the removal of 1, 2, or 3 electrons, not the addition of 7,6, or 5 electrons respectively.
Likewise for elements in groups 15,16, and 17, only in their case, they tend to gain 3,2, or 1 electrons, not lose 5,6, or 7 electrons respectively.

11. Examples
Sulfur . . . . . . S . .
Sodium . . . . . . Na . .

12. Get Out Your Periodic Tables!
What are the charges for groups 1, 2, 13, 15, 16, 17, and 18?
Group 1:
Group 2:
Group 13:
Group 15:
Group 16:
Group 17:
Group 18:

13. Atoms That Gain Or Lose Electrons Are Called Ions (we saw this last chapter!)
An ion is formed when the number of electrons differs from the number of protons in an atom.
This is accomplished by gaining or losing electrons, not by adjusting the number of protons!
The charge of an ion is denoted as a superscript in the atom’s isotope notation.
Ex: A Nitrogen ion: 14 3- N 7

14. 6.2 Ionic Compounds
When an atom that tends to lose electrons comes in contact with an atom that tends to gain them, an electron transfer occurs.
As a result, two oppositely charged ions are formed.
The two oppositely charged ions are attracted to each other via the electrostatic attractions.
These attractions create an IONIC BOND

15. Ionic bonds Result from a Transfer of Valence Electrons+ -

16. Ionic Bonds Result from a Transfer of Valence Electrons
Ionic Compounds - All chemical compounds containing ions (one cation, one anion)
Ionic compounds typically are formed from elements on opposite sides of the periodic table.
For all ionic compounds, positive and negative charges must balance out.
Formula must be neutral overall!

17. Properties of Ionic Compounds
Positive and negative ions come together to create an organized pattern called a lattice structure.
An example of a lattice structure for sodium chloride

18. Properties of Ionic Compounds
Ionic compounds tend to be brittle
They have high melting and boiling points
They tend to be soluble in water.
They can conduct electricity when dissolved in H2O
Many of the physical and chemical properties of the components of the compound are changed when they bond.

19. 6.3 Writing Ionic Formulas+1 +2 +3 +1 +3 +2 -1 -2 -1 -2 -2 -1

20. 6.3 Writing Ionic Formulas+3 +3 -2 -2 -2

21. Subscripts Show the Number of Atoms of Each Element in a Compound
A subscript notes how many atoms of each element are in a subunit of a compound.
1 is NEVER used as a subscript.
Ex: NaCl  1 Sodium atom : 1 Chlorine atom
Al2O3  2 Aluminum atoms : 3 Oxygen atoms.
The subscripts determine how many of each atom is required to balance out the positive and negative charges from the ions.

22. Writing Formulas for Ionic Compounds
We can use the charges of the ions in a compound to easily determine the formula of the compound.
All you do is crisscross the charges to make them subscripts (Reduce the subscripts if you can!!)
Ex: magnesium nitride +2 -3 Mg N 3 2
Don’t forget to drop the positive/negative signs once they’re subscripts!!

23. PRACTICE, PRACTICE, PRACTICE!!
Write formulas for the following ionic compounds: Magnesium chloride Calcium bromide Sodium oxide Beryllium oxide Potassium iodide

24. Polyatomic Ions
There are some ions that are made of more than one element
These ions are called POLYATOMIC IONS
“poly” meaning more than 1, “atomic” meaning atoms  more than 1 type of atom in the ion
Examples: (see chart in your notes for more)
Sulfate SO42-
Hydroxide OH-
Nitrate NO3-

25. Writing Formulas w/ Polyatomic Ions
Write both the ions with their accompanying charges.
Crisscross the charges to make them subscripts.
Use parenthesis around the polyatomic ion if there is more than one of the ion!
Drop the positive and negative signs once they are subscripts.
Reduce subscripts if possible.
Example: Barium phosphate

26. PRACTICE WITH POLYATOMICS!
Don’t forget to reduce subscripts that come from the charges if you can!
Write formulas:
calcium phosphate beryllium bicarbonate
lithium carbonate
Magnesium sulfate
Barium hydroxide

27. Writing Formulas w/ Transition Metals
Transition metals – all elements in groups 3-12 and elements in groups 13, 14, and 15 below the stair step line
Can have more than one charge in ion form!
Except Ag+ and Zn2+
When writing formulas for compounds involving transition metals, you will get the charge on the metal cation from the Roman numeral in parenthesis after the metal’s name

28. Writing Formulas w/ Transition Metals
Writing Formulas of Compounds w/ Transition Metals:
Write the metal’s atomic symbol with the value of the Roman numeral as a positive charge
Write the value of the anion (negatively charged ion) with its charge
Crisscross the charges to get your formula
Example: Lead (IV) hydroxide

29. Writing Formulas w/ Transition Metals
Don’t forget to reduce subscripts that come from the charges if you can!
Examples:
Copper (II) chloride
Tin (IV) sulfate
Gold (I) nitrate

30. On a separate piece of paper:
Please write formulas for the following compounds & turn them in to me on a separate piece of paper so I can check them.
Sodium fluoride
Magnesium phosphate
Copper (II) sulfate
Tin (IV) phosphide
Hydrogen cyanide

31. 6.4 Naming Ionic Compounds
When naming ionic compounds, we take the name of the metal followed by the name of the non-metal, only we drop the last syllable and add –ide as an ending.
Examples: NaCl: Sodium Chloride Al2O3: Aluminum Oxide

32. PRACTICE, PRACTICE, PRACTICE!
Write names for the given ionic compounds:
LiBr
AlCl3
BeO
Ra3N2

33. Naming Ionic Compounds with Polyatomic Ions
The rules for naming polyatomic ions are the same rules for naming normal ionic compounds except you use the polyatomic ion name for the second word
Example:
MgSO4 – magnesium sulfate
LiNO3 – lithium nitrate
Ba(OH)2 – barium hydroxide

34. PRACTICE WITH POLYATOMICS!
Name the compounds:
RbNO3
NaHCO3
Mg(OH)2

35. Naming Formulas w/ Transition Metals
Transition metals – all elements in groups 3-12 and elements in group 14 beneath the stair step line
Can have more than one charge in ion form!
Except Ag+ and Zn2+
When naming compounds involving transition metals, you need to include a roman numeral in parenthesis to indicate the charge of the ion

36. Naming Compounds w/ Transition Metals
Naming Compounds Containing Transition Metals:
Name the metal from its symbol
Determine the charge on the metal by multiplying the subscript on the anion by the charge the anion normally has (see box below)
Write the charge as a roman numeral in parenthesis
Write the name of the anion either as an element with an “–ide” ending or as the unchanged name of the polyatomic ion
Charge of metal = |charge of anion × anion subscript|

37. Naming Compounds w/ Transition Metals
Examples:
CrCl6
Fe(OH)2
NiO

38. Covalent Bonding Covalent Bonds
Bonds in which e- are shared between two atoms
Most common type

39. 6.5 Covalent Bonds Result from a Sharing of Valence Electrons
Elements that have a partially filled valence shell can complete them by sharing electrons with another element
The mutual attraction for shared electrons is called a covalent bond (co- signifies sharing, -valent signifies valence electrons)

40. Shared Electrons Complete ShellsF F

41. 6.5 Covalent Bonds Result from a Sharing of Valence Electrons
A substance made up of atoms which are held together by covalent bonds is a covalent compound.
They are also called molecules.

42. Drawing Electron Dot Diagrams for Molecules
Chemists usually denote a shared pair of electrons as a straight line.
F F
Sometimes the nonbonding pair of electrons are left off of the electron dot diagram for a molecule

43. Nonmetals Tend to Form Covalent Bonds
A covalent bond is formed when two atoms that tend to gain electrons come into contact with each other.
Hydrogen tends to form covalent bonds because it has a fairly strong attraction for an additional electron.
The number of covalent bonds an atom can form is equal to the number of additional electrons it can attract.

44. Examples H2
H H O H
H2O H

45. More Examples H
CH4 H C H H H N H
NH3 H

46. O O O C O Multiple Bonds O2 CO2
Sometimes two atoms share more than just one pair of electrons which results in a double bond or triple bond. O O O2 O C O
CO2

47. Writing Formulas for Covalent Compounds (2 non-metals bonded together)
Write the symbol for the first element listed
If there is a prefix on the first element, write the prefix value as a subscript attached to the first element’s symbol
Write the symbol for the second element listed
If the prefix on the second element is “di-” or greater, write the prefix value as a subscript attached to the second element’s symbol
Ex: Diphosphorous monoxide 

48. Writing Formulas for Covalent Compounds (2 non-metals bonded together)
Examples:
Diphosphorous hexoxide
Carbon tetrachloride
Nitrogen trihydride

49. Naming Covalent Compounds (2 non-metals bonded together)
Given the formula:
Write the names of the 2 elements
Attach a Greek prefix (see chart) to the beginning of the first element only if it has a subscript greater than or equal to 2
Attach a Greek prefix to the second element no matter what!
Change the ending of the second element’s name to an –ide ending
Ex: Se2O3 

50. Naming Covalent Compounds (2 non-metals bonded together)
Examples:
P2O6
CCl3 NO

51. In review… Chemical Bonds Types of Chemical Bonds
Force that holds atoms together
It’s all about the electrons (e-)
Types of Chemical Bonds
Ionic
Bond between metal and nonmetal due to “electrostatic interactions”
Electrons are transferred from metal to nonmetal

52. In review… Covalent Bonds Metallic Bonds Bonds in which e- are shared
Most common type
Metallic Bonds
Atoms are bonded to one another (not to other elements)
Positive ions in a “sea” of negative charge (e-)
Example: Cu, Ag