1. Bonding Unit Today we will: -Define Ionic, and Covalent Bonding
-Discuss ionic and covalent properties
-Learn to draw Lewis Structures
-Be Chemistry Match Makers

2. Ionic Compounds
Ionic compounds are composed of both metals and nonmetals. The bond that is formed is based on electrostatic forces between negatively(anion) and positively(cation) charged ions.
Electrostatic forces remember are the attractions between positive and negative ions that hold the bond together.
Ionic bonding occurs by the transfer of one or more electrons from one atom to another

3. Properties usually solid at room temperature have high melting points
usually do not conduct electricity as a solid
usually dissolve in water
usually conduct electricity when in solution or molten state
1+2-Strong bonds are formed so it is hard to break them apart
3-ions can’t move, and when compound is formed it doesn’t want to give away any electrons
4-since water is a polar molecule it attracts ions

4. Ionic Bonds
An ion is an atom or group of atoms that have a charge. Atoms normally have a neutral charge because most often they have the same number of electrons and protons. They become ions by the loss or addition of one or more electrons. This process is called ionization.
To understand ionic bonding we will develop an understanding of ions.
An ion that has more electrons than protons is called an anion, and an ion that has fewer electrons than protons is called a cation.

5. Ionic Bonds
The interaction of ionic bonds is when atoms gain or lose electrons until the outer shell of electrons is full and stable with 8 electrons. This is part of the octet rule.
Recall octet rule:
When atoms combine to form molecules they generally each lose, gain, or share valence electrons until they attain or share eight and reach a noble gas electron configuration
Stress that the octet rule is useful but it does not work in every case, and there are exceptions to the rule.

6. Ionic Bonds
The number of electrons the atom gains, loses, or shares is called its Valence.
Nonmetals usually have four or more electrons in their outer shell. To make their outer shell full, it’s easier(it takes less ionization energy) for them to gain three or four electrons than to lose four or five electrons.
When you look at the metals, they usually have three or less electrons in their outer shell. Opposite from nonmetals, it is easier for metals to lose three or less electrons than to gain four or more.
Therefore it makes sense that metals and nonmetals bond together easily.

7. Lewis Dot Structure
In 1902 Gilbert Newton Lewis invented the valance bond theory.
Lewis came up with an easy way to represent electrons in the outer shells of ions. His invention is called “Lewis Dot Symbols”.
Lewis structures are used to visualize the valence electrons of elements. In the Lewis model, an element symbol is inside the valence electrons of the s and p subshells of the outer ring.
Lewis was a famous American physical chemist.

8. Lewis Dot Structures
Here are the Lewis dot structures for some elements. The transition metals, lanthanides and actinides are not displayed by Lewis Dot structures because they do not follow the octet rule. Their first outer orbital shell only has the capacity for two electrons.

9. Lewis Dot Structure
1)Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
2)Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
3)Complete an octet for all atoms except hydrogen
4)If structure contains too many electrons, form double and triple bonds on central atom as needed.

10. Lewis Dot Structure Example:
Sulfur has 16 electrons. You can show this by using the method you learned earlier in the year [Ne]3s23p4. You can see from using this method that there are 6 electrons in the outer shell of sulfur. Sulfur can be represented in Lewis dots when you follow these rules:
1. The element symbol is placed in the center Dots, representing electrons, are placed on all four sides of the element Each side can hold 2 electrons You must put an electron on each side (if possible) before filling up a side.

11. Lewis Dot Structure
This is what sulfur looks like according to the Lewis Dot Chart: S

12. Lewis Dot Structure
Now it’s your turn to try and draw some elements using the Lewis Dot Structure Potassium Germanium Phosphorus Neon 5. Aluminum

13. Lewis Dot Structure

14. Covalent Compounds
Covalent compounds are made up of two nonmetals. A single covalent bond is formed when a pair of electrons is shared between two atoms
There are two types of covalent bonds: non- polar covalent and polar covalent.

15. Properties
simple molecular substances have low melting and boiling points
larger more complex compounds will have higher melting and boiling points
usually do not conduct electricity as a solid or when molten or in solution
usually do not dissolve in water
1-weak intermolecular bonds
2-many strong covalent bonds
3-no mobile electrons, no way to transfer electrons
4-molecules are not charged, thus polar water
molecules do not attract them

16. Non-Polar Covalent Bonding
Accounts for the bond that keeps two atoms of the same element together. (Cl2, H2)
Atoms share electrons from ½ filled orbitals to achieve noble gas configurations
Shared electrons are attracted to both nuclei, which keeps atoms together
Electrons involved in bonding are called shared electron pairs, ones that are not are called lone electron pairs

17. Polar Covalent Bonds Account for the bonding found in HF
One electron from each. atom is shared but not equally due to unequal attraction for shared electrons
The bond is referred to as polar because 2 poles are formed
(+ and -)
Electronegativity values allows us to determine which atom has a greater pull
The atom with the greater electronegativity becomes the negative end of the polar bond.
The atom with the lower electronegativity becomes the positive end of the polar bond

18. Example-HF H F F H
e- rich
e- poor + -

19. Electronegativity
Electronegativity is the tendency of an atom to draw or attract the electrons in a bond toward itself
Electronegativity is like a game of tug-of-war, atom's ability to pull determines what kind of bond it forms
To form a covalent bond, two or more atoms with similar electronegativities will share electrons
Values fall between a low of 0.7 for Fr and a high 4.0 for F
The greater the difference in electronegativity the more polar the bond.

21. Double Covalent Bonds
Compounds sometimes share two pairs of electrons and form a double bond. This often occurs when two atoms of the same element bond, but also occur between different elements.
This is called Double Covalent bonds
Examples: O2, CO2

22. Triple Covalent Bond Same idea as single and double
Two atoms of the same element or two different elements share three pairs of electrons and form a triple bond
Example: N2

23. Coordinate Covalent Bond
A covalent bond in which one of the atoms donates both electrons
Properties are same as previous covalent bonds
Distinction is useful when keeping track of valence electrons and when assigning formal charges.
Example formation of ammonium ion

24. Methods to Classify Bond Type
Subtract the electronegativity of the least electronegastive atom from the most
Divide the difference by the greater value
Covalent Bonds are in the range of less than 0.17, polar covalent bonds are between and 0.45, and ionic bonds are greater than 0.45
Method 2
Covalent bonds=no difference in electronegativity
Polar covalent bonds=difference less than 1.7
Ionic Bonds= difference of 1.7 or more.

25. Chemistry Match Makers
Draw element
Research it
Make dating profile
½ class draw profiles
Get into group
Decide if it’s a match and draw it/classify it/give reasons why
Report results