1. Warmup:
For each of the following, (1) determine if they will gain or lose electrons, (2) how many they will gain/lose, (3) write out the ionic symbol.
Oxygen (#8) Potassium (#19) Calcium (#20)

2. New Objectives: Red, Yellow, Green
Goal 1: I can identify all 3 types of chemical bonds.
Goal 2: I know how to use subscritps and put together cations and anions to make an ionic compound.
Goal 3: I know the difference between monatomic and polyatomic ions

3. 3 types of bonds Ionic: very strong, creates solids.
Polar: medium strength bond. Some solids, liquids, and gases
Covalent: very weak bond. Some solids, liquids, gases

4. Bond type depends on how electrons are shared between atoms
Electronegativity: An atom’s attraction for electrons
High EN = strong attraction
Low EN = low attraction

5. How to figure out bond type:
You must calculate at the electronegativity difference between the two atoms
Use the approximate guideline for electronegativity differences:
0.0 – < 0.5 = covalent (electrons shared nicely)
0.5 – < 2.0 = polar covalent (electrons being pulled from stronger atom)
2.0 – 4.0 = ionic (electrons stolen/transferred by stronger atom

6. Examples: H and Cl 2.1 – 3.0 Difference = 0.9 = polar C and H
2.5 – 2.1
Difference = 0.4 = covalent

7. Examples: Na and Cl 0.9 – 3.0 Difference = 2.1 = ionic
Also notice that you have a metal bonded to a nonmetal!
Follow this link for a video recap:
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8. Use these values provided to help:
EN increases
EN decreases

9. Atomic Size

10. Why?
As you go across a period, the principal energy level (# of rings) remains the same. As the nuclear charge (# protons) increases, it draws those electrons in closer to the nucleus. Atomic size therefore decreases going across the periodic table,
Atomic size increases when you move down the periodic table because more rings of electrons are being added.

11. Ionization Energy
Energy required to remove an electron from a gaseous atom. This causes a +1 charge.
The energy required to remove the first outermost electron is called the 1st ionization energy.
The larger the atom is, the easier it is to rip away an electron. Electrons that are far away from the nucleus are easier to pull away.

12. Electronegativity Trends
Increases
Decreases

14. Ionic Bonding

15. Ionic Bonding Bonding of ions Ionic Bond Ionic Compound
The electrostatic force that holds oppositely charged particles together in an ionic compound. One or more electrons are transferred from cation (metal) to anion (nonmetal or group of nonmetals).
Ionic Compound
Always formed between a metal and a nonmetal (or a group of nonmetals)

16. Compound Formation and Charge
Want compounds to have a neutral charge
Number of electrons lost = number electrons gained
Example: CaF2
Calcium wants to lose 2 electrons to be like Argon.
Fluorine can only gain 1 electron to be like Neon.
This compound must use 1 Calcium and 2 Fluorine atoms.

17. Formula Unit
The simplest ratio of ions represented in an ionic compound
Overall charge of the formula unit is zero

18. Formula Writing Symbol of the cation is always written first
Symbol of the anion is always written second
Subscripts represent the number of ions present
Example: What is the formula for an ionic compound containing magnesium and bromine?

19. Naming Ionic Compounds
Recall:
Cations keep the original element’s name
Ex. A sodium ion is still called sodium
Anions place “ide” at the end of their element’s name
Ex. A chlorine ion is called chloride
Since ionic compounds are made of cations (first) and anions (second) we will use the above rules to name ionic compounds
Example: NaCl is called sodium chloride
What would MgS be called?
Magnesium sulfide

20. Polyatomic ions “poly” = many “atomic” = atoms
Group of atoms bonded covalently that have an overall charge.
Ex: PO4-3 = “phosphate” ion (5 atoms: 1 P atoms and 4 O atoms)
Notice the ending of the name: -ate. This is NOT a monatomic anion from the periodic table
You WILL NOT be required to know their formulas or charges. We have a chart for you.

21. You will be provided a LIST of polyatomic ions and their charges

22. Rules for using polyatomic ions
If you need more than one polyatomic ion in a formula, you must group it together in parenthesis with a subscript outside.
Ex: Try to make a balanced compound using magnesium and Chlorate.
Remember: if the name of an ion ends in –ide, it came from the periodic table and is a monatomic ion. The end of this name is –ate
If the ion ends in –ate or –ite…look it up on your list!!!

23. Some examples mixed up:
Try these:
Sodium oxide
Calcium hydroxide (this is a polyatomic anion event though it ends in –ide)
Strontium sulfite (anion ends in –ite…look it up on your list…it’s polyatomic)
Barium nitrate
Silver chloride
Rubidium sulfate

24. Writing the Formula from the Name
Find the elements that make up the compound
Determine the charge of the ions
Determine the ratio of the ions in the compound
Write the formula, using () if you need more than one polyatomic ion.
Example: What is the formula for calcium iodide?

25. Red? Yellow? Green?
Goal 1: I can identify all 3 types of chemical bonds.
Goal 2: I know how to use subscritps and put together cations and anions to make an ionic compound.
Goal 3: I know the difference between monatomic and polyatomic ions