1. Before Bell Rings Turn in Nucleosynthesis Packet on front table Grab two papers from the side table Chill out/Cram for test

2. Nuclear Unit Test
Everything away.

3. Unit 3-Naming

4. Learning Targets
I can determine how two ions will bond to form an ionic compound.

5. Writing an Ionic Formula
Key: Total charge on the compound must be neutral!
Biggest idea is that the charge for the entire compound has to be zero. A compound must have a amount of negative charge equal to the positive. So if there is a total of +6 charge from the cations there has to be a -6 total charge from the anions in the compound.

6. Barium and Oxygen BaO – Barium Oxide
If we look at barium and oxygen forming a compound notice that barium (Ba) wants to give up two electrons to have a full valence shell. Oxygen on the other hand needs to take in two electrons to have a full valence shell. The charge of the barium ion is 2+, since it will give up two electrons, and the charge of the ion of oxygen, the oxide ion, is 2- since it takes in two electrons. This means we would need one of each to have a neutral charge of zero = 0

7. Sodium and Oxygen Na2O – Sodium Oxide
Sodium wants to give up 1 electron, so it will have a +1 charge for its ion. Oxygen wants to take in two electrons, so it will have a 2- charge. This means I need two sodium ions for every one oxygen ion.
Na2O – Sodium Oxide

8. Aluminum & Sulfur Al2S3 – Aluminum Sulfide
Hint: To make the compound neutral
Look for the least common multiple and make the total charge provided by of each type of ion equal to this
6+ vs 6-
Same thing here. Each aluminum atom gives away three electrons, each sulfur wants to take in two electrons. If I bond just one sulfur to one aluminum, the aluminum still wants to give away one electron, so we have to add another sulfur. This makes us unbalanced again, so we add another Al. We then have two extra electrons so we add one more sulfur to end up with two Al and 3 S or Al2S3.
When making compounds with ions with different charges a good method to figure out how many of each you need is to fine the least common multiple (LCM) of the two charges. In this case the charges were 3+ and 2- meaning the LCM was 6. This means I need to have enough Al to make 6+ charge and enough S to make 6-. This means 2 Al and 3 S.

9. Naming Rules for Chemical Formulas
Chapters 3 and 4

10. Protons Stay the Same Protons = ID
Atomic Mass is the mass of protons and neutrons – it is the mass of the nucleus of an atom. An element has weight but it is ridiculously small compared to that of a proton. Example of a person on a scale being handed a pen – how much would the weight change.
The number of protons determines and element’s identity. You can change the number of electrons or neutrons and the element will still be Carbon. But if you change the number of protons it isn’t carbon anymore. Also an atom can not be an atom unless it has protons, electrons, and neutrons. An atom is composed of all three.

11. But e- can change

12. Ionization: Turning atoms into Ions
Stealing their e-

13. - + Metals vs Non- Metals Metals =Losers Non-Metals = Thieves (Takers)
Positive = Ca+ions
Non-Metals =
Thieves (Takers)
Neg = Anions - +
Metals (elements to the left of the staircase on the PT) are all going to lose electrons, making their ions positive. We call positive ions cations. A good way to remember this is the T in cation looks like a + sign, so you know it is the positive ion.
Non-metals (elements to the left of the staircase) all want to gain electrons. This makes their ions negative and we call negative ions anions. A clue here is that at the start on anion we see an N just like the start of Negative.

14. + - Greediest Givers Losers
Cations give electrons away, anions take them in.
Givers
Losers

15. Ionic Bond A bond formed because e- transferred
Between a metal & non-metal
Electronegativity Difference =
LARGE!!!
The electronegativity difference has to be large because one has to want to attract electrons and the other doesn’t want to

16. Why Share if your strong enough to steal?
Stealing e- = Ionic
Why Share if your strong enough to steal?

17. Ionic Bond Often between + metal ion and – non-metal ion
Bond is an “attraction” between + and - charges
Remember an ionic bond forms when one atom loses an electron to another atom, becoming a cation since it will have a positive charge. The atom that took the electron is the anion because it will have a negative charge. The bond comes from the attraction between the positively charged cation and the negatively charged anion since opposites attract.

18. Ionic Compounds

19. Why does one Na bond with just one Cl?
How many Cl would bond with one Ca?
CaCl2
We only need one sodium (Na) to bond with one chlorine (Cl) because sodium only needs to give away one electron to have a full outer shell and chlorine only neds to take in one electron to have a full outer shell.
Calcium (Ca) on the other hand needs to give away two electrons to have a full outer shell, so it will have to bond with two chlorine (Cl) atoms to make a compound, since each chlorine only takes one electron.

20. Writing Ionic Formulas
Sodium & Oxygen
Na2O
Al2S3
BaO
1) Metal always1st
2) Total charge must = Ø
3) Subscripts = number of atoms
4) Subscripts should be reduced when possible
Here are some of the rules for writing ionic formulas.
When we write formulas, the cation is always the first ion we write. This is usually a metal, but remember there are non-metal cations as well.
Just like we have been working on, the overall charge for the compound has to equal zero. So if the cation has a total of a 2+ charge, the anions have to have a total of a 2- charge so they will balance out to zero. This is the main rule for ionic compounds, so make sure you understand it.
If there is a subscript, that represents the number of atoms of that element. So if it says Na2, this means the compound has two sodium ions in it.
Compounds should always be the smallest ratio of the ions. So if a compound had Fe4S6, we could reduce that to Fe2S3 to give the smallest ratio. 2+ 2- Ø

21. You try writing the formula
K3P- Potassium Phosphide CaBr2 - Calcium Bromide
NaI - Sodium Iodide Mg3N2 – Magnesium Nitride
AlF3 - Aluminum Fluoride BaS - Barium Sulfide
Potassium & Phosphorus
Calcium & Bromine
Magnesium & Nitrogen
Sodium & Iodine
Aluminum & Fluorine
Barium & Sulfur

22. Polyatomics: Many atoms, but One Ion with One Charge1-
Polyatomic ions act just like an ion made-up of single atoms, but the whole polyatomic ion has a given charge, not each atom in it. So if I wrote NH41+, it means the charge for the entire ion, ammonium (NH4), is 1+.

23. You Need To Memorize Names and Charges
Need to memorize these, as well as the ones that are starred in your ions sheet and the underlined ones on the next slide.

24. Polyatomic Ion Names ClO-1 – hypochlorite ClO2-1 – chlorite
ClO3-1 – chlorate
ClO4-1 – perchlorate
Several atoms that are bonded together with one ionic charge
Ex. CO3-2  carbonate ion, whole ion wants to gain 2 electrons
SO2-2 – hyposulfite
SO3-2 – sulfite
SO4-2 – sulfate
SO5-2 - persulfate
When naming polyatomic ions, the name changes based on the number of oxygen atoms in the ion. The ion with the lowest number of oxygen atoms has hypo- as a prefix and -ite as a suffix. The next least number of oxygen atoms has the –ite suffix. Next comes the –ate suffix, and finally the per- prefix and –ate suffix.
Note that an ion with the -ate ending always has a larger number of oxygen atoms than the -ite suffix. This helps when keeping the different polyatomic ions straight.
-ate is always bigger than –ite
You only need to know the underlined ones

25. Ionic Formula w/ Polyatomics
Same rules, but Use Parenthesis for more than one polyatomic
Ca(ClO3)2
Magnesium & Nitrate Lithium & Phosphate
Mg+2 and NO3-1  Mg(NO3)2 Li+1 and PO4-3  Li3PO4
not MgNO6 or MgN2O6
Ammonium & Sulfur Barium & Hydroxide
NH4+1 and S-2  (NH4)2S Ba+2 and OH-1  Ba(OH)2
not BaOH2!!
When writing formulas involving polyatomic ions, treat it the same as without but if you need multiples of the polyatomic ion you don’t just add a subscript, you have to put the polyatomic ion in parenthesis first. The helps keep the number of atoms in the compound correct and makes sure we know what ions are present in the compound.

26. Endings for non-metals!
IDE  Single non-metal (S-2)
ATE  Poly with the larger subscript (SO4-2)
ITE  Poly with the smaller subscript (SO3-2)
NEVER change the original polyatomic – just put it in parenthesis
To get the name of the anion when it is on its own (not part of a polyatomic ion-i.e. just S, just Cl, just F), we change the ending to ide to indicate it is the ion.
When determining what name to give a polyatomic ion, the one with the larger subscript (more oxygen atoms) will get the –ate ending, the one with less will get the –ite ending.

27. Ide ATE BIG, with little bites
S = sulfIde
SO4-2 = SulfATE
SO3-2 = sulfite
Mnemonic device to help remember the different endings.
Ide ending is on its own, ate ending for the bigger # of oxygen atoms/subscript
Ite ending for the littler/smaller # of oxygen atoms/subscript

28. Learning Targets
I can determine how two ions will bond to form an ionic compound.

29. Homework
The WS I gave you is NOT HW!!!
Chill out