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NUMBERS NAMING & NET CHAPTER 2 | 1

Chapter 2 OXIDATION NUMBERS, NAMING, NET IONICS

I. Oxidation Number: · Used to keep track of · Shows the general distribution of electrons NOT absolute charge! · Oxidation numbers will typically fall in the range: ─4 to +8, and can include zero!

SPECIES RULE Uncombined elements in OXIDATION NUMBER IS ZERO natural state. Ex: Fe(s), O2(g), P(s) Binary Ionic Compounds OXIDATION NUMBER = IONIC CHARGE Ex. Fe3+ = +3, S─2 = ─2 Metals in Compounds FOLLOW THEIR IONIC CHARGE. Gr 1 = +1, Gr 2 = +2, Gr 3 = 2+ + 2+ Ex. FeCl2, Fe = +2 +3, Zn , Ag , Cd ALWAYS ─1 +1 EXCEPT WHEN WITH METALS Ex. NH3 H = +1, CaH2 H = ─1 Oxygen ─2 EXCEPT: PEROXIDES = ─1 (If Oxygen is with a Gr. 1 or Gr 2 Ex: Na2O O = ─2, NaO O = ─1 element CALCULATE OXYGEN! EXCEPT: OF2 = +2 (F must be ─, therefore O = +2) Covalent/molecular ASSIGN THE MOST ELECTRONEGATIVE ELEMENT ITS MOST NEGATIVE VALUE ACC. TO THE PERIODIC TABLE. CALCULATE THE OTHER (Gr. 17 = ─1, Gr. 16 = ─2, Gr. 15 = ─3) The sum of oxidation numbers in a neutral compound is zero. CaCl2: (+2) + 2(-1) = 0 The sum of oxidation numbers for a polyatomic equals the charge -3 on the . PO4 : (+5) + 4(-2) = -3

Let's do it 1. Assign oxidation numbers to each of the elements in the following compounds and

Substance Element one Element two Element three

H2

Zn

Rb2S

Cr3P4

Ba(NO3)2

─ H2PO3

PBr5

PCl3

NUMBERS NAMING & NET CHAPTER 2 | 2 Substance Element one Element two Element three

ICl3

SO2

─ NO3

H2O2

B2H6

NO3

KMnO4

II. Don’t forget your diatomics!

There are 7 elements that cannot exists as single and must be written as two atoms together when they are pure and not combined into compounds.

These elements are called Diatomic Elements and they must be memorized.

H2 hydrogen gas, pure hydrogen

O2 oxygen gas, pure oxygen, oxygen

N2 nitrogen gas, pure nitrogen, nitrogen I Bring Clay For Our

F2 fluorine gas, pure fluorine, fluorine New House

Cl2 chlorine gas, pure chlorine, chlorine

Br2 bromine, pure bromine, bromine

I2 iodine , solid iodine, pure iodine, iodine

NUMBERS NAMING & NET CHAPTER 2 | 3 IONIC COMPOUNDS

One of the most important concepts to understand when naming and writing ionic compounds is the concept of charge. You must be able to predict which ions will have which charge. Therefore, let’s have a little review!

Let's do it 2. Fill in the charge (or charges) for each of the following groups.

Group Metal or Non-Metal? Lose or Gain Electrons? Charge(s)?

Group 1 Metal Lose 1

Group 2 Metal Lose 2 electrons

Group 13 Metal Lose 3 electrons

Non-Metal Gain 4 electrons Group 14 Metal Lose 2 or 4 electrons

Non-Metal Gain 3 electrons Group 15 Metal Lose 3 or 5 electrons

Group 16 Non-Metal Gain 2 electrons

Group 17 Non-Metal Gain 1 electron

In addition to the Gr 1 and Gr 2 metals, you need to memorize six others with only one oxidation number. Fortunately they form a nice pattern on the periodic table. Notice the stair-step: One element (silver) with +1, two (zinc & cadmium) with +2, and three (aluminum, gallium, and indium) with +3. ASSUME ALL OTHER METALS NEED A ROMAN NUMERAL IN THE NAME TO INDICATE THE ! Note necessarily true, but will work for all AP/IB questions.

A. Binary Ionic Compounds

Now that you know which elements have which charge, we are ready to start writing formulas of ionic compounds! Remember from our last unit that we represent ionic compounds as the lowest whole number ratio of atoms that make a NEUTRAL ! REMEMBER: The goal of ionic compounds is to make electrons lost equal electrons gained – or, in other words, TO MAKE CHARGES CANCEL OUT. Therefore, you might need more than one of an element to make them cancel!

NUMBERS NAMING & NET CHAPTER 2 | 4

Let's do it 3. Notice no charges Remember “-ide” means single are shown in the final element anion EXCEPT compound! and cyanide

Elements Ions Make Charges Cancel! Final Formula

! +! + ! −! = ! +2 -1 Ca and Cl Least common multiple = 2 CaCl2 (need 1 Ca +2 and 2 Cl -1) ! +! + ! −! = ! Al2O3 +3 -2 Aluminum Al and O LCM = 6 (need 2 Al +3 and 3 O -2)

Magnesium phosphide

Iron (III) oxide

Lead (IV) oxide

Tin (II) phosphide

REMEMBER: CATIONS (metals) GO FIRST AND ANIONS (non-metals) GO SECOND! We always write the positive ions first and the negative ions second. We name in the same order as well. For BINARY (2-element) ionic compounds, simply name the metal, then name the non-metal with an “– ide” ending. Metals that have more than one oxidation number possible need a roman numeral to indicate the charge.

Let's do it 4. Name the following binary ionic compounds. If it is a metal with more than one oxidation number, you will have to do some quick algebra to determine the roman numeral.

Cation Name Non-Metal Ion Find Charge of cation Formula charge if (use roman numeral for with Charge if multiple possibilities firm charge!) -1 ! +! + ! −! = ! PbCl2 Cl x = +2 Lead (II) chloride

Gr. 2 -1 BaI2 I NA Barium iodide always +2

Cu2S

Ag N 3

SnO

CoF 3

CdBr 2

NUMBERS NAMING & NET CHAPTER 2 | 5 B. Ternary Ionic Compounds USING POLYATOMIC IONS!

Finally, we need to learn how to name and write compounds that have polyatomic ions as part of the formula. Luckily for us, this is not much different than what we’ve already learned! By the way, these are called TERNARY ionic compounds because they have three or more elements. With formula writing, the process is the same as it was with binary ionic compounds – the only difference is that if you need more than one of a polyatomic ion, we use parentheses before we put the final subscript. Notice parentheses were used when Let's do it 5. Find the final formula for each of the following compounds. more than one polyatomic ion was needed!

Compound Ions Make Charges Cancel! Final Formula

+3 -1 ! +! + ! −! = ! Aluminum nitrate Al and NO3 +3 -1 Al(NO3)3 (need 1 Al and 3 NO3 )

Ammonium phosphate

Calcium hydroxide

Copper (II)

Zinc chlorate

Naming Flowchart to help with Ionic Compound naming:

Ionic Compound

Metal + nonmetal Metal + poly ion + NH4+ + nonmetal NH4 + poly ion

1. Name metal or 1. Name metal or ammonium 2. Name non-metal with 2. Name polyatomic ion "-ide" ending

If metal has more than one possible oxidation number, show using rom an numerals (Group 3-15 metals, except Al +3, Zn+2 , Ag+1 , & Cd +2)

NUMBERS NAMING & NET CHAPTER 2 | 6 NOTICE HOW IMPORTANT SPELLING IS!!!!! Ammonium Chloride: NH4Cl

Ammonium Chlorate: NH4ClO3

Ammonium Chlorite: NH4ClO2 -ATE and –ITE endings indicate POLYATOMIC IONS.

-IDE endings indicate non-metal ions, except cyanide and hydroxide.

C. Ionic Compounds that are

Hydrates are ionic compounds that have trapped # PREFIX # Waters PREFIX in their structure. They have different 1 mono- 6 hexa- physical properties from the anhydrate form. 2 di- 7 hepta- Use covalent prefixes to indicate the number of 3 tri- 8 octa- waters that are bound when writing the name. Use 4 tetra- 9 nona- a “•” (dot) to show how many waters are bound 5 penta- 10 deca- when writing the formula.

Let's do it 6.

FORMULA NAME Cu2SO3•3H2O Copper (I) sulfite trihydrate (note that you need to include because copper has more than one oxidation number) Na2SO4• 10H2O

LiNO3•3H2O

NAME FORMULA Zinc sulfate heptahydrate ZnSO4•7H2O Cobalt (II) tetrahydrate

COVALENT

A. INORGANIC Covalent Molecules

A binary covalent compound is composed of two different nonmetal elements. For example, a of chlorine trifluoride, ClF3 contains 1 of chlorine and 3 atoms of fluorine.

STEPS TO NAMING:

Step 1: Name the first non-metal using a prefix to indicate how many of that element was in the compound. Step 2: Name the second non-metal using a prefix and add the suffix -ide to the end of the name of the element. NUMBERS NAMING & NET CHAPTER 2 | 7

# ATOMS PREFIX # ATOMS PREFIX 1 mono- 6 hexa- 2 di- 7 hepta- 3 tri- 8 octa- 4 tetra- 9 nona- 5 penta- 10 deca-

EXCEPTION: If the compound contains one atom of the element that is written first in the name, the prefix "mono-" is not used.

Note: when the addition of the Greek prefix places two vowels adjacent to one another, the "o" at the end of the Greek prefix is usually dropped; "monooxide" would be written as "monoxide". The "i" at the end of the prefixes "di-" and "tri-" are never dropped.

Let's do it 7.

FORMULA NAME CO 2

N2 O3

SO3

NAME FORMULA

Dinitrogen pentoxide N2O5 Chlorine monofluoride

Nitrogen trifluoride

ACIDS

We tend to place into their own category, as they follow their own set of rules. Acids are always aqueous (aq). An example is HCl (aq) – named hydrochloric . If the substance is a gas (such as HCl (g)), then use your ionic naming rules – i.e. hydrogen chloride. 1. Naming Acids: • Name the anion present in the acid. • Change the suffix of the anion according the chart below • Add the word “acid” as a last name. • If the acid has sulfur or phosphorus, we add the “ur” and “or” back into the name. Ex. H2SO3  sulfite  sulfurous acid NUMBERS NAMING & NET CHAPTER 2 | 8 Let's do it 8.

Formula of Acid Anion Name Acid Name I “ate” something “icky”, all “nite” I H + poly ion (per-ate) Per_____ate Per_____ic acid was nauseous, when H + poly ion (-ate) _____ate _____ic acid I took a ride on a H + poly ion (-ite) _____ite _____ous acid hydraulic plane H + poly ion (hypo-ite) Hypo____ite Hypo___ous acid H + non-metal _____ide Hydro___ic acid

FORMULA ANION NAME ACID NAME

HNO2 (aq) nitrite Nitrous acid HF (aq)

H3PO3 (aq)

HIO4 (aq)

2. Writing Acid Formulas: • Change the acid suffix to an anion suffix • Write the anion formula • Add enough “H+” to balance out the charge on the anion

Let's do it 9.

ACID NAME ANION NAME ANION FORMULA ACID FORMULA Add H+ to anion 3- Phosphoric acid “-ic” came from “-ate” PO4 H3PO4 “Phosphate” Perbromic acid

Acetic acid

III. COMMON ORGANIC MOLECULES – In class activity

NUMBERS NAMING & NET CHAPTER 2 | 9 IV. NET IONIC EQUATIONS - Net ionic equations are streamlined to show only the species involved in the chemical change. In order to do these you need to re-memorize your strong acids, strong bases, and rules.

SOLUBLE COMPOUNDS EXCEPTIONS All Group 1 None + All ammonium (NH4 ) salts None − − − − All NO3 , ClO3 , ClO4 , and C2H3O2 salts None − − − + 2+ 2+ All Cl , Br , I salts Ag , Hg2 (mercury (I)), Pb All F− salts Mg2+ Ca2+, Sr2+, Ba2and Pb2+ 2− 2+ 2+ 2+ 2+, + 2+ All salts of SO4 Ca , Sr , Ba , Pb Ag , Hg2 Group 1 & 2 and NH4OH All other hydroxides

Strong Bases: dissociate 100% in water. All hydroxides of group I and II* except beryllium and magnesium.

* Completely dissociated in solutions of 0.01 M or less. These are insoluble bases which ionize 100%. The other five in the list can easily make solutions of 1.0 M and are 100% dissociated at that concentration.

Strong Acids 1 - dissociate 100% in water

Type Formula Hydrogen halides (aq) HCl HBr HI Oxyacids of HClO3 HClO4 HIO4 st + + ─ Sulfuric (1 H only!!) H2SO4  H + HSO4

Nitric Acid HNO3

There are three steps, although with practice you can skip one.

Complete 2AgNO3 (aq) + Na2CrO4 (aq)  2NaNO3(aq) + Ag2CrO4 (s) Formula + ─ + 2─ + ─ Complete 2Ag + 2NO3 + 2Na + CrO4  2Na + 2NO3 + Ag2CrO4 (s) Ionic

+ 2− ─ + Net Ionic 2Ag + CrO4  Ag2CrO4 Spectator Ions: NO3 & Na

Let's do it 10. Aqueous solutions of lead (II) nitrate and iodide are combined to form lead (II) iodide and potassium nitrate.

Complete Formula Complete Ionic Net Ionic Spectator Ions:

1 http://www.science.uwaterloo.ca/~cchieh/cact/c123/stacids.html NUMBERS NAMING & NET CHAPTER 2 | 10 Let's do it 11. Chlorine gas is bubbled through a of potassium bromide to form bromine and potassium chloride.

Complete Formula Complete Ionic Net Ionic Spectator Ions: