CHM 1032C   Tentative Grading Outline   Fall 2015


Chapter 3 Ionic Compounds Homework Packet

A. _____ (03) e-1 Configuration of Ions-lecture (3.1-3.3)  Answers                              

B. _____ (02) Periodic Ionic Character-Section 3.2 Answers

C._____  (02) Bond Recognition/Compound Classification-Sections 3.3, 4.1Answers

D  _____(02) Binary Ionic Compounds-Section 3.9  Answers                                              

E. _____(05) Polyatomic Ions-Section – section 3.8  Answers                                             

F. _____(04) Ternary Ionic Compounds-Section 3.9 Answers                                 

G. _____(02) Binary Acids/ Ternary Oxyacids-Section 3.11 Answers         

_______(20) Chapter 3 Total                                                                             


Chapter Three: Part A:   Electron Configuration of Ions    3 points

Given the following ions, use arrows to fill-in the electron configuration of the ion, then rewrite the configuration into the chemist’s shorthand:

 1.     Cl1-  ion      Chemist Shorthand: ___________________________




2.  K1+ ion  Chemist Shorthand: _____________________________


3.  H1+ ion  Chemist Shorthand: _____________________________

4.  H1-ion  Chemist Shorthand: _____________________________

Chapter Three: Part B     Periodic Ionic Properties      2 points

Using a periodic chart, write the ionic character (monoatomic ionic charge) of the following elements: (The number before the element is its atomic number)


1.  19 K     ________                  6.    9F      _____


2.  20Ca    _______                    7.    1H      _____    _____  


3.  7N        _______                   8.    16S     _____


4.  17Cl      _______                   9.    10Ne   _____


5.  53I       ______                   10.   15P     _____      


 Chapter Three: Part C     Bond Recognition            2 Points

Read the short discussion in Sections 3.1 and 4.1 on the difference between Ionic and covalent bonding.

There are three types of chemical bonds:
 Ionic, Covalent, and Metallic

There is a simpler way to predict if two atoms will transfer their electrons or share their electron in pairs making a compound. Read about the Pauling’s Scale of Electronegativity in Section 4.9. Corwin Figure 12.9 shows the electronegativity of each element on the periodic chart. This table will be needed in Chapter 4 Part II Bond Polarity.

 If the difference in electronegativity between two atoms is greater than 1.8 (Corwin), the electrons will transfer from one atom to the other to make ions and Ionic Compounds. Ionic (sometimes called Electrovalent) Compounds are also called salts and in nature they are called minerals and in Sports medicine Body Electrolytes. We will over simplify this concept to say if a metal meets a nonmetal ionic bonds are formed (Just a Rule of Thumb)(if a table of electronegativity is not included). Hein (14th) states on page 227 if the difference in electronegativity is greater than 2.0 the bonding is strongly ionic, while less than 1.5 strongly covalent. Then he states between 1.7-1.9 the bonding will be more ionic than covalent.

 For this course, if the difference between the electronegativity of two atoms is less than 1.7 then the two atoms will share electrons in pairs. Two types of sharing bonds are formed. Metallic and Covalent.

 Metallic Bonds are formed when two metals share electrons such as alloys of metals. 24 karat gold is pure gold and is very soft. But Jewelry is usually 10-18 Karat Gold, meaning that another metal is mixed with gold to make the solid harder. We will not study Metallic Bonds in this course, but you should know that two metals share electrons in pairs to make Metallic Bonds.

Metallic bonding occurs as a result of electromagnetism and describes the electrostatic attractive force that occurs between conduction electrons (in the form of an electron cloud of delocalized electrons) and positively charged metal ions. It may be described as the sharing of free electrons among a lattice of positively charged ions (cations). In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities.[citation needed] Metallic bonding accounts for many physical properties of metals, such as strength, ductility, thermal and electrical resistivity and conductivity, opacity, and luster.[1][2][3][4]

Metallic bonding is not the only type of chemical bonding a metal can exhibit, even as a pure substance. For example, elemental gallium consists of covalently-bound pairs of atoms in both liquid and solid state—these pairs form a crystal lattice with metallic bonding between them. Another example of a metal–metal covalent bond is mercurous ion (Hg2+2).“

 Covalent Bonds are formed when two nonmetals bond together. The elements carbon, oxygen, hydrogen, sulfur, nitrogen, phosphorus, chlorine, and bromine will be the main nonmetals studied in drawing dot structures of molecules. Bonds between these nonmetals are always Covalent.

Part C of Chapter 3 should now be easy. Predict what type of bond will be made if two atoms combine:

In General:

Metal-Metal = Metallic Bond  (example: Ag(5)-Au(14)-Cu(5) = 14 Karat Gold)

Metal-Nonmetal = Ionic Bond (example: Na-Cl)

Nonmetal-nonmetal = Covalent Bond (example: H2O)


Chapter Three: Part C  Sample   Bond Recognition   2 points

Using a periodic chart (Rule of Thumb), predict the bond that would form between the two elements:


1.   Fe-Al            ________________


2.    P-S                   ________________


3.    C-O               ________________


4.    B-Cl                ________________


5.    K-I                ________________


For the following element pairs use the electronegativity table below to determine if the bond is ionic or covalent.


6.    Na-P               ________________   9.    P-H                  ________________


7.    Ca-Br              ________________  10.   Be-Cl                ________________


8.    Ge-O               ________________



       Most Common Ionic Charges for Monatomic Ions


The element written first in either the name or the formula is a metal. 

The element written second is a nonmetal. 

 Salts are metallic and nonmetallic ionic compounds. 

 There are no molecules of salts-just macro ionic lattices. 

 Name the metallic element. 

If the metallic element has more than one ionic state, write a ROMAN NUMERAL after the element’s name (In Parenthesis) to indicate which charge state the metallic element is using to form the compound.

 Drop the suffix off the nonmetal’s name and add -ide which indicates the salt is binary (exceptions: cyanide & hydroxide which are polyatomic ions).

No prefixes are used to indicate how many atoms are present in the formula. 


NaCl                Sodium Chloride (table salt)

 Al2O3               Aluminum oxide

 FeS                  Iron(II) sulfide (Note: No space between the metal and the parenthesis)

 Fe2O              Iron(III) oxide (rust)


To write the formula from the name of the salt use the following procedure:

 (a) Write the symbols (or formulas for radicals) of the ions represented
For Example: 
Calcium nitride

 (a)                                Ca          N

(b)  Use the periodic chart to write the ion charge of each element (or polyatomic ion) as superscripts: 

                           Ca+2            N-3

  (c ) Find the L.C.M. (Least common multiple) of the positive and negative charge.

 The LCM is the smallest number that both charges will decide into evenly.  The LCM is  the total electrons transferred.  Therefore, it represents the total  positive charge created by the metallic ions and the total negative charge created by the nonmetallic ions.  This may be proved by drawing the dot structure of the compound showing all electrons transferred.

 The LCM of +2 and -3 is 6,   therefore 6 e-1 are transferred creating a total positive charge of +6, and the total negative charge of -6

         --> 6e-1-->

 (d) Divide the LCM by the positive charge, this dividend will represent the subscript behind the metallic ion in the formula.

+6 divided by +2 = 3; therefore half of the formula is:    Ca3Nx


 (e)  Divide the LCM by the negative charge, this dividend will represent the number of nonmetallic ions in the formula.

-6 divided by -3 = 2; therefore the other half of the formula is:   Ca3N2          


Example:           Potassium phosphide

 Write Symbols and the Charges:

                   K+1     P -3

   LCM:    3

        Balance the chemical formula:


 In addition to working the sample tests, you must practice on writing the names and formulas for Ionic Compounds. He following are online homework for 2 points each:


D. Binary Ionic Names:


D1. Binary Ionic Formulas:


Submit grades on separate grading Sheet when taking after taking Chapter 4 exam


Chapter Three:  Part D   Binary Ionic Compounds     2 points

Using a periodic chart, write the names or the balanced formulas for the following compounds depending on whether the formula or the name is given:


1.   Copper II phosphide                _________ (Cupric phosphide)


2.    Iron III Oxide (rust)                  _________  (Ferric Oxide)


3.    Lead IV sulfide                        _________  (Plumbic sulfide)


4.    Sodium chloride                      _________


5.    Tin II fluoride (in toothpaste)   _________  (Stannous Fluoride)


6.   MgCl2        ________________________


7.    NiF2         ________________________


8.   K3N          ________________________


9.   Al2O3       ________________________


10. CuBr         ________________________


Online Study Guide:


Chapter 3: PART E:   Polyatomic Ions


Ion Flowchart




You can predict the Monoatomic Anions or Cations by the position the element resides on the periodic chart. If the ion come from a Representative Element (IA-VIIIA) there is one and only one ionic charge. If cation is a transitional metal with several different charges you have to rely on the Name

                              Periodic Table of Selected Ions


Note the charges for groups IA, IIA, IIIA, VA, VIA, VIIIA. From book to book, the charges on the transitional metals will vary


By now you should have practiced: C-3 Part A, then try C-3 Part D and write the formulas for Binary Ionic Compounds.

Almost all chemistry textbooks have sections dedicated to polyatomic ions and include a list of common ions.


What is a polyatomic ion?


A group of atoms bound together (covalent bonds) that bears an overall negative or positive charge.


Corwin (7th) suggests that you use flash cards listing the name on one side and the formula with its charge on the other to aide your memorization of these formulas. Most chemistry teachers require you to know some of the common polyatomic ions by the end of the course whether it is from repetition of use with a help table or from memory from the first day of introduction. Below are tables from various chemistry books used:


Polyatomic Ion Charts from Textbooks
2045 McMurray:
Table 3.2       1025  Corwin: Table 7.03    1032 McMurry GOB: Table 3.
2045 Silverberg:
Table 2.5       1020 Tillery:  Table 9.3       1025 Hein  Table 6.5
2045 Kotz:           
Table 3.1       1020  Hill:      Table 5.04



Here is a sample polyatomic ion table:


Hill’s Table 5.4 (and Hill suggest for you to memorize the entire table):

 After you start memorizing, during the course the formulas may be swimming in your head and the charges too. To write balance Ternary Ionic Compounds, you must be able to write the formula and the charge of each polyatomic ion required.


Corwin suggests there is only one (Hill has two) common polyatomic Cation(s) and both end in –ium suffix. He notes most of the Anions have an –ate suffix, while a few have –ite, and two have –ide in their name. How do we accomplish this list?










Our McMurry GOB text suggests are students should learn the common polyatomic ions in Chapter 3 Section 3.8 Table 3.3:




Knowing dot structures of polyatomic ions McMurry GOB Section 4.7) (Corwin Chapter 12 section 12.5), and some keen observations you can boil it down to six questions:


1.   What is the formula for the –ate polyatomic ion?

2.   What is the charge on –ate polyatomic ion?

3.   What happens when you attach hydrogen atom(s) to the polyatomic

     2- and 3- anions?

4. What does –ite mean?

5.   How do the hypo- and per- prefixes apply to polyatomic ions?

6.   What are the two –ide polyatomic ions and two -ium positive Anions?













Your First task is to memorize the formulas and the charges for the polyatomic ions in your text book (Table 3.3) for a short test:


Progressive Polyatomic Ions McMuury GOB (1 point) 



Formula with charge

Acetate ion



Ammonium ion



Carbonate ion




Carbonate ion


Chromate ion



Dichromate ion



Cyanide ion



Hydroxide ion



Hypochlorite ion


Nitrate ion



Nitrite ion



Oxalate ion


Permanganate ion



Phosphate ion



Phosphate ion



Phosphate ion


Sulfate ion



Hydrogen sulfate ion


Sulfite ion







Progressive Polyatomic Ions Corwin (1 point)) 

Write the formula and the charge for the following polyatomic ions: Corwin(Table 6.3) 1 point if your text is Corwin.        


Formula with charge


























Hydrogen sulfate          































Progressive Polyatomic Ions Hein (1 point) 

Write the formula and the charge for the following polyatomic ions: Hein(Table 6.5 page 109) 1 point if Hein Text              


Formula with charge




























Hydrogen sulfate         
























So: it is time for you to discover, what I saw over 50 years ago. It is not in any textbook. The books just say know or memorize these tables. Go to:


When you go to the site above (which looks like the image below), click on the X for each polyatomic ion and note if the # of oxygen atoms are three or four in the formula.



To expose the threes and the fours in the lower left hand corner (Taylor’s ¾ rule) click the numbers 0,1…8,9 Border three rule, then 1,2..5,6 in the box of six rule. Also do the 0,1…7,8 Transitional O4 Rule.


Taylor’s ¾ rule is summarized at:





Then do the same for the box just to the right of Taylor’s ¾ Rule, and discover Taylor’s Charge Rule.


Taylor’s Charge Rule is summarized at:



The story behind how your instructor related the periodic table to a long list of polyions, read the abstract for his talk at 2YC3:


Now comes the big task!


You may either memorize 55 polyatomic ions or learn to read the periodic table with six rules and be able to write formulas and the charges for the required 1025/1032/2045 list:





Either make a hard copy set of polyatomic ion flash cards or practice the 65 polyatomic ions Flash Card web site for 2 points  at:



Chapter Three: Part E    Polyatomic Ions                 2 points


Using a periodic chart write the names or formulas of the following polyatomic ions depending on whether the formula or name is given:


1.   CO32-       _____________________


2.   SO32-       _____________________


3.   PO33-     _____________________


4.   ClO31-     _____________________


5.   NO31-     _____________________


6.   Hydroxide          ________


7.   Ammonium        ________


8.   Hypochlorite       ________


9.   Nitrite                  ________


10.  Phosphate          ________


Textbook (Corwin 7th) Reference: Chapter 6 Section 6.3 Table 6.3 Optional End of Chapter p184 #13-18

Hein (14th): Section 6.5 Memorize the formulas and charges of Hein Table 6.5 (18 ions)


Online Required  Homework (2 Points Each):


E: Polyatomic Ion Names Homework:   

E1. Polyatomic Ion Formulas:


Submit grades on separate grading Sheet (Goldenrod) when taking Chapter 4 exam or download the form from:








CHM 1032C Chapter 3 Homework Packet

In chemistry, a ternary compound is a compound containing three different elements. An example of this is sodium phosphate, Na3PO4. The sodium ion has a charge of 1+ and the phosphate ion has a charge of 3-. Therefore, three sodium ions are needed to balance the charge of one phosphate ion. Another example of a ternary compound is calcium carbonate. In naming and writing the formulas for ternary compounds, we follow rules that are similar to binary compounds.(CaCO3).


Ste that uses least common multiple balance method:


Sites (You-tubes) that use the crossing method(UGH):


Another You-Tube:


Chapter 3: Part F    Ternary Ionic Compounds        2 points

Using a periodic chart write the names or formulas of the following compounds depending on whether the formula or name is given:


1.   Na2CO3       _____________________


2.   K2SO4           _____________________


3.   (NH4)3PO4    _____________________


4.   Ca(ClO3)2     _____________________


5.   CuNO3            _____________________


6.   Aluminum Hydroxide          ____________


7.   Ammonium carbonate        ____________


8.   Sodium Hypochlorite          ____________


9.   Magnesium Nitrate              ____________


10.  Iron III sulfite                       _____________


McMurry GOB: Sections 3.9 & 3.10/Corwin Text Sections 6.4 & 6.6


Online Homework (2 Points Each Required):


F: Ternary Ionic Compound Names Homework:    


F1. Ternary Ionic Compound Formulas:


Submit grades on separate grading Sheet when taking Chapter-4 Exam



Chapter 3 Part G: Binary/Ternary Acids

What is an acid?

A substance that releases hydrogen ions (H+) when dissolved in water. Inorganic formulas of acids have ionizable hydrogen(s) written first in the formula.

      Strong Acids               Weak Acids

Strong acids ionize 100% in a water solution, while Weak Acids ionize
less than 5% in a water solution

There are Binary/Ternary Acid online homeworks for your practice for M-4 Part G:

G: Binary/Ternary Acid Names:

G1: Binary/Ternary Acid Formulas: 

( Chapter 6 Bishop Sections 6.3-6.4 )give you instructions for naming and writing formulas of acids. );
 (Chapter 6 Corwin 7th covers binary acids in section 6.8; while section 6.9 covers ternary acids.) (Hein 14th covers acids in section 6.6.

A brief tutorial for names and formulas of acids follows:

If hydrogen is written first in a chemical formula, there is two ways to name the compound. As a pure molecular compound or as an aqueous acid:

If the compound is a pure molecular compound then you name it just as if it were an ionic compound:
HCl          hydrogen chloride                                H2CO3      hydrogen carbonate

HClO        hydrogen hypochlorite                       H2SO4      hydrogen sulfate

HClO2      hydrogen chlorite                                 H2SO      hydrogen sulfite

HClO3      hydrogen chlorate                                HC2H3O2   hydrogen acetate 

HClO4      hydrogen Perchlorate                         H2C2O4     hydrogen oxalate

H3PO4      hydrogen phosphate                          HBr          hydrogen bromide

HF            hydrogen fluoride   

Writing hydrogen first in a chemical formula indicates that when you dissolve the compound in water, a water molecule has the ability to pull the hydrogen off  (from strong electronegative elements like oxygen)  the molecule HXO3 and creating hydronium ions, H3O1+ and  a negative ion XO31- (cation).

The way you indicate this ionic solution is to write the formula followed by (aq) meaning a water solution:  HXO3 (aq) .

The first step is to drop the first word hydrogen and
add a second word

HCl          hydrogen chloride acid (aq)

HClO        hydrogen hypochlorite acid (aq)

HClO2      hydrogen chlorite acid (aq)

HClO3      hydrogen chlorate acid (aq)

HClO4      hydrogen perchlorate acid (aq)

H3PO4     hydrogen phosphate acid (aq)

H2CO3     hydrogen carbonate acid (aq)

H2SO4     hydrogen sulfate acid (aq)

H2SO3     hydrogen sulfite acid (aq)

HC2H3O2   hydrogen acetate acid (aq)

H2C2O4    hydrogen oxalate acid (aq)

HBr          hydrogen bromide acid (aq)

HF            hydrogen fluoride acid (aq)

  The next step is to drop the suffix from the cation and make the following substitution with another suffix:

Change the -ate to -ic

Change the -ite to -ous

but the instead of coming up with a third suffix for -ide , they reused the -ic for -ide and added a prefix hydro- (Do not get this confused with the prefix hypo- which means 'under'.)


HCl          hydrochloric  acid (aq)

HClO        hypochlorous acid (aq)

HClO2      chlorous acid (aq)

HClO3      chloric  acid (aq)

HClO4      perchloric  acid (aq)


H3PO4     phosphoric  acid  (aq) (Put the -or- syllable back in the name)


H2CO3     carbonic  acid (aq)


H2SO4     sulfuric  acid  (aq) (Put the -ur- syllable back in the name)

H2SO3     sulfurous acid (aq) (Put the -ur- syllable back in the name)


HC2H3O2   acetic  acid (aq)  (Notice the three hydrogens written after carbon are NOT ionizable and not written first in the formula)


H2C2O4    oxalic acid (aq)


HBr          hydrobromic acid (aq)


HF            hydrofluoric acid (aq)

On Corwin 7th page 185 Questions 49-56 will give you more practice on writing names and formulas of acids.

At the end of chapter 6 Hein 14th exercises 17, 18, 19, and 20 pages 116-117 are additional acid nomenclature problems.

 McMurry GOB introduces acids in Section 3.11, then goes into acids but does  teach how to name in Chapter 10.

Chapter 3: Part G    Binary/Ternary Acids                 2 points

Using a periodic chart write the names or formulas of the following compounds depending on whether the formula or name is given:


1.    HCl       _____________________


2.    H2SO4     ____________________        


3.    HNO3      _____________________


4.  HNO2        ___________________


5.   H2CO3      ___________________


6.    Hypochlorous acid          _________


7.    Phosphoric acid              _________


8.    Sulfurous acid                 _________


9.    Perchloric acid                _________


10.  Hydrofluoric acid              ________



Corwin 7th Text Sections 6.8 and 6.9

Optional Additional Homework: p 185 Q #49-56


Online Homework (2 Points Each Required):


G: Binary/Ternary Acid Names Homework:     


G1. Binary/Ternary Acid Formulas:


Submit grades on separate grading Sheet when taking Chapter-4 Exam


Online Study Guide: