CHM 1025C Module Four Assignment Outline

Module Four covers four parts of Chapter 12 and all of Chapter 7. If you use the grading outline, the key reference sections are listed under each part:

Module Four Part I: Answers Language of Chemistry/Chemical Bonds (Chapter 12, 7)

A. _____(05) Bond Recognition/Compound Classification-Sections 7.1,12.1-.3,12.6 Answers acd

B. _____(30) Dot Structures of Molecules-Section 12.4, 12.5 Answers                                

C. _____(05) Binary Molecular(Covalent) Compounds-Section 7.7 Answers acd

D  _____(10) Binary Ionic Compounds-Section 7.2, 7.5  Answers acd

E. _____(10) Polyatomic Ions-Section – section 7.3, 7.6  Answers e

F. _____(10) Ternary Ionic Compounds-Section7.3, 7.6  Answers f                                      

G. _____(05) Binary Acids/ Ternary Oxyacids-Section 7.8,7.9 Answers g    

H. _____(12) Inorganic Compounds 7.2-7.9 Answers h                                                                                                                       

M._____ (20) Multiple Choice Chapter 12, 7 (answers at bottom)                                         

_______(107) Total = ______%

Now that you have completed Module Three Part F and can draw the Dot Structure of Any Representative Element on the Periodic Chart, it is time to make Ionic or Covalent Compounds.

 Part A: Bond Recognition

Read the short discussion in sections 12.1-12.3 on pages 310-317 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. Skip back to page 325. Read about the Pauling’s Scale of Electronegativities. Figure 12.9 shows the electronnegativity of each element on the periodic chart. This table will be needed in Module Four Part II Bond Polarity.

 If the difference in electronegativity between two atoms is greater than 1.7, 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. We will over simplify this concept to say if a metal meets a nonmetal ionic bonds are formed (if a table of electronegativity is not included).

 If the difference between the electronegativites 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.

 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 A of Module Four should now be easy. Predict what type of bond will be made if two atoms combine:

In General:

Metal-Metal = Metallic Bond

Metal-Nonmetal = Ionic Bond

Nonmetal-nonmetal = Covalent Bond

 

Part B: Dot Structures of Covalent Compounds

Section 12.4, pages 318-322 and Section 12.5 pages 322-324 are critical for you to be able to do Part B: Dot Structures of Covalent Compounds. For a take home Laboratory Assignment, you are to draw the dot structures for all the Compounds listed on Module Four Part B. You should alternate using colored pencils (pens) when showing the bonding between the two atoms. Each connected atom should be a different color (all oxygen atoms one color, hydrogens another, etc. We will only draw the dot structures of covalent compounds which follow the octet rule (duet rule for hydrogen) for this section. Page 322 gives a set of guidelines for drawing the dot structures of Covalent Molecules. There is an interactive drag and drop web site for these molecules. Draw the following dot structures of the 30+ molecules below: 

 

1.  NH3   CH4    H2O2     H2O    (all single bonds)  
2.  H2SO4   H3PO4   HClO4   HClO3  (all single bonds)
3.  HNO3   H2CO3  HNO2  (contains one double bond)             
4.  CO2    HCN    CO   SO3  SO2,  SiO2 (contains at least one multiple bond)
5.   HC2H3O2     H2C2O4  HCHO2    CH2O (carbon to carbon by single covalent bond in first two)        
6.  C2H4  C2H2   C3H8   C4H10    C2H6  (bond carbons to carbon BUT MAY HAVE A MULTIPLE BOND)
7.   CH3CH2OH      CH3COCH3 (carbon to carbons by single covalent bonds)
8.   CH3OCH3   (Oxygen separates the carbons)  CHONH2   (O & N both bond to C)
9.  CH2NH2COOH (carbon to carbons by single covalent bonds (-NH2 hooks to 1st carbon in#1)  
10. CH3COOCH2CH3        (-CH2CH3 also hooks to oxygen in#2)

 The dot structure of molecules interactive drag and drop web site menu page is:
http://www.lsua.us/chem1001/dragdrop/menu.html

Guidelines from Instructor’s Lecture:

  1. Place the nonmetal which is not oxygen or hydrogen in the center of the structure.
  2. First using all single bonds, hook all oxygens in the formula to the central nonmetal (Simple covalent or coordinate covalent bonds).
  3. Never hook oxygen to oxygen except in peroxides.
  4. If oxygen is present, hook the hydrogen written first in the chemical formula to an oxygen. Hydrogen requires only two electrons to fill it's orbital. If hydrogen is written second in the formula after another nonmetal, then hook that hydrogen to that nonmetal, not to oxygen such as:
     (CH3-COOH written organically) acetic acid HC2H3O2 (written inorganically)
    or oxalic acid H2C2O4 both hydrogens are attached to the oxygens.
  5. If using all single bonds connecting the oxygen to the central nonmetal, the count of electrons around each element should total eight electrons (octet rule). This includes the element's original outer surface (valence) electrons plus the electrons being shared from the bonded element. (in CHM 2045C we will investigate the Rules of six, ten and twelve, but for our course everything will follow the octet rule, except hydrogen (duet)
  6. If the count is 7-7, then add a second bond, a double covalent bond (four electrons being shared between two atoms.
  7.  If the count is 6-6, then make a triple covalent bond between the two elements (six electrons shared between the two atoms). (Try HCN, N2, and C2H2)

  8. If the count is 8-6, 9-7, then make a coordinate covalent bond by shifting the electrons. Hook the element with six electrons (usually oxygen) by creating a vacant orbital onto an unshared pair of the element with eight. A coordinate covalent bond is still a single bond. In making double, triple bonds you may use one or two coordinate covalent bonds to predict a structure using the octet rule.
  9. Extras: Never have an unshared pair or lone pair of electrons on a carbon atom (except carbon monoxide, CO). These are two dimensional structures, so there are many variations of the answer shown on the web site. Never have more than two bonds to any oxygen (except CO). If you place a hydrogen to an oxygen, then oxygen HAS to hook to another element by a single bond, never a double bond.

Part E: Polyatomic Ions 

Read carefully Section 7.3 on Polyatomic Ions.

The web sites for Polyatomic Ions are on your home page. Go to Polyatomic ion study guide:
http://www.hccfl.edu/faculty/john_taylor/chm1025/polyions/polyionstudyguide.html

 It goes through my rules for polyatomic ions and has links to the other pages. However, if you chose to memorize the polyatomic ions, then you are responsible for all the polyatomic ions on page 173 Table 7.3 plus all nonmetal –ate, -ite ions and the prefixes: Per- and Hypo- on the oxy-halogen ions.

Acetate, Cyanide, Chromate, Dichromate, and Permanganate ions in Table 7.3 are bonus ions which you will be required to know in CHM 2045C. You must also know hydroxide, ammonium, hydrogen phosphate, dihydrogen phosphate, and hydrogen sulfate ions not in Table 7.3. Module 4 Parts E, F,G,H will use these polyatomic ions.

 

 

Parts D & F: Binary/Ternary Ionic Compounds

 For M-4 D&F there are two interactive homework web sites and is covered in section 3.3 of the text:

Binary Ionic: http://www.lsua.us/chem1001/dragdrop/menu.html

Ternary ionic: http://www.lsua.us/chem1001/nomenclature/TernarySalts/ternaryIonic.html

A brief tutorial for Part D: 

PART D:   BINARY  (IONIC) COMPOUNDS

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 mane 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. 

Examples:

NaCl                Sodium Chloride (table salt)

 Al2O3               Aluminum oxide

 FeS                  Iron II sulfide

 Fe2O3               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-->
Ca+2                     N-3

 

(d   (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 Charges:

        Ba+2 O -2                         Cu +                                                                 K+1 P -3

   LCM:    3

         Ba+2 O -2                        Cu +1 S                                      Balance the chemical formula:
       K3P       

 In addition to working the sample tests, you may want to practice on writing the names and formulas for Ionic Compounds. 

On pages 188-9, questions 13 thru 42 are also good practice.

Part G: Binary Ternary Acids

 There is a Binary/Ternary Acid online homework for your practice for M-4 Part G:

http://www.lsua.us/chem1001/nomenclature/Acids/acids.html

Chapter 6 Bishop Sections 6.3-6.4 give you instructions for naming and writing formulas of acids.

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

HClO        hydrogen hypochlorite

HClO2      hydrogen chlorite

HClO3      hydrogen chlorate

HClO4      hydrogen perchlorate

H3PO4     hydrogen phosphate

H2CO3     hydrogen carbonate

H2SO4     hydrogen sulfate

H2SO3     hydrogen sulfite

HC2H3O2   hydrogen acetate 

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 acid: 

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)

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)

On page 190 Questions 47-54 will give you more practice on writing names and formulas of acids.

 

Part C Binary Molecular Compounds

 Binary Molecular compounds are explained after the ionic compounds in Chapter 7 section 7.7, and inorganic acids are not covered till last in the chapter.

 The Online Binary covalent Molecules Homework has a glitch. It freezes after working through several samples. There are 31 total items. Just restart if it freezes.

The web site is:
 http://www.fccj.us/chem1001/nomenclature/Molecules/binaryCovalent.html

 

Here is a brief tutorial for Part C: 

 PART C: BINARY COVALENT COMPOUNDS

 Both elements are nonmetals attached by covalent bonds.  These bonds may be single, double, or triple covalent.  Due to the covalent bonding there are many ratios of the same two elements making many different compounds.  For this reason, the chemist states how many atoms of each element is present in the chemical formula in the formal name of the compound.

Prefixes are attached to each element to indicate how many.  Each student should learn the following prefixes:

 

MONO            =          ONE                                        HEXA              =          SIX

DI                    =          TWO                                       HEPTA            =          SEVEN

TRI                  =          THREE                                   OCTA             =          EIGHT

TETRA            =          FOUR                                      NONA            =          NINE

PENTA            =          FIVE                                        DECA              =          TEN

The element that is shown first in the chemical formula is written first using the proper prefix to indicate how may atoms of that element is contained in the compound.  If there is only one atom of that element it is often found without the prefix mono.  If you leave the prefix off then it is understood that you mean mono.

The element which is written second in the chemical formula is written second in the chemical name, but in addition to the prefix indicating how many, the suffix of the element’s name is changed to -ide. 

 

 carbon becomes carbide                             chlorine becomes chloride

sulfur becomes sulfide                                oxygen becomes oxide

hydrogen becomes hydride                        nitrogen becomes nitride 

Therefore, the following formulas of binary compounds would be spoken:

CCl4                 carbon tetrachloride                                                     

SO2                  sulfur dioxide

CO2                 carbon dioxide

N2O3                dinitrogen trioxide

 BH3                 boron trihydride

We use common names for NH3, and H2O. What would be their correct binary molecular names? Methane, CH4, is the organic name for CH4, what would its inorganic name be?

 For more practice on page 189 try problems 43 thru 46 for binary nonmetal compounds.

 

 

Part H: Inorganic Compounds 

For Part H, we mix the compounds from Parts C, D, F, and G. There is a long interactive homework for Part H at:

http://www.fccj.us/chem1001/nomenclature/Inorganic/inorganic.html 

The key to deciding which system to use in Part H is to look at the element written first.

 1. If a Metal is written first (or a polyatomic ion), then use the rules for ionic compounds (salts).

2. If a nonmetal is written first, then use the Covalent/Molecule System with prefixes. (If the compound is Organic Nomenclature of Organics is covered in Chapter 11, but for now use the prefix system of binary molecular nomenclature.

3. If hydrogen is written first (and it is in aqueous solution) then name it as an Acid

There is a second sample Module 4 exam posted for you to practice you names and formulas:

Module Four Sample Exam:

Module Four Part I:  and with  Answers