Ionic Compound

Ionic compounds are basically defined as being compounds where two or more ions are held next to each other by electrical attraction

Ionic Compounds

This is a summary of the past and present nomenclature or naming conventions for ionic compounds

Naming Compounds

An ionic compound is one in which at least two of the elements or compounds in the group are oppositely-charged ions held together

Ionic Compounds

Ionic compounds generally are very hard and have very high melting points. They are solids at room temperature

Face Centered Cubic

When metals react with non-metals they form an ionic compound. Ions have a charge because electrons are lost or gained in forming an ionic bond.

Sunday, July 31, 2011

Balancing Charges on Ions


Formula of an Ionic Compound
Balancing Charges on Ions


Introduction
Atoms of different elements combine with one another to form compounds. The empirical formula of an ionic compound indicates the kinds of atoms that are present in the compound as well as the relative number (ratio) of each kind of atom. Let’s investigate how the formula of an ionic compound can be determined.

Concepts

•Ionic compounds             •Empirical formula
•Polyatomic ions               •Precipitation reaction

Background

An ionic compound is composed of ions – atoms or groups of atoms that have a positive or negative charge. Oppositely charged ions arrange themselves into an extended, three-dimensional structure called a crystal lattice. The net attractive forces among oppositely charged ions in the crystal structure are called ionic bonds. Although composed of charged ions, ionic compounds are electrically neutral. The ratio of oppositely charged ions in the crystal structure is such that the positive charge contributed by the cations in equal to or balanced by the negative charge contributed by the anions. There is no net or overall charge on an ionic compound.

The empirical formula of an ionic compound indicated the smallest whole number ratio of each type of ion in the crystal structure and is called a formula unit. For example, magnesium chloride has the empirical formula MgCl2. Magnesium cations (Mg2+) and chloride anions (Cl-) combine in a 1:2 ratio to form the MgCl2 formula unit. The overall charge on ionic compounds is always a zero.

Some ions consist of a charged group of covalently bonded atoms. Such ions are called polyatomic ions. An example is the nitrate ion (NO3-), which contains one nitrogen atom and three oxygen atoms and has an overall charge of –1. In calcium nitrate, calcium (Ca2+) ions combine with nitrate ions in a 1:2 ratio in order to balance the positive and negative charges. The empirical formula for calcium nitrate is Ca(NO3)2. Parentheses are used around the nitrate ion to show that the subscript “2” pertains to the nitrate ion as a whole.

Many ionic compounds can be prepared in the lab using precipitation reactions. When solutions of two ionic compounds are combined, the ions may rearranged to form a new ionic compound that is insoluble in water. An example of this type of reaction is the formation of solid barium sulfate when barium chloride and sodium sulfate are combined in solution (Equation 1a). In Equation 1b, only the ions that form the precipitate are represented. This makes it easier to recognize what happens in the precipitation reaction.

                  BaCl2(aq)  +  Na2SO4(aq)  à  BaSO4(s)  +  2NaCl(aq)               Equation 1a
                                Ba2+(aq)  +  SO42-(aq)  à  BaSO4(s)                              Equation 1b

According to the balanced equation for this reaction, barium ions (Ba2+) combine with sulfate ions (SO42-) in a 1:1 ratio to form barium sulfate (BaSO4). This ratio can be observed experimentally in the lab by mixing BaCl2(aq) and Na2SO4(aq) solutions containing equal amounts (concentrations) of barium and sulfate ions, respectively. The maximum amount of precipitate will be obtained when equal volumes (a 1:1 ratio) of the two solutions are combined. A similar approach can also be used to determine the formula of an unknown ionic compound.

Experiment Overview

The purpose of this experiment is to determine the empirical formula of an unknown ionic compound. Two solutions containing equal amounts (concentrations) of two reactant ions will be combined in a series of reactions. In each reaction, the totally volume of the two solutions will be held constant while the volume ratio of the reactants is varied. The amount of precipitate obtained in each reaction will be measured and plotted against the volume ratio to find the empirical formula of the product.

Pre-Lab Questions

1.      Many common drugstore chemicals are ionic compounds. Write the correct empirical formula for each of the following compounds.
Common name:     Milk of magnesia                    Washing soda           Epsom salt
Chemical name:    Magnesium hydroxide    Sodium carbonate    Magnesium sulfate

2.      Solutions of iron(III) chloride and sodium hydroxide were mixed in a series of precipitation reactions, as described in this experiment.
(a)    Name the two possible products in this precipitation reaction and predict
      their empirical formulas.
(b) Which product is likely to be insoluble in water and precipitate out as a
      red solid?
(c) What volume ration of reactants gave the most precipitate (see Table 1)?
      Explain.

Table 1.
Test tube
1
2
3
4
5
6
7
FeCl3, 0.1 M, mL
5
10
12
15
17
20
24
NaOH, 0.1 M, mL
55
50
48
45
43
40
36
Volume of precipitate, mL
1
10
14
20
4
1
0







Materials
Copper(II) chloride solution, CuCl2,                 Metric ruler, marked in mm
0.1M, 6ml                                                        Pipets, Beral-type, 2
Sodium phosphate solution, Na3PO4,                Stirring rod or wood splints
0.1M, 6ml                                                        Test tubes, small, 7
Marking pen or wax pencil                                Test tube rack or 24-well
     reaction plate

Safety Precautions
Copper(II) chloride and sodium phosphate solutions are skin and eye irritants are slightly toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure
1.      Label seven small test tubes #1-7 with a marking pen and place them in a test tube rack or in a 24 well-reaction plate.
2.      Cut the stems of two Beral-type pipets at a 45º angle about 5cm from the bulb.
3.      Fill one pipet with 0.1M copper(II) chloride solution and record the color of the solution in the data table.
4.      Carefully add the appropriate number of drops of copper(II) chloride solution to each test tube #1-7, as shown in Table 2. Note: Exact volumes are very important – hold the pipet vertically to obtain uniform size drops.
5.      Fill the second pipet with 0.1M sodium phosphate solution and record the color of the solution in the data table.
6.      Carefully add the appropriate number of drops of sodium phosphate solution to each test tube, as shown in Table 2.

Table 2.
Test Tube
1
2
3
4
5
6
7
CuCl2, 0.1M (mL)
5
4
3
2
1
0

Na3PO4, 0.1M (mL)
0
1
2
3
4
5


7.      Use a clean stirring rod or wood splint to stir each reaction mixture in test tubes #1-7. Let the tubes sit undisturbed for 10-15 minutes to allow the precipitates to settle.
8.      During this time, determine the volume (drop) ratio of copper(II) chloride and sodium phosphate solutions in each test tube. Write this ration in the data table. Example: In test tube #1, 3 drops of CuCl2 and 27 drops of Na3PO4 correspond to a 1:9 ratio of CuCl2:Na3PO4.
9.      After the precipitates have settled, observe the appearance of the products (both the solid and the solution). Record the observation in the data table in the space provided. Be as detailed as possible.
10.  Use a metric ruler to measure the height of the precipitate of millimeters in each test tube. Read from the top of the solid material to the bottom center of the test tube. Record each height in mm in the data table.
11.  Dispose of the contents of the test tubes as directed by your instructor.



Formula of an Ionic Compound

Data Table
Color of CuCl2
Color of Na3PO4 Solution
Appearance of Products

Precipitation Reactions
Test tube
1
2
3
4
5
6
7
Volume Ratio *       (Drops of CuCl2:Drops Na3PO4







X
 
 
Height of Precipitate (mm)






X
 
 

*Reduce the volume ratio to the simplest whole-number ratio


Post-Lab Questions
1.      (a) Name the two possible products in the precipitation reaction of copper(II) chloride with sodium phosphate. Use the charges on the ions to predict the empirical formulas of the products.
(b) Based on common knowledge, which product is likely to be insoluble in water and to precipitate from solution?

2.      Complete the following bar graph to show the height of the precipitate in each test tube.


3.      Which test tube had the greatest amount of precipitate? Does this result agree with the prediction made in Questions #1 concerning the empirical formula of the product? Explain.

4.      Write a balanced chemical equation for the precipitation reaction of copper(II) chloride and sodium phosphate. Include abbreviations for the physical state of each reactant and product, using (aq) for aqueous solution, (s) for solid, (l) for liquid), and (g) for gas.

5.      (a) Which test tubes showed evidence of unreacted Cu2+ ions in the supernatant when the reaction was complete? Explain why unreacted Cu2+ ions were present in these tubes based on the volume ratio of the solutions used.

(b) How could you tell that all of the Cu2+ ions had reacted in a particular test tube? Which test tubes showed such evidence? Explain, based on the volume ratio of solutions used.

6.      What was the totally number of drops of solution in each test tube? Why was it necessary to keep the totally volume of reactant constant in each test tube?

7.      (Optional) Does the height of precipitate in each test tube accurately reflect the amount of precipitate in each case? Hint: Compare the shape of a test tube to that of a graduated cylinder. What effect does this error have on the conclusions reached in this experiment?

Saturday, July 30, 2011

Naming Ioning Compounds



Nomenclature of Compounds

Introduction

Naming Ionic Compounds
An ion is a single atom or a group of atoms with a net charge.  A positively charged ion is called a cation.  A negatively charged ion is called an anion.  An ion occurs when the total numbers of electrons and protons are not equal.  For example, a neutral magnesium atom has 12 protons and 12 electrons.  The magnesium ion, however, has 12 protons, but only 10 electrons.  The net charge is 2+ and the formula for the ion is written as Mg2+.  This is an example of a monatomic ion (single atom ion).  Groups of covalently bonded atoms can also have a net charge.  For example, the hydroxide ion consists of one atom each of oxygen and hydrogen.  The group as a whole has 9 protons and 10 electrons giving a net charge of 1-.  The formula of this polyatomic ion ("many" atom ion) is written as OH-.  Most polyatomic ions are anions; the ammonium ion, NH4+, is an important polyatomic cation.

Naming monatomic ions follows some simple rules.  The ionic forms of some transition metals and the names and formulas of the polyatomic ions, however, must be memorized. 

There are some helpful hints and trends for the monatomic ions to make learning easier.  Metals tend to form cations and nonmetals tend to form anions.  The common charges of monatomic main group ions are related to their position on the periodic table.  All elements in Group IA (the first column - the alkali metals) form cations with a 1+ charge.  All elements in Group IIA (the second column - the alkaline earth metals) form cations with a 2+ charge.  Metals in Group IIIA form cations with a 3+ charge. Elements in Group VIIA (the halogens) form anions with a 1- charge.  Elements in Group VIA (the chalcogens) form anions with a 2- charge.

Most main group elements have only one monatomic ionic form.  In such cases, the monatomic cations are simply named with the element name followed by the word "ion".  For example Na+ is the sodium ion.  The monatomic anions are named by changing the suffix of the element name to "ide" and adding the word "ion".  For example F- is the fluoride ion and O2- is the oxide ion.

Most transition metals (and even some heavier main group elements), however, have more than one possible monatomic ionic form.  For example, iron has two ionic forms, Fe2+ and Fe3+.   These two forms are distinguished in the written names of the ions by writing the charge as a Roman numeral in parentheses after the name of the element (without a space).  Thus Fe2+ is the iron(II) ion (pronounced "iron two ion").  This ion was formerly called the ferrous ion from the Latin name for iron, ferrum. Many other transition metals also have old names that are still used.  For more examples see Table 3.1.

Ionic compounds consist of oppositely charged ions present in a ratio to yield a uncharged substance.  Usually only two kinds of ions are present.  Occasionally three kinds are present, but this is uncommon. 

The nomenclature of the ionic compounds is simple.  By convention, the name of the cation is written first and the name of the anion is written second.  The word "ion" is omitted and the names are separated by a space.  The same order (cation first, anion second) applies to writing the chemical formulas of ionic compounds.  The chemical formula, however, also indicates the smallest whole number ratio between the ions as subscripts.  For example, table salt is composed of a 1:1 ratio of Na+ and Cl- ions.  Its name is sodium chloride and the formula is NaCl.  (Note that omitted charge numbers and subscripts imply a "1".)

Table 3.1

Name of Ion
Formula of Ion
Classic Name of Ion
copper(I) ion
Cu+
cuprous ion
copper(II) ion
Cu2+
cupric ion
iron(II) ion
Fe2+
ferrous ion
iron(III) ion
Fe3+
ferric ion
lead(II) ion
Pb2+
plumbous ion
lead(IV) ion
Pb4+
plumbic ion
gold(I) ion
Au+
aurous ion
gold(III) ion
Au3+
auric ion


Consider the following examples:

Example 1:  Write the formula for magnesium chloride.
Solution:  The cation is Mg2+, and the anion is Cl-.  When the anion and the cation are combined, the charges must "cancel".  In other words, the total number of positive charges has to be the same as the total number of negative charges.  The smallest whole number ratio of cations to anions to cancel the charges can be found mathematically using the least common multiple method.  However, an easier way of accomplishing this is to write the cation followed by the anion, and criss-cross the numbers where the number (without its sign) of the charge on the cation becomes the subscript for the anion, and the number of the charge on the anion becomes the subscript for the cation as shown below:

Mg2+     Cl1-      à     Mg1Cl2     à     MgCl2

Note that in writing the final formula, if there is no subscript, then the number is understood to be 1. 

Example 2:  Write the formula for calcium phosphate.
Solution:  The cation is Ca2+ and the anion is PO43-.  Use the criss-cross method to determine the subscripts.

Ca2+        PO43-     à    Ca3(PO4)2

Note that the polyatomic ion is in parentheses.  Whenever you have more than one of a polyatomic ion in a formula, enclose the polyatomic ion formula in parentheses and put the subscript outside the parentheses.

Example 3:  Write the formula for tin(IV) sulfide.
Solution:  The cation is Sn4+ (remember the Roman numeral tells what the charge is on the cation) and the anion is S2-.  Use the criss-cross method to determine the subscripts.

Sn4+             S2-    à    Sn2S4    à   SnS2

Notice that both of the subscripts are divisible by two.  In writing formulas for ionic compounds, you need to write the empirical formula with respect to the cation and the anion.  In other words, the subscripts must be in the lowest whole number ratio.  In order to accomplish this for tin(IV) sulfide divide the subscripts by 2 to get SnS2.

Example 4:  Write the formula for iron(III) phosphate.
Solution:  The cation is Fe3+ and the anion is PO43-.  You can do the criss-cross method if you wish, but that is not necessary here.  Since the numbers on the charges are the same for the positive and the negative charge, you only need one of each (in other words, no subscripts).  Therefore, the formula is FePO4.

Example 5:  Name the compound K2S.
Solution:  The key to naming the compound is to identify the cation and anion in the compound.  If you cannot find a distinct anion and a distinct cation in the compound, then the compound is not ionic.  Another way to determine if the compound is NOT ionic is that all of the atoms involved are non-metals or metalloids (unless it begins with NH4).

When attempting to identify the cation and the anion, remember that the cation is always written first, and except for ammonium (NH4+), the cation only consists of one atom.  Here, the cation is K+, and the anion is S2-.  In this case, the charges are deduced by the positions of potassium and sulfur on the periodic chart.

Once you have identified the cation and the anion, simply name the cation first, then the anion (do not write the word "ion" in the name of the ionic compound).  The name of K2S is potassium sulfide.

Example 6:  Name the compound Ba(NO3)2 .
Solution:  Let's identify the cation and the anion.  There are more than two different atoms here.  Since "NH4" is not in the formula, it is the anion that is polyatomic.  The cation is Ba2+, and the anion is NO3-.  The charge on barium must be 2+ since it is an alkaline earth metal.  The charge on nitrate must be memorized.  The name of this compound is barium nitrate.

Example 7:  Name the compound CuBr2
Solution:  Identifying the anion is easy.  The anion is Br-.  Since the cation is a transition metal (or heavy main group metal), we need to look at the charge from the anion to help us determine the charge on the cation.  Each Br- anion has a 1- charge.  Since there are two bromide ions, the total negative charge is 2-.  This means that the cation(s) must have a total charge of 2+.  Since there is only one cation, the charge for each cation is 2+.  Therefore, the cation must be Cu2+, and the compound is named copper(II) bromide, or cupric bromide. 

A list of ions is provided for you to help you with this exercise.

Cations with only one possible charge:

Li+       lithium ion                                          Al3+     aluminum ion
Na+      sodium ion                                          NH4+   ammonium ion
K+        potassium ion                                      Zn2+     zinc(II) ion
Mg2+    magnesium ion                                    Ag+      silver(I) ion
Ca2+     calcium ion                                          Ni2+     nickel(II) ion
Sr2+      strontium ion                                       Ba2+     barium ion

Cations which can have more than one charge:

Fe2+     iron(II) ion or ferrous ion                    Sn2+     tin(II) ion or stannous ion      
Fe3+     iron(III) ion or ferric ion                     Sn4+     tin(IV) ion or stannic ion
Cu+      copper(I) ion or cuprous ion                Pb2+     lead(II) ion or plumbous ion
Cu2+     copper(II) ion or cupric ion                 Pb4+     lead(IV) ion or plumbic ion
Co2+     cobalt(II) ion or cobaltous ion                        Mn2+    manganese(II) ion or manganous ion
Co3+     cobalt(III) ion or cobaltic ion              Mn3+    manganese(III) ion or manganic ion

Anions:

H-                    hydride ion                              C2H3O2-          acetate ion
F-                    fluoride ion                             MnO4-             permanganate ion
Cl-                   chloride ion                             SO42-               sulfate ion
Br-                   bromide ion                             HSO4-             hydrogen sulfate ion
I-                     iodide ion                                SO32-               sulfite ion
O2-                  oxide ion                                 Cr2O72-            dichromate ion
O22-                 peroxide ion                            CrO42-             chromate ion
S2-                   sulfide ion                               PO43-               phosphate ion
N3-                  nitride ion                                C2O42-             oxalate ion
NO3-               nitrate ion                                CN-                 cyanide ion
NO2-               nitrite ion                                 OH-                 hydroxide ion
CO32-              carbonate ion                           ClO3-               chlorate ion
HCO3-             hydrogen carbonate ion          ClO-                hypochlorite ion

So far, only the naming of ionic compounds has been illustrated.  Now let's discuss how to name compounds where there are no distinct cations or anions, the covalent compounds.

Naming Covalent Compounds
Covalent compounds are composed of atoms joined by the sharing of pairs of electrons, called covalent bonds.  Binary compounds, substances composed of two nonmetallic elements, are the simplest covalent compounds. Usually, one of the elements is more electronegative than the other.  Fluorine is the most electronegative element.  Generally, the closer an element is to fluorine on the periodic table, the greater its electronegativity.  Thus, electronegativity tends to increase as you go up any given column or to the right on any given row in the periodic table (ignoring noble gases). 

In writing formulas and names, the less electronegative element is written first, followed by the more electronegative element.  In the name, the more electronegative element is written as if it were an anion, so it has the "ide" suffix. The names of the elements are separated by a space.

By the law of multiple proportions two elements may combine in more than one ratio to form different compounds.  For example, carbon and oxygen can form either CO or CO2. These compounds must have different names.  We use Latin prefixes to indicate the number of atoms of each element type as shown in the table below (Table 3.2):

Table 3.2

Number of Atoms
Prefix
1
mono-
2
di-
3
tri-
4
tetra-
5
penta-
6
hexa-
7
hepta-
8
octa-
9
nona-
10
deca-

If there is only one atom of the first element in the name / formula, the prefix "mono-" is omitted.  So CO is carbon monoxide and CO2 is carbon dioxide.  Note the "o" in "mono" is dropped before the word oxide.

Some compounds have a "common" name.  For example, trihydrogen nitride is ammonia, and dihydrogen monoxide is called water.

Example 8:  Write the formula for carbon tetrafluoride.
Solution:  Since the first atom has no prefix, there is only 1 carbon atom.  The prefix "tetra-", followed by fluoride, means there are four fluorine atoms.  Therefore the formula is written CF4.

Example 9:  Write the formula for dinitrogen trioxide.
Solution:  The first atom has a prefix "di-" meaning two, so there are two nitrogen atoms that are written first.  The prefix "tri-" means there are three oxygen atoms, so the formula is written N2O3.



Example 10:  Name the compound, IF3.
Solution:  There is only one iodine, so we do not need the "mono-" prefix.  For fluorine, the prefix "tri-" is used since there are three of them.  Also, fluorine is written like the anion, fluoride.  So the name of this compound is iodine trifluoride.

Example 11:  Name the compound, P4S6.
Solution:  There are four P atoms written first, so the first part of the name is tetraphosphorus.  The prefix "hexa-" is used since there are six of the second atom.  Since the second atom is sulfur, it is written as the anion, sulfide.  So the name of this compound is tetraphosphorus hexasulfide.

The tricky part about naming these compounds is classifying them as ionic or covalent.  Once you do that, you know which set of rules to apply.  Use the following guidelines to help you choose. 

                        Composed of only nonmetals or metalloids   Covalent
                        and no ammonium, NH4+ , is present 

                        Metal and nonmetal                                        Ionic

                        Ammonium and a nonmetal                            Ionic

                        Metal and a polyatomic ion                            Ionic

                        Ammonium and a polyatomic ion                   Ionic

                        Name contains prefixes like "mono",              Covalent
                        "di", "tri", etc.

                        Name has no prefixes and is not a                  Ionic
                        special name like water.

If you can distinguish metals from nonmetals and can recognize the polyatomic ion groups as formulas and names, the easier this will be.



Experiment 3 - Nomenclature of Compounds

REPORT

Name________________________________________      Section _________

25 bottles are set out for you with either the name or the formula on the label.  Record this on the table below.  Then fill in the rest of the missing information.  All of the compounds on this table are ionic.  See the example below.

#
Name
Formula
Cation
Anion
Color
EX.
sodium chloride
NaCl
Na+
Cl-
white
1
copper(II) bromide




2

PbCr2O7



3

Cr2(SO4)3



4
barium nitrate




5
silver(I) acetate




6

KClO3



7

CuCl



8

KMnO4



9
potassium nitrite




10

MnCl2



11

Ni(C2H3O2)2



12
ammonium sulfate




13
chromium(III) chloride




14
cobalt(II) nitrate




15

Na3PO4



16
potassium Iodide




17
cobalt(II) sulfate




18
iron(III) chloride




19

K2CO3



20
ammonium phosphate




21

ZnSO4



22
mercury(II) chloride




23
copper(I) sulfide




24

NaBr



25
iron(II) sulfate





Again, fill in the blanks in the table.  This is a list of compounds that are found in the "real world".  Some of the compounds below are ionic and others are covalent.  You must be able to tell the difference.

#
Name
Formula
Real World Use
1

AgCl
self tint sunglasses
2
dinitrogen monoxide

laughing gas
3
diphosphorus pentoxide

drying agent
4
lithium carbonate

medication for bipolar disorder
5

NaHCO3
baking soda
6
tin (II) fluoride

fluoride in toothpaste
7

Sr(NO3)2
used in road flares
8

CaCO3
Antacids
9
magnesium sulfate

Epsom salts
10

CuSO4
root eater
11
cobalt(II) chloride

humidity indicator
12

(NH4)2CO3
smelling salts
13

SO2
pollutant in coal burning plants
14
barium sulfate

nasty drink before X-rays
15

K2CrO4
used in breathalyzers
16
dihydrogen monosulfide

rotten eggs smell
17
chromium(III) oxide

dye used to color money
18
lead(II) sulfate

used in lead acid batteries
19

Fe2O3
Rust
20

MgO
magnesium supplement
21
sodium hypochlorite

Bleach
22
calcium sulfate

plaster of Paris
23

CO2
bubbles in soft drinks
24

FeSO4
generic iron supplement
25
manganese(IV) oxide

used in flashlight batteries
26
tetraphosphorus trisulfide

used in match heads
27
silicon dioxide

sand (life's a beach)



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