Chapter Five From Atoms to Compounds

Chapter Five

From Atoms to Compounds

Covalent Vs. Ionic Compounds

Recall that atoms want to have a full outer/valance energy level with an octet of eight electrons.

There are two ways that an atom can do this; they can either share electrons, or transfer electrons. We will consider the sharing of electrons first.

Covalent Compounds

When two or more non-metals form a compound they share electrons to form what are known as molecular or covalent compounds. At room temperature, these compounds can be solids, (like sugar, C6H12O6), liquids (H2O), or gases (carbon dioxide, CO2). Covalent compounds exist as distinct individual molecules. This means that if we were to look at a sample of a covalent compound under an “ultra-powerful” microscope, we would see separate/distinct molecules floating around everywhere as the images below illustrate.

O=C=OO=C=O O=C=O O=C=O O=C=O O=C=O

I put “ultra-powerful,” in quotes because atoms are far too small to be seen with an ordinary microscope like the ones found in a biology class. Atoms are extremely tiny. To see atoms, a microscope about 1,000,000 times more powerful than a regular light microscope is needed.

We will learn more about how to draw structures of these types of compounds later in the course.

For now, we will simply learn how to name them.

Naming of covalent compounds makes use of prefixes (listed below) to tell the reader how many of each atom are present in a given compound.

Number Prefix Number Prefix

1 mono 6 hexa

2 di (not bi) 7 hepta

3 tri 8 octa

4 tetra (not quad) 9 nona

5 penta 10 deca

To name a molecular compound follow the following formula

1^st prefix (unless mono) + name 1^st element + 2^nd prefix + first part of 2^nd element + “ide”

Using this formula would give the following names N2S4 = dinitrogen tetrasulfide

S2O3 = disulfur trioxide

To shorten names, chemists do NOT include the prefix mono if there is only one of the first element. In addition, when adding a prefix would give double vowels like “oo” or “ao” the extra vowel is dropped. The following examples illustrate these common errors.

CO = carbon monoxide, not monocarbon monoxide, or carbon monooxide CO2 = carbon dioxide, not monocarbon dioxide

N2O4 = dinitrogen tetroxide, not dinitrogen tetraoxide

Do the following practice problems before moving onto naming ionic compounds.

Problem 1 – Fill in the following table of names and formulas of covalent compounds.

Name Formula Nitrogen dioxide

P2O10

Cl2O7

Phosphorus trichloride

ICl3

Phosphorus pentabromide Sulfur hexafluoride

SO

This marks the end of the section on naming of covalent compound. The rest of this chapter is dedicated to ionic compounds. In naming ionic compounds, prefixes like di, tri, tetra, etc. are NOT used. If you use a prefix to do any of the naming problems after this point in this chapter you are making one of the most common errors students make. Prefixes are for naming covalent compounds only. Do NOT use prefixes to name ionic compounds.

Ionic Compounds

When metals and a non-metals form compounds, they do so by transferring (gaining and losing) electrons. Recall that metals lose electrons to form positively charged cations, while non-metals gain electrons to form negatively charged anions. Because opposite charges attract, cations and anions will be attracted to each other and form what are known as ionic compounds. Ionic compounds are compounds made up of oppositely charged ions that are held together by ionic bonds; bonds that result from the attraction between oppositely charged ions. Ionic compounds are nearly always crystalline solids, and are often referred to as “salts.” In chemistry the word

“salt,” can refer to any ionic compound. “Table salt,” refers specifically to sodium chloride, one of the most well-known ionic compounds. In the case of table salt, the sodium atom loses an electron to form the sodium ion (Na⁺), while the chlorine atom gains that same electron to form the chloride ion (Cl^-); the attraction between these two oppositely charged ions forms the ionic bond. This process can be shown by writing a set of reactions where the symbol e^- is used to represent the transferred electron:

Na  Na⁺ + e^- Cl + e^-  Cl^-

In looking at the reaction above you will notice that the number of electrons lost by sodium (one) is equal to the number of electrons gained by chlorine. This makes sense given that salt is a neutral compound. Chemists say that the charges must balance. With sodium chloride, this is easy because everything has a charge of one, but there are many other ionic compounds that can form where the charges may be different. For example, based on its location on the periodic table, magnesium forms a plus two ion (Mg²⁺). Therefore, if magnesium were to react with chlorine, we would have to write the formula as MgCl2 because two minus one chloride ions would be needed to cancel out the plus two charge from the magnesium ion.

Unlike covalent compounds that exist as distinct molecules, ionic compounds exist in a repeating 3D array of billions and billions of cations and anions arranged into what is known as a crystal lattice. Because there are no distinct “molecules” for ionic compounds, scientists do not use the terms “molecule” or “molecular” when discussing ionic compounds, but rather use the term formula unit to refer to the formulas of ionic compound. Below is an image (obtained from

Wikipedia) of the crystal lattice of sodium chloride, NaCl. If it were possible to look at a sample of NaCl under a ultra-powerful microscope this is what you would see; a seemingly endless array of zillions of cations and anions stacked neatly, and repeating over and over again. The purple spheres represent Na⁺ cation and the green spheres represent Cl^- anions.

Notice that there is no one distinct NaCl “molecule.” Instead, the formula NaCl simply tells the reader that the ratio of Na⁺ cations to Cl^- anions is 1:1, as a result there are equal numbers of green and purple spheres in the crystal lattice diagram above.

For compounds where the ions are not 1:1 the crystal lattice looks a little different. Below is the crystal lattice of MgCl2. In this compound the ratio of Mg²⁺ cations to Cl^- anions is 1:2 and thus we see one green (Mg²⁺) for every two purple spheres (Cl^-). For simplicity, I have drawn the crystal lattice as 2D, however the actual lattice is three dimensional.

Unlike with covalent compounds where prefixes are used to indicate the number of atoms of each element in a compound, in ionic compounds the charges of the ions determine the ratio of the elements in a compound. Because chemists know what charge ions the different elements form, they are able to use those charges to figure out the ratio of elements in an ionic compound and as a result, they do NOT use prefixes to name ionic compounds.

Summary of the Ion Charges of the Elements

In the previous chapter, we learned how to predict the charge ion that many elements form based upon their location on the periodic table by applying the octet rule. For your convenience I have included that table again below.

IA 1 valence

e^-

IIA 2 valence

e^-

IIIA 3 valence

e^-

IVA 4 valence

e^-

VA 5 valence

e^-

VIA 6 valence

e^-

VIIA 7 valence

e^-

VIIIA 8 valence

e^-

H⁺ He

no ions

Li⁺ Be²⁺

Octet Rules does not apply to transition metals _N3-

nitride

O^2-

oxide

F^-

fluoride

Ne

no ions

Na⁺ Mg²⁺ Al³⁺ P^3-

phosphide S^2-

sulfide Cl^-

chloride Ar

no ions

K⁺ Ca²⁺ As^3-

arsenide Se^2-

selenide Br^-

bromide Kr

no ions

Rb⁺ Sr²⁺ Te^2-

telluride

I^-

iodide

Xe

no ions

Cs⁺ Ba²⁺ Rn

no ions

All of the metal elements in the table above only form one charge of ion. In addition to these elements, silver, zinc, and cadmium only form the following ions: Ag⁺, Zn²⁺, Cd²⁺. Every other metal on the periodic table can form multiple charges. Chemists memorize those charges for important elements, and look up the others when needed. In this class, you will only be required

Fe +2/3

Ag +1 Cu +1/2

Cd +2 Zn +2

Al +3

Sn +2/4 Pb +2/4

O -2 S -2 Se -2 Te -2 N -3

P -3 As -3

F -1 Cl -1 Br -1 I -1 14

IVA 15 VA

16 VIA 12

IIB 13 IIIA

18 VIIIA 17 VIIA 7

VIIB 8 ----

9 VIIIB 5

VB 6 VIB

11 IB 10 ---- 3

IIIB 4 IVB

H +1

Li +1 Na +1

Rb +1 K +1

Cs +1

Be +2 Mg

+2 Ca +2 Sr +2 Ba +2 1 IA

2 IIA

*

**

to memorize the charges that the elements copper (Cu), iron (Fe), tin (Sn), and lead (Pb) form as summarized in the following table:

Element Ions Formed Ion Names Copper (Cu) Cu⁺ Copper (I)

Cu⁺² Copper (II)

Iron (Fe) Fe²⁺ Iron (II) Fe³⁺ Iron (III)

Tin (Sn) Sn²⁺ Tin (II)

Sn⁴⁺ Tin (IV)

Lead (Pb) Pb²⁺ Lead (II)

Pb⁴⁺ Lead (IV)

To summarize, you are expected to know the ions listed in the periodic table below:

Using Ion Charges to Write Formulas of Ionic Compounds

Ionic compounds are neutral compounds. This fact requires that the formulas of ionic compounds are written such that the total positive charge from the cations is equal to the total negative charge from the anions. With compounds like sodium chloride (NaCl) or magnesium chloride (MgCl2) this is easy. Sodium has a charge of plus one and chloride has a charge of minus one.

MEMORIZE THESE IONS

Therefore, one Na⁺ cation is needed for every one Cl^- anions giving the formula Na1Cl1; which is always written as NaCl, as ones in formulas are understood, they are never written. With

magnesium two chloride ions are needed to cancel out the positive two charge from the Mg²⁺ ion.

The paper cutouts below, that were used in lecture, provide a visual representation of this process.

When the ion charges are not simple number like one or two, it can be useful to use the “crissy- crossy” trick to help write correct formulas of ionic compounds. We will use aluminum and oxygen as an example. Aluminum forms a +3 ion (Al³⁺) and oxygen forms a -2 ion (O^2-). Our goal is to write a formula that gives these ions in a ratio such that the charge cancels out. To do this we criss-cross the numbers (just the numbers, not the charges) as illustrated below.

Al³⁺ O^2-

Al2O3

What this formula says is that the combined positive charge to two Al³⁺ ions (+6) is canceled by the combined negative charge (-6) of three O^2- anions. Because a 2:3 ratio can not be simplified the formula is written as Al2O3.

For a slightly different example let’s consider the formula that results when magnesium and oxygen form a compound. We begin by writing our ions with their charges and then criss-cross the numbers as before

Mg²⁺ O^2-

Mg2O2

giving the formula Mg2O2. Because a 2:2 ratio can be simplified this formulas is written as MgO.

Mg2O2  MgO

This last simplification step is important. Writing the formula as Mg2O2 is wrong. Chemists always simplify formulas of ionic compounds when possible.

Problem 2- write formulas for the ionic compounds forms by the ions in the table below.

F^- Cl^- O^2- S^2- N^3- P^3-

Li⁺ Na⁺ K⁺ Mg²⁺

Ca²⁺

Zn²⁺

Cd²⁺

Fe²⁺

Fe³⁺

Sn²⁺

Sn⁴⁺

Cu⁺ Cu²⁺

Ag⁺

Naming Ionic Compounds

How ionic compounds are named depends upon the metal ion involved. If the metal can only form one charge of ion, then the charge is NOT given in the name. In an ionic compound, the charges of the ions determine their ratios in the formula of a compound. As a result prefixes are NOT used when naming ionic compounds because they are not necessary. To name ionic compounds we apply the following formula:

Name of cation + 1st part of anion name + “ide”

Look at the following examples and see how they fit the above formula NaCl – Sodium chloride

KCl – Potassium chloride

CaCl2 – Calcium chloride (NOT calcium dichloride - no prefixes in ionic naming) Na2S – sodium sulfide (NOT disodium monosulfide – no prefixes in ionic naming) Al2O3 – aluminum oxide (NOT dialuminum trioxide – no prefixes in ionic naming) ZnBr2 – Zinc bromide (NOT zinc dibromide – no prefixes in ionic naming) Ag2O – silver oxide (NOT disilver oxide – no prefixes in ionic naming) CdS – Cadmium sulfide

If the metal in question can form more than one charge, it is necessary to specify the charge in the name. To keep things as simple as possible, we will only consider ions of Cu, Fe, Sn, and Pb as listed in previously. To name such compounds we apply the following formula:

Name of cation + (charge in Roman Numerals) + 1st part of anion name + ide Below are some examples for iron:

FeCl2 – Iron (II) chloride (NOT iron dichloride – no prefixes in ionic naming) FeCl3 – Iron (III) chloride (NOT iron trichloride – no prefixes in ionic naming)

Common Naming Errors for Ionic Compounds

If both ions have a charge other than one, the formulas can become more complicated. When this happens students making some common naming errors. For example, when iron reacts with oxygen (O^2-) two different compounds can be formed depending on if its Fe²⁺ or Fe³⁺ that combines with the oxygen:

FeO – Iron (II) oxide Fe2O3 – Iron (III) oxide (NOT Iron (I) oxide) (NOT Iron (II) oxide)

Notice that there is only one iron atom in the formula of iron (II) oxide (FeO) and that there are two iron atoms in the formula of iron (III) oxide (Fe2O3). It is the charge on iron, not the number of iron atoms, that belongs inside the ( )’s. Another example of compounds that students often name incorrectly are the oxides of copper:

Cu2O – Copper (I) oxide CuO – Copper (II) oxide (NOT copper (II) oxide) (NOT copper (I) oxide)

Problem 3- Fill in the following table of names and formulas for ionic compounds.

Name Formula Potassium bromide

Na3N Ca3P2

Magnesium sulfide

PbO Pb3As4

Fe2S3

Barium selenide Aluminum iodide

AgF Copper (I) oxide

Copper (II) oxide

Li2S PbS PbO2

ZnCl2

Tin (II) phosphide Tin (IV) phosphide

CdO

Ionic Compounds Containing Polyatomic Ions

Sometimes several atoms bond together to make one big ion known as a polyatomic ion. These multi-atom ions are so common that they are given their own names. Chemists memorize the names and formulas of dozens of these ions; luckily you will only have to memorize the nine common ones given below.

Formula Ion Name NH4+ Ammonium OH^- Hydroxide

CN^- Cyanide

NO3- Nitrate HCO3- Bicarbonate C2H3O2- Acetate CO32- Carbonate SO42- Sulfate PO43- Phosphate

Writing formulas for compounds containing these ions is no more difficult than writing formulas of any other ionic compound. We still criss-cross the numbers. If more than one of a polyatomic ion is needed we put the polyatomic ion in parentheses. As usual, the formula is simplified whenever possible.

For an example consider the formula that results when aluminum reacts with nitrate (NO3-).

Because nitrate has a minus one charge, three nitrates will be needed.

Al³⁺ NO3-

Al(NO3)3

In the formula above the nitrate ion is enclosed in ( ) to indicate that the subscript three applies to the entire ion.

When possible formulas are simplified. For example consider the compound of Al³⁺ and PO43-

Al³⁺ PO43-

Al3(PO4)3

which must be simplified to give AlPO4 as the final formula.

Al3(PO4)3  AlPO4

Notice that no ( ) were used in the formula AlPO4. Parentheses are only used if there are

multiple polyatomic ions. To use parentheses when there is only one of a polyatomic ion, such as Al(PO4) is wrong.

Problem 4- Write formulas for ionic compounds containing the following polyatomic ions.

OH^- NO3- SO42- CO32- PO43-

Li⁺ Na⁺ NH4+

Mg²⁺

Zn²⁺

Cd²⁺

Fe²⁺

Fe³⁺

Pb²⁺

Pb⁴⁺

Cu⁺ Cu²⁺

Ag⁺

Naming compounds containing polyatomic ions is no different that naming any other ionic compound. For elements that can only form one charge, the name is simply the name of the cation followed by the name of the polyatomic ions.

NaHCO3 – sodium bicarbonate (a.k.a. baking soda) KOH – potassium hydroxide

Ammonium (NH4+) is the only polyatomic ion with a positive charge. This sometimes confuses students. Below are a couple examples containing the ammonium ion. In these compounds, the ammonium cation is playing the role that the metal ion usually plays.

NH4Cl – ammonium chloride NH4NO3 – ammonium nitrate

If multiple polyatomic ions are needed to balance the charges, then the polyatomic ion is enclosed in ( ).

(NH4)2SO4 – ammonium sulfate Ca(OH)2 – calcium hydroxide Zn(C2H3O2)2 – zinc acetate

If the metal in question can form multiple charges, the charge must be included in the name.

Fe(NO3)2 – Iron (II) nitrate Fe(NO3)3 – Iron (III) nitrate SnSO4 – Tin (II) sulfate Sn(SO4)2 – Tin (IV) sulfate

Problem 5- Fill in the following table regarding ionic compounds containing polyatomic ions.

Name Formula

Potassium bicarbonate

(NH4)3PO4

Ca(NO3)2

Magnesium sulfate

Pb(C2H3O2)2

Pb(OH)4

Barium carbonate Aluminum cyanide

AgCN Copper (I) carbonate

Copper (II) sulfate

NH4NO3

Pb(C2H3O3)4

Pb(CN)4

ZnCO3

Tin (II) phosphate Tin (IV) bicarbonate

CdSO4

Answers to In Chapter Questions

1.

Name Formula

Nitrogen dioxide NO2

Diphosphorus decoxide P2O10

Dichlorine heptoxide Cl2O7

Phosphorus trichloride PCl3

Iodine trichloride ICl3

Phosphorus pentabromide PBr5

Sulfur hexafluoride SF6

Sulfur monoxide SO

2.

F^- Cl^- O^2- S^2- N^3- P^3-

Li⁺ LiF LiCl Li2O Li2S Li3N Li3P Na⁺ NaF NaCl Na2O Na2S Na3N Na3P

K⁺ KF KCl K2O K2S K3N K3P

Mg²⁺ MgF2 MgCl2 MgO MgS Mg3N2 Mg3P2

Ca²⁺ CaF2 CaCl2 CaO CaS Ca3N2 Ca3P2

Zn²⁺ ZnF2 ZnCl2 ZnO ZnS Zn3N2 Zn3P2

Cd²⁺ CdF2 CdCl2 CdO CdS Cd3N2 Cd3P2

Fe²⁺ FeF2 FeCl2 FeO FeS Fe3N2 Fe3P2

Fe³⁺ FeF3 FeCl3 Fe2O3 Fe2S3 FeN FeP Sn²⁺ SnF2 SnCl2 SnO SnS Sn3N2 Sn3P2

Sn⁴⁺ SnF4 SnCl4 SnO2 SnS2 Sn3N4 Sn3P4

Cu⁺ CuF CuCl Cu2O Cu2S Cu3N Cu3P Cu²⁺ CuF2 CuCl2 CuO CuS Cu3N2 Cu3P2

Ag⁺ AgF AgCl Ag2O Ag2S Ag3N Ag3P

3.

Name Formula

Potassium bromide KBr

Sodium nitride Na3N

Calcium phosphide Ca3P2

Magnesium sulfide MgS

Lead (II) oxide PbO

Lead (IV) arsenide Pb3As4

Iron (III) sulfide Fe2S3

Barium selenide BaSe

Aluminum iodide AlI3

Silver fluoride AgF

Copper (I) oxide Cu2O

Copper (II) oxide CuO

Lithium sulfide Li2S

Lead (II) sulfide PbS

Lead (IV) oxide PbO2

Zinc chloride ZnCl2

Tin (II) phosphide Sn3P2

Tin (IV) phosphide Sn3P4

Cadmium oxide CdO

4.

OH^- NO3- SO42- CO32- PO43-

Li⁺ LiOH LiNO3 Li2SO4 Li2CO3 Li3PO4

Na⁺ NaOH NaNO3 Na2SO4 Na2CO3 Na3PO4

NH4+ NH4OH NH4NO3 (NH4)2SO4 (NH4)2CO3 (NH4)3PO4

Mg²⁺ Mg(OH)2 Mg(NO3)2 MgSO4 MgCO3 Mg3(PO4)2

Zn²⁺ Zn(OH)2 Zn(NO3)2 ZnSO4 ZnCO3 Zn3(PO4)2

Cd²⁺ Cd(OH)2 Cd(NO3)2 CdSO4 CdCO3 Cd3(PO4)2

Fe²⁺ Fe(OH)2 Fe(NO3)2 FeSO4 FeCO3 Fe3(PO4)2

Fe³⁺ Fe(OH)3 Fe(NO3)3 Fe2(SO4)3 Fe2(CO3)3 FePO4

Pb²⁺ Pb(OH)2 Pb(NO3)2 PbSO4 PbCO3 Pb3(PO4)2

Pb⁴⁺ Pb(OH)4 Pb(NO3)4 Pb(SO4)2 Pb(CO3)2 Pb3(PO4)4

Cu⁺ CuOH CuNO3 Cu2SO4 Cu2CO3 Cu3PO4

Cu²⁺ Cu(OH)2 Cu(NO3)2 CuSO4 CuCO3 Cu3(PO4)2

Ag⁺ AgOH AgNO3 Ag2SO4 Ag2CO3 Ag3PO4

5.

Name Formula

Potassium bicarbonate KHCO3

Ammonium phosphate (NH4)3PO4

Calcium nitrate Ca(NO3)2

Magnesium sulfate MgSO4

Lead (II) acetate Pb(C2H3O2)2

Lead (IV) hydroxide Pb(OH)4

Barium carbonate BaCO3

Aluminum cyanide Al(CN)3

Silver cyanide AgCN

Copper (I) carbonate Cu2CO3

Copper (II) sulfate CuCO3

Ammonium nitrate NH4NO3

Lead (IV) acetate Pb(C2H3O3)4

Lead (IV) cyanide Pb(CN)4

Zinc carbonate ZnCO3

Tin (II) phosphate Sn3(PO4)2

Tin (IV) bicarbonate Sn(HCO3)4

Cadmium sulfate CdSO4