unit 4 notes chemistry i cp page 1 -...

Unit 4 Notes – Chemistry I CP page 1

Chemistry Annotation Guide If you are NOT using the following annotation, put in your key to the left of each item.

Your

Key

Mr. T’s

Key Items to be annotated

Mr. T’s

or

headings and subheadings.

key content vocabulary

Write DEF next to definitions (sometimes #2 and #3 will be the same and in that case I expect to see a box AND DEF)

Underline important ideas (include captions for visuals).

Write ―Eq‖ in the margin next to an important equation.

Underline key steps to solving an example problem, or write those steps in the margin.

Write any questions in the margin.

Write a learning objective next to each subheading.

Using this handout effectively

The studying and annotations and other homework assignments are designed to require a maximum of 45 minutes to an hour of your time. This time limit, however, does not include things like test make-up points or extra credit activities. Students are expected to plan their time. If students wait until the last minute to complete assignments that are assigned multiple days in advance, then the time required to complete the assignment may require more than the 1 hour limit.

This document is a ―work in progress.‖ I have made an attempt to put in as much about the knowledge you need for this unit of study as I can. This document, however, DOES NOT contain everything that you need to know to make 100% on the Unit 4 test. You must

also rely on the knowledge that you should have acquired in previous units of study, classroom explanations, your notes, the textbook sections that were assigned for study, and—in some cases—creative thinking and problem-solving skills.

Molecular Structure and Nomenclature Unit 4: Chemistry I Honors

The Standards in This Unit

According to the South Carolina Science Standard C-3, students will demonstrate an understanding of

the structures and classifications of chemical compounds. Most of the material in this unit falls under

this standard.

Key questions for this next section:

1) What is the oxidation number of a monoatomic ion equal

to?

2) Define a monoatomic ion.

3) What is the oxidation number of monoatomic ions

formed from elements in the group or family of the periodic table labeled column 1 or column IA?


formed from elements in the group or family of the periodic table labeled column 2 or column IIA?

5) What is the oxidation number of monoatomic ions formed from elements in the group or family of

the periodic table labeled column 16 or column VIA?

Circle

Box

Your learning objective for this section:



formed from elements in the group or family of the

periodic table labeled column 17 or column VIIA?

7) What is the difference between a polyatomic and a monoatomic ion?

8) What other kind of particle do polyatomic ions act like?

9) When water autoionizes and there are only water molecules around, what does the hydrogen ion do?

10) When water autoionizes and there are ammonia molecules around, what does the hydrogen ion do?

11) What other name do we give to ions that have a positive charge?

12) What other name do we give to ions that have a negative

charge?

13) How do polyatomic ions often behave in chemical reactions?

Indicator C-3.1 – The charge on an ion

According to Indicator C-3.1 in the South Carolina Science Standards students should be able to

predict the type of bonding (ionic or covalent) and the shape of simple compounds by using Lewis dot

structures and oxidation numbers. Some of the material in this indicator was covered in unit 3. More

information about oxidation numbers and polyatomic ions will be presented in this unit.

In this unit of chemistry students should:

Know that the oxidation number of a monoatomic ion is equal to its charge

Explanation: A monoatomic ion is an atom that has either lost or gained an electron or multiple

electrons and therefore is a charged particle called an

ion.

Example: The sodium atom (symbol Na) typically

looses 1 electron to form a positively charged ion

(positively charged ions are called cations).

Electrons are negatively charged, protons are

positively charged, and an atom has an equal number

of each so that it’s overall charged is 0. The loss of

an electron means that the ion has one more proton

than it has electrons. Since the sodium ion has one

more proton than electrons it has a charge of 1+,

which we simply write a +. The plus charge is

placed in a superscript position to the right of the

symbol for the element. The sodium ion, therefore,

has the symbol Na+.

Know the oxidation number of the monoatomic ions

formed from elements in the following groups of the

periodic table

Group 1 (IA), 1+ (this is written not as 1+ but

simply as +)

What is a learning objective?

A learning objective is an outcome

statement that captures specifically

what knowledge, skills, and attitudes that you should be able to exhibit after

having studied and annotated the

section.

In what form should it be written? ―After studying this section, I will be

able to ____.‖

For example: ―After studying this section I should be able to figure out whether a

substance is ionic or covalent by its

properties and to write and recognize

molecular formulas, empirical

formulas and structural formulas.‖

Polyatomic ions - Covalently bonded

groups of atoms (similar in structure

to molecules) that have a positive or

negative charge. Polyatomic ions act

like monoatomic ions in forming

ionic compounds. Monoatomic ions

are single atoms that have gained or

lost electrons and, therefore. have

obtained a negative charge or a

positive charge. Monoatomic ions,

then have more protons than electrons

or more electrons than protons.

Similarly, when you add up all the

protons and electrons for the entire

polyatomic ion structure, the

polyatomic ion will have more

protons than electrons or more

electrons than protons.


Group 2 (IIA), 2+

Group 16 (VIA), 2 (this almost always applies to binary ionic compounds and sometimes

applies to other kinds of compounds)

Group 17 (VIIA), 1 (this almost always applies to binary ionic compounds and sometimes

applies to other kinds of compounds; this is written not as 1 but simply as )

Understand that some covalently bonded groups of atoms (similar in structure to molecules) act

like single atoms in forming ionic compounds. These charged groups of covalently bonded

atoms are called polyatomic (many-atomed) ions and may be positive or negative.

Formation of these polyatomic ions frequently occurs when a molecule loses one or more

hydrogen ions (H+), leaving the species negatively charged, such as the disassociation of water

into a hydroxide ion (OH-) and a hydrogen ion (H

+)

This equation illustrates the autoionization of water.

The H+ ion also joins with a different water molecule to make a hydronium ion: H3O

+.

water hydrogen

ion hydronium

(a polyatomic ion)

The ammonium ion is formed when a molecule of ammonia, (NH3), combines with a

hydrogen ion, (H+), resulting in a positively charged species: (NH4

+).

ammonia hydrogen

(a gas) ion ammonium

(a polyatomic ion)

A species such as ammonium (above) is called a polyatomic ion. In this particular case, it

is a positively charged polyatomic ion and ions with positive charges are called cations.

Ions with negative charges are called anions

Explanation: Students will be expected to know a few polyatomic ions from memory but

they will be able to use the following list for most polyatomic ions.

H+ + H O

H [ ]

O H

H+ + H O H [ ] +

H O H

H

H+ + H N H

H

+

H N H

H

H


Table 1. This table contains many of the most common polyatomic ions. Students do NOT need to memorize

these. These ions are found on the student’s test references. Students WILL have to learn to recognize and use

these ions in writing formulas and equations.

Polyatomic Anions Ion Name Ion Name Ion Name

CH3COO

or CH3CO2

or C2H3O2

acetate Other ways to write the

acetate ion

H2PO4 dihydrogen phosphate

CO32 carbonate

HCOO formate CrO42 chromate

HCO3 hydrogen carbonate (bicarbonate is a widely used common name)

Cr2O72 dichromate

NH2 amide

HPO42 hydrogen phosphate

C6H5COO benzoate HSO4 hydrogen sulfate

(bisulfate is a widely used common name) MnO4

2 manganate

BrO3 bromate MoO42 molybdate

BrO2 bromite HS- hydrogen sulfide C2O42

(OOCCOO

2-)

oxalate ClO3 chlorate OH hydroxide

ClO2 chlorite ClO hypochlorite O22 peroxide

OCN or CNO cyanate BrO hypobromite C8H4O42 phthalate

CN cyanide IO hypoiodite SiO44

& SiO42 silicate

IO2 iodite SO42 sulfate

Polyatomic Cations IO3 iodate SO32 sulfite

Ion Name NO3 nitrate C4H4O62 tartrate

NH4+ ammonium

NO2 nitrite

S2O32 thiosulfate (thiosulphate)

FeSCN2+

ferrocyanate BrO4 perbromate BO33 borate

Ag(NH3)2+

diamine silver (I)

ClO4 perchlorate Fe(CN)63 hexacyanoferrate (III)

IO4 periodate PO43 phosphate

*Hg22+

(or Hg+) mercury(I)

MnO4 permanganate

PO3

3 phosphite

*Hg2+ mercury(II) SCN– thiocyanate Fe(CN)64 hexacyanoferrate (II)

*Mercury(I) ions always bond together in pairs to form Hg22+

.

Understand that the oxidation number of a polyatomic ion is equal to its charge

Understand that polyatomic ions often react exactly the same as monoatomic ions in

chemical reactions

Explanation: Often in chemical reactions polyatomic ions stay together from one side of the

equation to the other.

Example: In a reaction between barium nitrate—Ba(NO3)2—and sodium sulfate—

Na2SO4—the nitrate ions (NO3 ) and sulfate ions (SO42

) remain intact:

Ba(NO3)2(aq) + Na2SO4(aq) → 2NaNO3(aq) + BaSO4(s)

Notice that there are intact nitrate ions (NO3 ) and sulfate ions (SO42

) on both sides of

the equation, so they remained intact.

Ba(NO3)2(aq) + Na2SO4(aq) → 2NaNO3(aq) + BaSO4(s)

barium nitrate sodium sulfate sodium nitrate barium sulfate

Practice Problems:

1. What is the correct formula for an ion of strontium?


Key questions for this next section: 14) Explain how to use oxidation numbers to find the correct ratio

of atoms in a chemical formula.

15) Explain how to use oxidation numbers to find the correct ratio

of atoms in a chemical formula with polyatomic ions. 16) Explain how to determine if a compound is ionic or covalent

(be complete in your explanation).

17) Explain how to determine the name of a covalent compound (be complete in your explanation).

18) Explain how to determine that name of an ionic compound (be

complete in your explanation).

Indicator C-3.2 – The relationship between non-organic chemical names and formulas

According to Indicator C-3.2 in the South Carolina Science Standards students should be able to interpret the names and formulas for ionic and covalent compounds.

In Physical Science students learned to

Predict the ratio by which the representative elements combine to form binary ionic compounds,

and represent that ratio in a chemical formula. (PS-4.5)

Example:

Magnesium—symbol Mg—has a 2+ oxidation number. In an ionic compound this 2+ charge is

real charge. This makes magnesium— with a 2+ charge—an ion. Since it’s a positively charged

ion it is also called a cation. It’s symbol is Mg2+

.

Bromine—symbol Br—has a 1 oxidation number. In an ionic compound this 1 charge is real

charge. This makes bromine— with a 1 charge— an ion and since it’s a negatively charged ion it

is also called an anion. Its symbol is Br . Note that 1 charges are simply written as and 1+

charges are simply written as +.

Compounds want to combine to be more stable. Isolated charges are inherently unstable, so

cations seek out anions to balance their charge and anions seek out cations for the same reason.

When this magnesium cation is chemically combined with the bromine ion they combine in a ratio

that neutralizes the charges. Two bromines, each with a 1 charge, are needed to neutralize the 2+

charge on the magnesium. So, the formula of a magnesium bromide compound is MgBr2. The

subscript 2 in the formula shows that there are 2 bromines. The subscripts work just like the

coefficients in algebra—they are just written in a different place. Instead of X2Y, it’s written in

the form of XY2, where there is one X and 2 Ys. So, instead of writing Mg2Br to show that there

is 1 Mg and 2 Br symbols, chemists write it MgBr2 to show 1 Mg and 2 Br.

Note: Coefficients are also used in chemistry, but they are used to show the number of unbonded

formula units, molecules, atoms, or other chemical entities. For example, 2H2O means that there

are 2 water molecules, where each water molecule has 2 Hs and 1 O for a total of 4 Hs and 2 Os in

2H2O.

What is your learning objective for this section?


Key questions for this next section: 19) Explain how to use oxidation numbers to find the correct ratio of atoms or monoatomic in a chemical

formula.

20) Explain how to use oxidation numbers to find the correct ratio

of atoms and polyatomic ions in a chemical formula. 21) Explain in a complete manner how to determine if a

compound is ionic or covalent.

22) Explain in a complete manner how to determine the name of a covalent compound from the formula.

23) Explain in a complete manner how to determine the name of

an ionic compound from the formula.

In this unit of chemistry students should be able to:

Name and write the chemical formulas for binary molecular compounds

All binary compounds (with a few exceptions), whether molecular or ionic, will end with an ―-ide.‖

The number of atoms in non-organic molecules are indicated with prefixes (organic molecules

contain at least 2 carbon atoms covalently bonded to each other). Molecules are always bonded

covalently and therefore almost always are made of 2 or more non-metals. The prefixes are:

mono – 1 (this is very rarely used; primarily for ―carbon monoxide‖—symbol CO)

di – 2

tri – 3

tetra – 4

penta – 5

hexa – 6

hepta – 7

octa – 8

nona – 9

deca – 10

Example: Water is H2O. But it’s systematic name would be ―dihydrogen oxide.‖ Ammonia is

NH3. But it’s systematic name is ―nitrogen trihydride.‖ N2O5 is ―dinitrogen pentoxide.‖

Later we will learn to identify binary ionic compounds and covalent compounds by the difference

in electronegativity. For now, use these general rules:

When metals and non-metals bond together they form ionic compounds.

In this class, any formula unit that contains a polyatomic ion will also be an ionic

compound. This can happen when:

a metal combines with a polyatomic anion

a polyatomic cation combines with a non-metal

a polyatomic cation combines with a polyatomic anion

When 2 non-metals bond together, they form covalent compounds. Most covalent

compounds form units called molecules although they sometimes form extended lattice

structures called covalent network crystals. Diamonds and quartz glass are covalent

network crystals.

Another category of covalent compounds is organic compounds. The general rule is that

covalent compounds that include carbon atoms bonded to other carbon atoms is an organic

compound.

Name and write the chemical formulas for ionic compounds including those that contain common

polyatomic ions

Explanation: All binary compounds (with a few exceptions), whether molecular or ionic, will end

with an ―-ide.‖ Polyatomic ions are named what they are according to the table of common ions.

Metals in ionic compounds are named for their element name except when they have more than 1

oxidation number. Metals in ionic compounds that have more than one oxidation number must

have a Roman numeral in parentheses after the element name to designate which ion it is.



Binary ionic compounds: Binary ionic compounds are typically metals (which will be positively

charged and are, therefore, cations) bonded with non-metals. If the metal cation has only one

oxidation number (see the periodic table of oxidation numbers), then the metal is named for the

element. The non-metal is named for the element but with an ―-ide‖ ending.

Example: CaCl2 is named for the metal, calcium, and the non-metal with an ―-ide‖ ending,

chloride. The compound name is, therefore, calcium chloride.

Larger than binary ionic compounds:

Ionic compounds that have more than 2 elements also have polyatomic ions. To name these you

must 1st identify the polyatomic ions. You can find them in the table of common polyatomic ions.

The charges on the polyatomic ions do not appear in the formula for the compound because a

properly balanced compound with have an equal number of positive and negative charges which

cancel each other out. Polyatomic ions are named just as they are in this table.

Metals bonded to polyatomic anions: If the polyatomic ion is an anion (a negatively charged ion)

bonded to a metal, then the same rules apply to naming it. A metal that only has one oxidation

number is named for its element name but a metal that has more than one oxidation number must

have a Roman numeral in parentheses to indicate which metal ion it is.

Example #1: Barium has only one oxidation number (a 2+). If it is bonded to a chlorate ion

(ClO4‾), then the formula is Ba(ClO4)2 and its name is barium chlorate.

Example #2: Iron has 2 oxidation numbers (a 2+ and a 3+). If iron(III)—symbol Fe3+

—is

bonded to a sulfite ion (SO32‾), then the formula is Fe2(SO3)3 and its name is iron(II) sulfite.

Polyatomic cations bonded to non-metals: If the polyatomic ion is a cation (a positively charged

ion) bonded to a non-metal, then the non-metal will be named for its element but with an ―-ide‖

ending.

Example #1: Ferrocyanate is FeSCN2+

. If it is bonded to a sulfur then that sulfur’s name is

changed to have an ―-ide‖ ending, ―sulfide,‖ which has a symbol S2–

. The formula is FeSCNS

and its name is ferrocyanate sulfide.

Polyatomic cations bonded to polyatomic anions: If the ionic compound is made of 2 polyatomic

ions (a cation and an anion), then the compound is simply named for the 2 polyatomic ions.

Example #2: Ammonium is NH4+. If it is bonded to a thiocyanate ion (SCN‾), then the formula

is NH4SCN and its name is ammonium thiocyanate.

Before we leave the subject of naming compounds (nomenclature) let’s learn to name non-

organic acids. Naming acids isn’t really in the standards until much later, but I find it easier for

students to learn the basic rules now while we’re on the subject of naming.

Acids are usually (but not always) written with a hydrogen on the front of the formula.

Example: H2SO4 is sulfuric acid.

To name an acid, name it 1st as an ionic compound. Binary compounds such as H2S would be

named hydrogen sulfide. Compounds with polyatomic ions such as H2SO4 would be named

hydrogen sulfate. Compounds with polyatomic ions such as H2SO3 would be named hydrogen

sulfite.

Binary acids – The ―hydrogen‖ in the ionic compound name is changed to ―hydro-‖ and is

attached as a prefix to the anion name. The ―-ide‖ suffix in the anion name is changed to the

suffix ―-ic.‖ The word ―acid‖ is attached as a second name. So, H2S which was named as the


ionic compound hydrogen sulfide now becomes hydrosulfuric acid. Another example would be

HCl. As an ionic compound it would be named hydrogen chloride and as an acid it would be

changed to hydrochloric acid.

Acids with polyatomic anions – The ―hydrogen‖ in the ionic compound name is dropped

altogether. The ―-ate‖ suffix in the anion name is changed to the suffix ―-ic‖ and the ―-ite‖

suffix in the anion name is changed to the suffix ―-ous.‖ The word ―acid‖ is attached as a

second name. So, H2SO4 which was named as the ionic compound hydrogen sulfate now

becomes sulfuric acid. H2SO3 which was named as the ionic compound hydrogen sulfite now

becomes sulfurous acid. Another example would be H3PO4. As an ionic compound it would

be named hydrogen phosphate and as an acid it would be changed to phosphoric acid.

A note about oxidation numbers:

If given a formula, sometimes you have to do a little math to figure out the oxidation number on a

particular atom. For example, in sulfuric acid (H2SO4) the hydrogen always has a 1+ charge unless

it’s called ―hydride.‖ You can almost always assume that oxygen has a 2 charge. Since there are

2 hydrogen atoms, that accounts for a 2+ charge. Since we have 4 oxygen atoms, that accounts for

an 8 charge. In order for the compound to have a neutral charge, the charge on sulfur must be 6+.

The math can be written like this:

2(H+) + 1(S

6+) + 4(O

2) is neutral because 2(1+) + 1(6+) + 4(8 ) = O.

Identify substances as molecular or ionic compounds.

Explanation: Compounds that only contain non-metals will typically be molecular. Compounds

that contain metals bonded with non-metals, metals bonded with polyatomic anions, polyatomic

cations bonded with non-metals, or polyatomic cations bonded with polyatomic anions will be

ionic compounds.

Practice Problems: 2. What is the correct formula for an ion of strontium?

3. What class of compounds would P2S3 fall into: ionic or covalent?

4. What class of compounds would Cr3O2 fall into: ionic or covalent? 5. What class of compounds would N2O5 fall into: ionic or covalent?

6. What class of compounds would Al2(SO3)3 fall into: ionic or covalent?

7. What class of compounds would (NH4)2CO3 fall into: ionic or covalent? 8. Name this compound: P2S3.

9. Name the compound N2O5.

10. Name the compound Na2Se.

11. Name the compound Cr3O2.

12. Name the compound Al2(SO3)3.

13. Name the compound (NH4)2CO3.



24) Typically, how are molecular compounds different from ionic compounds in electrical conductivity in aqueous solution?

25) Typically, how are molecular compounds different from ionic compounds in electrical conductivity when they are molten (in

liquid form)?

26) Typically, which is harder: a molecular compound or an ionic compound?

27) Typically, which is more brittle: a molecular compound or an ionic compound?

28) Typically, which has a higher boiling/condensation point: a molecular compound or an ionic compound?

29) Typically, which has a higher melting/freezing point: a molecular compound or an ionic compound?

30) What is the difference between a molecular formula, an

empirical formula, and a structural formula?

Properties of molecular and ionic compounds

Chemistry students should be able to compare molecular

and ionic compounds according to their properties

Explanation:

Electrical conductivity of the compound in aqueous

solution.

An aqueous solution is one in which water is the

solvent (water is dissolving the other substance or substances). For example, if sugar is

dissolved in water the result is an aqueous solution of sugar.

Ionic Molecular

Ionic compounds typically conduct electricity

when dissolved in water, because the dissociated

ions can carry charge through the solution.

Molecular compounds don't dissociate

into ions and so don't conduct electricity

in solution. Note: Recall that the word ―typically‖ infers that there are exceptions to the rule.

Electrical conductivity of the compound in liquid form.

Ionic Molecular

Ionic compounds conduct

electricity well when melted.

Covalent molecular compounds typically do not

conduct electricity very well, because they usually

don't transfer electrons unless they react. Theory: molten ionic compounds can

transfer electrons

Theory: molten covalent compounds do not transfer electrons

Note: Metallic solids conduct electricity even when they are in a solid state.


What is a learning objective?

A learning objective is an outcome statement that captures specifically

what knowledge, skills, and attitudes

that you should be able to exhibit after having studied and annotated the

section.

In what form should it be written? ―After studying this section, I will be

able to ____.‖

For example: ―After studying this section I should

be able to figure out whether a

substance is ionic or covalent by its properties and to write and recognize

molecular formulas, empirical

formulas and structural formulas.‖


Hardness.

Ionic Molecular

Ionic crystals are harder but often quite brittle. Squeezing an ionic

crystal can force ions of like charge in the lattice to slide into

alignment; the resulting electrostatic repulsion splits the crystal.

Molecular solids are

usually much softer than

ionic materials. Note: there are covalently bonded crystals—called covalent network crystals—that are even harder than ionic crystalline solids. The difference is that these are crystals and not molecules.

Melting points and boiling points.

Melting and boiling points are the temperatures at which substances start to melt or start to boil.

Ionic Molecular

In an ionic compound, the melting

and boiling points are usually much

higher than covalent compounds.

In covalent molecular compounds, the melting and

boiling points are usually much lower than for ionic

compounds.

Theory: The forces of attraction between

positive and negative ions are strong and

high temperatures are required to

overcome them.

Theory: A smaller amount of energy is required to overcome the

weak attractions between covalent molecules.

Note: Many compounds in this class are liquids or gases at room

temperature.

Enthalpies of fusion and vaporization.

The enthalpy of fusion is the amount of heat required to melt one mole of the compound in

solid form, under constant pressure. The enthalpy of vaporization is the amount of heat required

to vaporize one mole of the compound in liquid form, under constant pressure.

Ionic Molecular

Enthalpies of fusion and vaporization are

typically 10 to 100 times larger for ionic

compounds than they are for molecular

compounds.

Enthalpies of fusion and vaporization are

typically 10 to 100 times smaller for

molecular compounds than they are for ionic

compounds.

Theory: The forces of attraction between positive and

negative ions are strong and high temperatures are required to overcome them.

Theory: A smaller amount of energy is required to

overcome the weak attractions between covalent molecules.

Practice Problems: For these 2 compounds, PCl3 and ScCl3, answer the questions below and explain

why you know your answer to be true. 14. Which substance is more likely to conduct electricity when dissolved in water? Explain.

15. Which substance is more likely to conduct electricity when molten? Explain. 16. Which substance is more likely to form a harder substance when in solid form? Explain.

17. Which substance is more likely to have a higher melting point? Explain.

18. Which substance is more likely to have a higher boiling point? Explain. 19. Which substance is more likely to have a higher enthalpy of fusion? Explain.

20. Which substance is more likely to have a higher enthalpy of vaporization? Explain.


Differentiate and write molecular formulas, empirical formulas and structural formulas.

Formula

Type

How they are

different Examples

Molecular

formulas

show the actual

number of atoms in

a molecule.

Propane’s molecular formula is C4H10. There are 4

atoms of carbon (C) and 10 atoms of hydrogen (H) in

each molecule of ethane.

Empirical

formulas

show the lowest

whole number ratio

of atoms in a

molecule.

Propane’s molecular formula is C4H10, but its empirical

formula is C2H5. The lowest whole number ratio of the

atoms in propane is 2 atoms of carbon (C) for every 5

atoms of hydrogen (H) in each molecule of ethane.

Structural

formulas

show the way that

atoms are

connected together

in a compound and

especially in a

molecule or

polyatomic ion.

There are several kinds of structural formulas. The

Lewis formula for propane (also called the complete

structural formula) is:

C C

H

H

H

H

C

H

H

H

C

H

H

H

. Recall: These kinds of drawings were discussed in the last unit. A

Lewis formula is also called a Lewis structural formula, a

complete structural formula, or a bond-line formula.

More condensed but complete structural formulas are: CH

3CH

2

CH

2

CH3

CH3 CH2 CH2 CH3 CH3CH2CH2CH3

A skeletal structural formula is:

A skeletal structural formula is also called a line

drawing. An organic chemist looks at the line drawing

and ―sees‖ a carbon with hydrogen atoms at the end of

every line or at a corner. Since carbon needs 4 bonds to

complete its valence level octet of electrons, a carbon

on the end of the skeletal drawing would have 3

hydrogen atoms bonded to it and a carbon at the corner

of the drawing above would have 2 hydrogen atoms

bonded to it.

Practice Problems:

List of compounds for the following problems: P2S3, Cr3O2, N2O5, Al2(SO3)3, and (NH4)2CO3.

21. Which compounds above would be likely to conduct electricity in aqueous solution?

22. Which compounds above would be likely to conduct electricity in molten (melted or liquid) form?

23. Create 2 groups, harder compounds and softer compounds. 24. Create 2 groups, compounds with higher melting points and compounds with lower melting points.

25. Create 2 groups, compounds with higher boiling points and compounds with lower boiling points.

26. Create 2 groups, compounds with higher enthalpies of fusion and compounds with lower enthalpies of

fusion.

27. Create 2 groups, compounds with higher enthalpies of vaporization and compounds with lower

enthalpies of vaporization.



31) What range of electronegativity typically produces a non-polar covalent bond?

32) What range of electronegativity typically produces a polar

covalent bond?

33) What range of electronegativity typically produces an ionic bond?

34) How do you determine the END of a bond?

35) Why do most substances with ionic bonds usually have high melting and boiling points when compared to substances with most substances wit covalent bonds?

36) What is the percent ionic character in bonds between identical non-metals (diatomic compounds)?

37) What is the range of percent ionic character in bonds between non-metal atoms that are not identical?

38) What symbol is used to show that an atom has a significantly stronger attraction for the electrons than the other atom with which it is covalently bonded?

39) What symbol is used to show that an atom has a significantly weaker attraction for the electrons than the other atom with which it is covalently bonded?

40) What creates a polar bond?

41) What creates a non-polar bond?

42) What creates a polar molecule?

43) What creates a non-polar molecule when the bonds within that molecule are polar?

44) What kinds of molecules have significant attractions for each other?

45) What kinds of molecules have significant attractions for ionic compounds?

46) Explain how to determine when a molecule is polar and when it is non-polar.

Indicator C-3.3 – The range of bond characteristics

According to Indicator C-3.3 in the South Carolina Science Standards students should be able to

explain how the types of intermolecular forces present in a compound affect the physical properties of

compounds (including polarity and molecular shape).

Students should be able to understand that ionic bond and covalent bonds are relative terms and

that most bonds that we characterize as ionic or covalent actually have a character that lies

somewhere between 100% ionic and 100% covalent.

WARNING! What you are about to discover about electronegativity differences is about

individual bonds between 2 atoms and NOT about the entire molecule. Later we will see how the

electronegativity differences between different atoms in a molecule contribute to the overall

polarity of a molecule.

Students should understand how the electronegativity difference can be used to classify the type of

bond in a substance

Explanation: We will consider electronegativity differences (END) from 0 to 0.3 to be non-polar

covalent. Electronegativity differences between 0.5 and 1.5 will be considered polar-covalent.

Electronegativity differences greater than 1.7 will be considered ionic. To determine the percent

ionic character of a bond, therefore, you divide the electronegativity difference of the bond by 1.7

and multiply the answer by 100.



0 0.3 0.5 1.5 1.7

Electronegativity difference 0 to 0.3 0.3 to 0.5 0.5 to 1.5 1.5 to 1.7 1.7 to ∞

Classification of bond non-polar

covalent

unclear polar-covalent unclear ionic

Figure 4.1. This illustration shows the range of electronegativity differences in compounds.

Example: The electronegativity difference between nitrogen and chlorine is 0.12. This was

determined by subtracting the electronegativity of nitrogen (3.04) from the electronegativity of

chlorine (3.16): 3.16 – 3.04 = 0.12. The percent ionic character of the bond between nitrogen

and chlorine, therefore, is 7.1%. Here’s the math:

0.12 elecronegativity difference between nitrogen and chlorine = 0.0705882 0.071.

1.7 elecronegativity difference for an ionic compound

0.071 × 100 = 7.1% ionic character.

Bonds between active metals and active nonmetals are characterized by a high degree of ionic

character because electron transfer is virtually complete.

Because ionic bonds are very strong, substances with ionic bonds usually have high melting

and boiling points.

Bonds between identical non metals (diatomic compounds) are characterized by zero percent

ionic character because electrons are shared equally.

Explanation: These bonds have NO polarity. There is NO difference between the

electronegativity values assigned to the 2 atoms in the bond.

Example: Oxygen in the atmosphere is mostly O2. Since both oxygen atoms have the same

electronegativity value (3.44) (see the periodic table of electronegativities) there is no

electronegativity difference and therefore a 0 percent ionic character.

Bonds between other substances (such as the bond between oxygen and hydrogen) have an

intermediate nature; the shared electrons are not shared equitably but spend more time with

whichever atom is more electronegative.

The atom with the stronger attraction for electrons becomes partially negatively charged

( ).

Explanation: In your table of electronegativity values oxygen is assigned a value of 3.44

whereas hydrogen is assigned a value of 2.1. The oxygen in a covalent bond with hydrogen

is the more electronegative and the difference (3.44 – 2.1 = 1.34 1.3) is in the highly

polar range. An oxygen to hydrogen bond is, therefore, highly polar with the oxygen end

being the partially negative end and the hydrogen being the partially positive end.

Note: The symbol for partially positive is + and the symbol for partially negative is

–. A

particle is completely positive or negative when it is an ion and, indeed, may have positive

and negative values of 1+ (which is symbolized simply at +), 2+, 3+, … n+.

The atom with the lower electronegativity value becomes partially positively charged (+).

Covalent bonds that do not share the electrons equally are called polar covalent bonds

non-polar covalent

unclear polar covalent unclear ionic


Explanation: Electronegativity is a complex concept. The structure of a molecule affects

the electronegativity such that the electronegativity between 2 atoms is different in one

molecular structure than it is in another. For simplicity in this class, we will assume that

electronegativity differences from 0 to 0.3 are considered non-polar covalent,

electronegativity differences between 0.3 and 1.7 are considered polar-covalent, and

electronegativity differences greater than 1.7 are considered ionic.

Note: As was stated, other factors in the structure of a compound contribute to polarity

and/or non-polarity, so bonds with electronegativity differences slightly above 0.3 could

still behave as non-polar bonds. Bonds with electronegativity differences slightly less than

1.7 could still be ionic and those slightly above 1.7 could still be polar-covalent. For

now—in THIS unit—assume that those exceptions do not exist but PLEASE try to keep it

in your mind that these exceptions exist and DO NOT try to tell your college professor that,

―My high school chemistry teacher told us that ANYTHING below an electronegativity

difference of 1.7 was ALWAYS polar covalent.‖

Covalent bonds that do share the electrons relatively equally are called non-polar covalent

bonds.

Recall: Electronegativity differences from 0 to 0.3 are considered non-polar covalent, so

electronegativity differences from 0 to 0.3 are considered ―relatively equal.‖

If the polar bonds in a molecule are all alike, the polarity of the molecule as a whole

depends only on the arrangement in space of the bonds (water molecules are polar due to

bent structure). See Table 4.4 for some generalized examples. An even distribution of

polar bonds creates a non-polar molecule. An uneven distribution of polar bonds creates a

polar molecule. See figures 4.2 through 4.5 for examples.

Examples:

Figure 4.2. If we use an arrow to show the

direction of the greatest electronegativity in

each of the bonds in CCl4 you can see that the direction of END is equally spaced and

is equally spread around the carbon.

Further, the directions of those ENDs are all away from carbon. Like a tug-o-war in

which neither team is winning, the equal

pull of the END cancels out the polarity of the molecule.

Figure 4.3. The bonds between carbon and chlorine in CCl4 have an

electronegativity difference (END) of

0.61. This END makes the carbon-chlorine bond polar with chlorine pulling

the electrons in the bond harder than

carbon does. Chlorine, therefore, is

partially negative (–) and carbon is

therefore partially positive (+).

C

Cl

Cl Cl Cl

–

– –

– C

Cl

Cl Cl

Cl

Nonpolar

molecules

that have

polar

bonds


Add in illustrations about END running in the same direction.

Explanation: Any uneven arrangement of polar bonds can result in a polar molecule.

Polar molecules are attracted to one another, but the attraction is not a chemical bond so it

is broken easily. These substances usually have moderate melting and boiling points.

Note: You can think of polar molecules like tiny little magnets that are attracted to each

other but only if the north pole is aligned with the south pole. A polar molecule’s partially

positive (+) end is attracted to another polar molecule’s negative end (

–).

Polar molecules are attracted to one another and to ionic substances as well.

Example: A polar molecule’s partially positive (+) end is attracted to a negative ion (an

anion) such as nitrate (NO3–). A polar molecule’s negative end (

–) is attracted to a

positive ion such as iron(III) (Fe3+

).

Note: This is why polar liquids (such as water) can dissolve ionic compounds (such as table salt).

Practice Problems:

Determine whether these molecules are polar or non-polar. Do as many as you need to feel

comfortable with your ability to be successful.

Note: You must first determine the molecular shape of the molecule, then determine if the bonds are polar or non-polar, and

lastly determine if the polar bonds are unevenly distributed. If the bonds are all non-polar, then the molecule is non-polar.

If there are polar bonds but they are evenly distributed, then the molecule is non-polar. If there are polar bonds and they are UNevenly distributed, then the molecule is unusually polar. Things can get a little tricky if you have more than one kind of

bond around the central atom in the molecule. You will then need to consider the overall direction of the electronegativity

to determine if the molecule is polar or non-polar.

28. H2S

29. PCl3

30. BF3

31. HCN

32. PF3

33. CH4

34. CCl4

35. SiH4

36. NH3

37. H2O

38. OF2

39. SF2

40. CO2

41. C2H2

Figure 4.4. The bonds between nitrogen

and fluorine in NF3 have an

electronegativity difference (END) of 0.94. This END makes the nitrogen-

fluorine bond polar with fluorine pulling

the electrons in the bond harder than

nitrogen does. Fluorine, therefore, is

partially negative (–) and nitrogen is

therefore partially positive (+).

Figure 4.5. If we use an arrow to show the

direction of the greatest electronegativity in

each of the bonds in NF3 you can see that the direction of END is NOT equally spaced

and is equally spread around the nitrogen.

The directions of those ENDs are all away

from nitrogen. Like a tug-o-war in which one team is winning, the UNequal pull of

the END creates polarity in the molecule.

N

F F

F – –

– N

F F F

Polar

molecules


Answers to Practice Problems:

Determine whether these molecules are polar or non-polar. Do as many as you need to feel

comfortable with your ability to be successful.

21. Polar

22. Polar

23. Non-polar

24. Polar

25. Polar

26. Non-polar

27. Polar

28. Polar

29. Polar

30. Polar

31. Polar

32. Polar

33. Non-polar

34. Non-polar

35. Key questions for this next section:

47) What is electronegativity?

48) What is the general pattern of electronegativity of the elements as they are arranged on the periodic table?.

Indicator C-3.8 – The effect of electronegativity and

ionization energy on the type of bonding in molecule

According to Indicator C-3.8 in the South Carolina Science

Standards students should be able to explain the effect of

electronegativity and ionization energy on the type of bonding in a molecule.


Infer relative electronegativity values for elements based on the element’s position on the periodic

table.

You should recall that the highest electronegativity values for elements is found in the upper right

corner of the periodic table if we don’t include the noble gases. This was covered in the previous

unit of instruction.

H

2.1

Periodic table of the elements with electronegativities. Note: No values are shown for those elements whose electronegativities are not known.

He

0

Li

0.98

Be

1.57

B

2.04

C

2.55

N

3.04

O

3.44

F

3.98

Ne

0

Na

0.93

Mg

1.31

Al

1.61

Si

1.9

P

2.19

S

2.58

Cl

3.16

Ar

0

K

0.82

Ca

1

Sc

1.36

Ti

1.54

V

1.63

Cr

1.66

Mn

1.55

Fe

1.83

Co

1.88

Ni

1.91

Cu

1.9

Zn

1.65

Ga

1.81

Ge

2.01

As

2.18

Se

2.55

Br

2.96

Kr

0

Rb

0.82

Sr

0.95

Y

1.22

Zr

1.33

Nb

1.6

Mo

2.16

Tc

1.9

Ru

2.2

Rh

2.28

Pd

2.2

Ag

1.93

Cd

1.69

In

1.78

Sn

1.96

Sb

2.05

Te

2.1

I

2.66

Xe

2.6

Cs

0.79

Ba

0.89

La

1.1

Hf

1.3

Ta

1.5

W

2.36

Re

1.9

Os

2.2

Ir

2.2

Pt

2.28

Au

2.54

Hg

2

Tl

2.04

Pb

2.33

Bi

2.02

Po

2

At

2.2

Rn

0

Fr

0.7

Ra

0.89

Ac

1.1

Rf Db Sg Bh Hs Mt Ds Rg Uub Uuq Uuh

Ce

1.12

Pr

1.13

Nd

1.14

Pm

1.13

Sm

1.17

Eu

1.2

Gd

1.2

Tb

1.1

Dy

1.22

Ho

1.23

Er

1.24

Tm

1.25

Yb

1.1

Lu

1.27

Th

1.3

Pa

1.5

U

1.38

Np

1.36

Pu

1.28

Am

1.3

Cm

1.3

Bk

1.3

Cf

1.3

Es

1.3

Fm

1.3

Md

1.3

No

1.3

Lr



You should, therefore, be able to estimate relative

electronegativities by the relative position of elements on the

periodic table. In other words, because phosphorous is above

and to the right of germanium, you should know by their

position that phosphorous has a higher electronegativity value

than germanium.

You should, therefore, be able to estimate relative electronegativities

by the relative position of

elements on the periodic

table. In other words,

because phosphorous is

above and to the right of

germanium, you should

know by their position that

phosphorous has a higher

electronegativity value than

germanium:

Use a table of

electronegativity values

to assign values to

elements represented in

the structural formula of

a substance.

See the ―Periodic table

of the elements with electronegativities‖

above.

Determine the percent ionic character

of a bond based on the

electronegativity difference of the

elements involved

Percent ionic character is simply another

way of saying the difference in

electronegativity and some other factors

that affect the ability of one atom to

remove the electrons from another. In

general there are no bonds that are 100%

ionic or 100% covalent. There is always

some small component of both ionic and

covalent character in every bond. In

compounds we think of as ionic (such as

metallic salts) the covalent component is

so small that we can ignore it. In a diatomic molecule such as H2 or Cl2, the amount of ionic

component is also so small that it can be ignored. In other compounds there may be substantial

contributions from both the ionic component and the covalent component. For example, HF (a

P

2.19

Ge

2.01

Figure 4.6. Note the relative

positions of phosphorous and

germanium on the periodic

table.

12 3 4

5 6 78 9 10 11

12 13 14 1516 17 18

S1

S2

S3

S4

S5

S6

S7

S8

S9

S10

0

0.5

1

1.5

2

2.5

3

3.5

4

Electronegativity

Groups or Families

Periods

Periodic Table of Electronegativity

Figure 4.7. Three dimensional graph showing the relative

electronegativity of the elements.

Figure 4.8. Outline of the periodic table showing the

relative electronegativity of the elements. Note:

Electronegativity does NOT include the noble gases

(except krypton and xenon).

Higher

electronegativity

Lower

elecronegativity


weak acid and sometimes behaves like an ionic compound and sometimes behaves like a covalent

compound) has 41% ionic character and 59% covalent character. The factors that influence ionic

and covalent character are complex and many of them are too complex for this level of chemistry.

You will only be expected to say that ionic compounds have a high percentage of ionic character

and a small percentage of covalent character, highly polar molecules will have a high percentage of

both ionic and covalent character, and non-polar covalent compounds will have a high percentage

of covalent character and a low percentage of ionic character.

For example, the electronegativity difference between hydrogen and fluorine is is about 1.9 and the

percent ionic character is about 41%. In general however, an electronegativity difference of 1.6

gives about 51% percent ionic character. On the other hand KBr which an electronegativity

difference of 2.0 has 78% ionic character. This difference can be attributed to the ability of one

atom to distort the electrical field of the other atom. In KBr we see less distortion that we do with

HF.

In the section below, students should be able to

infer relative ionization energy values for elements based on the element’s position on the periodic

table.

use a table of ionization energy values to assign values to elements represented in the structural

formula of a substance.

understand how the relative ionization energies of two elements can be used to predict the type of

bonding that form between them.

Just as a large difference in electronegativity can be used to predict the type of bonding between 2

atoms, ionization energies can be used for this purpose as well. We have only looked at 1st

ionization energy, which is the energy that must be added to an atom to get it to release one

electron (see the ―Periodic table of the elements with 1st ionization energies‖ that follows). Getting

an atom to release more electrons requires increasing amounts of energy.

Since ionic bonding between 2 atomic species is the result of one species giving up one or more

electrons and another gaining one or more electrons, it should be easy to see that atoms with a high

ionization energy are much less likely to give up electrons and those with a low ionization energy

are much more likely to give up

electrons.

As a result, if 2 atoms have

large differences in ionization

energy, they are likely to form

an ionic bond.

Conversely, if 2 atoms have

high ionization energy but a

small difference in ionization

energy, they are likely to hold

onto their electrons but would

more likely share them to get a

more stable electron

configuration.

Figure 4.9. Outline of the periodic table showing the

relative first ionization energies of the elements. Note:

First ionization energies DOES include the noble gases.

Higher 1st

ionization energy

Lower 1st

ionization energy


Also, if 2 atoms have low ionization energy but a small difference in ionization energy,

they are likely to give up their electrons and share them in what in a kind of electron soup.

This electron soup is the theory behind the ―electron sea model‖ for the structure and

behavior of metals.

H

13.5

Periodic table of the elements with 1st ionization energies (in eV).

Note: No values are shown for those elements whose 1st ionization energies are not known. He

24.6

Li

5.4

Be

9.3

B

8.3

C

12.3

N

14.5

O

13.6

F

17.4

Ne

21.6

Na

5.1

Mg

7.6

Al

6.0

Si

8.2

P

10.5

S

10.5

Cl

13.8

Ar

15.8

K

4.3

Ca

6.1

Sc

6.5

Ti

6.8

V

6.7

Cr

6.8

Mn

7.4

Fe

7.9

Co

7.8

Ni

7.6

Cu

7.7

Zn

9.4

Ga

6.0

Ge

7.9

As

9.8

Se

9.8

Br

11.8

Kr

14.6

Rb

4.2

Sr

5.7

Y

6.4

Zr

6.8

Nb

6.9

Mo

7.1

Tc

7.3

Ru

7.4

Rh

7.5

Pd

8.3

Ag

7.6

Cd

9.0

In

5.8

Sn

7.3

Sb

8.6

Te

9.0

I

10.5

Xe

12.1

Cs

3.9

Ba

5.2

La

5.6

Hf

6.7

Ta

7.9

W

8.0

Re

7.8

Os

8.7

Ir

9.1

Pt

9.8

Au

9.2

Hg

10.4

Tl

6.2

Pb

7.4

Bi

7.3

Po

8.4

At

Rn

10.7

Fr

Ra

5.3

Ac

5.2

Rf Db Sg Bh Hs Mt Ds Rg Uub Uuq Uuh

Ce

5.5

Pr

5.4

Nd

5.5

Pm

5.6

Sm

5.6

Eu

5.7

Gd

6.2

Tb

5.8

Dy

5.9

Ho

6.9

Er

6.1

Tm

6.2

Yb

6.3

Lu

5.4

Th

6.1

Pa

5.9

U

6.2

Np

6.2

Pu

6.1

Am

6.0

Cm

6.0

Bk

6.2

Cf

6.3

Es

6.4

Fm

6.5

Md

6.6

No

6.7

Lr


Using molecular geometry and bond polarity to determine molecular polarity


49) What combination of molecular shapes and polar bonds results in a polar molecule?

Students should be able to interpret the polarity of a

molecule based on its geometry bond type.

Explanation: In the last unit students were taught a

system for determining molecular shape. Molecular

shape is one of 2 major factors in determining whether or

not a molecule is polar. The other major factor is the

polarity of all the bonds in the molecule.

If all the bonds in a molecule are non-polar, then the molecule is typically non-polar.

If all the bonds in a molecule are polar but the polar bonds are all evenly distributed around the

central atom or atom of interest, then the molecule is non-polar.

On the other hand if all the bonds in a molecule ARE polar but the bonds are NOT all evenly

distributed around the central atom or atom of interest, then the molecule IS likely to be polar.

If the atoms that are tied to the central atom are different from

each other and some of the bonds are polar and some are non-

polar, then you must look at the total polar imbalance of the

molecule to determine polarity. If a molecule which has

polar bonds on one side, but non-polar bonds on the other, the

molecule is likely to be polar. See figure 4.10 for another

example.

See the table of simple polar and non-polar molecules that

follows. Note that this table is simplified. There are

circumstances that can lead to polar molecules that were not

included in the table. Students must use some common sense as

well as knowledge of electronegativity in determining the overall

polarity or lack of polarity of a molecule.

Figure 4.10. For the molecule

above, the electronegativity

difference (END) in each

bond is in the unclear range

between polar and non-polar

(see figure 4.1). In a

molecular arrangement such

as this with the arrows turned

in the same direction, the

END for the 2 bonds are

additive, creating an overall

electronegativity for the

molecule of 0.94. This

molecule, therefore, is polar.

C N H END = 0.45 END = 0.49

Overall END = 0.94



Examples:

Table 4.1. Selected examples of simple polar and non polar molecules.

NON-POLAR POLAR POLAR Polar bonds ↔ Non-polar molecule Polar bonds ↔ Polar molecule A-B non-polar bonds; A-C polar bonds ↔ Polar molecule

Even if the bonds in these molecular shapes are the same and are polar, the molecule is NOT polar because

the bonds are evenly distributed around the central atom or atom of

interest.

If the bonds (A-to-B) in these molecular shapes are polar, then the molecule is polar

because the bonds are NOT all evenly distributed around the central atom or atom

of interest.

If the bonds between the central atom (A) and atom B are non-polar, but the bonds between the central atom (A) and atom C are polar, then the molecule is

polar because the polarity of the bonds is NOT evenly distributed around the central atom or atom of

interest.

A B B

Bent

molecular

shape

A

B

B

Linear

molecular

shape

A

B

B

B

Trigonal

pyramidal

molecular

shape

B

A

B B

Trigonal

planar

molecular

shape

B

A

B B

B

Bent

molecular

shape

B

A

B

Trigonal

planar

molecular

shape

B

B

B

A B

B

Irregular

tetrahedron

molecular

shape

B

B

B

A

B

Tetrahedral

molecular

shape

B

B

B

B

B

A

B

T-shaped

planar

molecular

shape

B

A

B

B

Tetrahedral

molecular

shape

B

B B

A

B

Square

pyramidal

molecular

shape

B

B

B

B

A

B

Tetrahedral

molecular

shape

Note: the central atom or atom of interest in these illustrations are labeled A.



50) What are allotropes?

51) What is a prominent example of allotropes?

52) An atom with a total of 4 single bonds plus lone pairs is

likely what hybridization?

53) An atom with one double bond and a total of 2 single

bonds plus lone pairs is likely what hybridization?

54) An atom with a triple bond or 2 double bonds is likely

what hybridization?

55) What is the hybridization of the carbon atoms in graphite?

56) What is the hybridization of the carbon atoms in a diamond?

57) What is the hybridization of a carbon atom in a carbon ring structure with all single bonds?

58) Explain the unique bonding

characteristics of carbon.

Indicator C-3.5 – Unique bonding

characteristics of carbon

According to Indicator C-3.5 in the

South Carolina Science Standards

students should be able to explain the

unique bonding characteristics of

carbon that have resulted in the

formation of a large variety of organic

structures.

In this unit of chemistry students should

be able to:

Understand bonding in the allotropic

forms of carbon (see figure 4.6):

diamond and graphite (and it’s

always fun to know to know about

―bucky balls‖).

Explanation: Allotropes for different

forms of the same element. The

difference is in how the atoms of the

pure element is structured or bonded

together.

Example: Graphite is a form of

carbon that comes in six sided 2

dimensional shapes all linked

together to make sheets.

Diamond has a much more

complex 3-dimension structure

that is very hard to illustrate on

Figure 4.11. A photo of a diamond (upper left) is seen in

this set of illustrations immediately above an illustration of

small section of the geometric arrangement of carbon

atoms in a diamond (lower left). A photo of graphite

(upper right) are seen in this set of illustrations

immediately above an illustration of small section of the

geometric arrangement of carbon atoms in the benzene ring

sheets that are the carbon structure of graphite (lower

right). This file is a composite of Image:GraphiteUSGOV.jpg (public

domain), en:Image:Brillanten.jpg (GFDL), and parts of Image:

Eight_Allotropes_of_Carbon.png (GFDL).


http://schools-wikipedia.org/images/497/49753.jpg.htm


a sheet of paper. Think of interlinking carbons all with a tetrahedral carbon center. Graphite,

diamond, and ―Bucky‖ balls are all allotropes. They are all pure carbon, but the carbon atoms

are bonded together in different ways for these 3 forms.

Watch the video ―Crystal Structures: Graphite and

Diamond‖ on the class website.

Describe hybridization (sp, sp2, and sp

3) of simple molecules.

This was described in the previous unit.

The headline version of hybridization is:

1. Any atom with a total of 4 single bonds plus lone pairs

is likely to have an sp3 hybridization

2. Any atom with one double bond and a total of 2 single

bonds plus lone pairs is likely to have an sp2

hybridization

3. Any atom with a triple bond or 2 double bonds is likely

to have an sp hybridization

Understand how the capacity to form four covalent bonds

results in several bonding possibilities for carbon, including

Single, double, and triple bonds

Ring structures

Covalent network

Explanation:

Single, double, and triple bonds were explained in the

previous unit of instruction.

Ring structures result from a carbon atom’s ability to

bond to it other carbon atoms in tetrahedral and trigonal

planar arrangements with hydrogen atoms forming the

terminating part of the structure. Graphite’s structural

unit is illustrated in figure 4.9.

Graphite is an infinite number of the individual units

bonded together in ―sheets.‖

These sheets (illustrated to right)

have delocalized electrons formed from pi (π)

bonds that keep the ―sheets‖ from sticking to

each other. The result is a slippery solid.

If we were to look at any individual carbon

atoms in the benzene ring that carbon forms in

graphite, we would see that each carbon has 2

single bonds and a double bond on it. That

means that each carbon has an sp2 hybridization

and it forms a trigonal planar bonding

arrangement around each carbon atom. That

trigonal planar bonding structure is true for every carbon in the graphite

sheet.

Figure 4.12. Three dimensional

illustration of a

Buckminsterfullerene, also

called a ―Bucky Ball.‖ This is a

60-carbon sphere, C60.

Figure 4.13. A 2 dimensional

structural drawing of a small section of

graphite.

CH

CHCH

CH

CHCH

Figure 4.14. A two

dimensional structural

drawing of benzene

which is the structural

unit of graphite.

CH2

CH2

CH

2

CH2

CH2

C

H2

CH2

CH

2

CH2

CH2

C

H2

CH2

C

H2

CH2

CH2 CH

2

CH

2

CH

2


This structural unit for graphite is called a ―benzene ring,‖ but the chemical benzene is made

entirely of these individual units that are not bonded to other such units. Note that a benzene ring

has 6 carbons forming a ―ring‖ and each carbon has at least one double bond. All of these carbons,

therefore, have an sp2 hybridization. These benzene rings can link up with an infinite number of

very complex ring structures to make an infinite number of organic compounds.

Example from your life: Pull out the drug insert from a prescription and look for the structural

formula for that drug. Chances are, the structure of that drug includes many of the complex ring

structures bonded together.

To the left are some examples of other ring structures in which carbon only has single bonds: Each

of the carbon atoms in these structures have a tetrahedral bonding arrangement and an sp3

hybridization. Notice that these structural formulas have line bonds, as in a Lewis formula, but

they also contain >CH2 groups, much like a typical molecular formula. This combination of Lewis

formula and molecular formulas is one of the condensed structural formulas that you will learn

about shortly. For now, it is enough that you know that each carbon in a >CH2 group has 4 single

bonds and that makes it an sp3 hybridization.

Check out the web site:

http://www.avogadro.co.uk/structure/chemstruc/network/g-molecular.htm to explore how these

complex structures can work.

You should watch the video that shows these allotropes at

http://www.youtube.com/watch?v=vYkyUqUa6vU&feature=related . Or you can see this video on

the Honors Chemistry web page at: http://tedderchemistry.com/honors.html .

Carbon’s ability to form a covalent network structure (a diamond) results from its ability to form

tetrahedral bonds to itself infinitely (see figure 4.6).

Practice Problems:

36. Each carbon in a bucky ball (Buckminsterfullerene ) or other fullerenes is singly bonded to 2 other carbon atoms and double bonded to a 3

rd carbon atom. What is the hybridization on these carbons atoms in a

fullerene?

37. In the simplest ring structures of carbon, each carbon atom is singly bonded to 2 other carbon atoms and is also singly bonded to 2 hydrogen atoms. What is the hybridization on these carbons atoms in such a ring

structure?



59) What does the acronym IUPAC mean?

60) What is a straight chain alkane?

61) How are straight chain alkanes named?

Indicator C-3.5 – Naming and writing structural

formulas for simple hydrocarbons

According to Indicator C-3.5 in the South Carolina Science

Standards students should be able to illustrate the structural

formulas of and name simple hydrocarbons (including alkanes and their isomers and benzene rings).


Understand International Union of Pure and Applied Chemistry (IUPAC) organic nomenclature

IUPAC nomenclature is the system used to name most organic compounds. The rules would fill a

large book and students will only be expected to understand and be able to use a few of these.

When the term ―systematic name‖ is used in this text, it means that it is named according to the

IUPAC naming system.

Name and write the formula for alkanes (up to 10-carbon), their isomers and benzene rings

Explanation: Alkanes are chains of carbon atoms with all single bonds. The naming system is

based on the number of carbons in the chain. There is a prefix for each number of carbons

followed by the ―-ane‖ ending for alkanes. In the case of alkanes, these ―-ane‖ endings are the

―root‖ name. The listed in table 4.1.

Table 4.2. IUPAC prefixes and suffixes for naming straight chain alkanes and the formulas indicating carbon

chain length of alkanes. Carbons in chain prefix

Root name Condensed structural formula Skeletal structural formula

1 meth- -ane CH4 none

2 eth- -ane CH3CH3 3 prop- -ane CH3CH2CH3

4 but- -ane CH3CH2CH2CH3 5 pent- -ane CH3CH2CH2CH2CH3

6 hexa- -ane CH3CH2CH2CH2CH2CH3

7 hepta- -ane CH3CH2CH2CH2CH2CH2CH3

8 octa- -ane CH3CH2CH2CH2CH2CH2CH2CH3

9 nona- -ane CH3CH2CH2CH2CH2CH2CH2CH2CH3

10 deca- -ane CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3

Note 1: The condensed structural and skeletal structural formula examples shown above are ―straight chain‖ hydrocarbons.

This means that there are no side chains are functional groups which you will learn about later in this unit.

Note 2: The prefixes meth-, eth-, prop-, and but- come from common names for chemicals that existed long before the

structures of compounds were known. Meth-, for example, comes from methanol which comes from the Greek words

methyl, which means wine, and hyle, which means wood. Methanol can be prepared b y heating wood in the absence or air

and is often called wood alcohol. Methanol’s formula is CH3OH. Notice that it only has one methyl group, which is a

carbon single bonded only to other carbon atoms or to hydrogen atoms.



Draw the structural formulas for alkanes up to a

10-carbon chain. See table 4.1 above.

If you write the Lewis formula for propane, you

get the illustration in figure 4.10:

If you group all the hydrogen atoms with the

carbons to which they are bonded, you get a

condensed structural formula, as illustrated in

figure 4.11.

An even more condensed structural formula for propane is CH3CH2CH3.

The skeletal formula eliminates the carbon and the hydrogen symbols (see figure

4.12). The end of each line or a corner in a skeletal formula indicates the location

of a carbon atom and the chemistry student then must infer where the hydrogen

atoms are located.

Students are expected to draw complete structural formulas, condensed structural

formulas, and skeletal structural formulas for organic compounds.

Practice Problems:

38. What is the name of CH3CH2CH2CH3?

39. Write the complete structural formula for the substance above.

40. Write the skeletal formula for the substance above.

41. Write the complete structural formula for hexane.

42. Write the condensed structural formula for hexane.

43. Write the skeletal formula for hexane.

44. What is the organic chemical name of this skeletal formula?

45. How many carbon atoms are there in the skeletal formula in the last question?

46. How many hydrogen atoms are there in the skeletal formula in the last question?

47. Write the skeletal formula for propane.

48. Write the condensed structural formula for propane.

49. How many carbon atoms are there in ethane?

50. What are the 10 prefixes used to indicate the length of the carbon chain in a straight chain alkane what does each prefix mean?

51. What is the suffix used to indicate that a carbon chain has all single bonds between carbon atoms?

Figure 4.15. A Lewis

formula or complete

structural formula for

propane.

Figure 4.16. A

condensed structural

formula for propane.

Figure 4.17. AA skeletal

formula for

propane.


Flow Chart for Naming Non-organic Compounds from Formulas Begin Here

This is an ionic

compound.

Determine the

correct oxidation

number. Then

name the 1st element

followed by the

oxidation number

written in

parentheses in

Roman numeral

form. Example: Fe3+ is

iron(III)

Name the 2nd

element as it

appears on the periodic table but

change the ending of the 2nd

element name to –ide.

All polyatomic ions are named just they appear on the

common ions chart.

This is an acid

Name the polyatomic ion as

it appears on the common

ions chart and then change

the –ate ending to -ic. Note: If the polyatomic ion is sulfate then change ending to –uric (sulfuric). If the polyatomic ion is phosphate then

change ending to –oric (phosphoric).

Name the polyatomic ion as

it appears on the common

ions chart and then change

the –ite ending to -ous. Note: If the polyatomic ion is sulfite then change the ending to –urous (sulfurous). If the polyatomic ion is phosphite then change ending to –orous (phosphoric).

Write the prefix hydro-;

name the 2nd

element, name

the polyatomic ion as it

appears on the common ions

chart, and then change the

ending to an -ic ending. Note: If the polyatomic ion is sulfate then change ending to –uric (sulfuric). If the polyatomic ion is phosphate then change ending to –oric (phosphoric).

Roman numerals: 1 = I 2 = II

3= III 4 = IV

5 = V 6 = VI

7 = VII 8 = VIII

9 = IX 10=X

Name the

metal just as

the name

appears on the

periodic table.

For ALL acids put the word ―acid‖ on the end of

the name.

Does the formula contain a polyatomic ion?

YES NO

Does the polyatomic ion contain an

-ate ending? YES NO

Are both elements non-metals? YES NO

This compound is covalent. Name the

first element using the proper prefix (but

never use the mono- prefix). The, name

the 2nd

element using the proper prefix

(but ONLY use mono- prefix for carbon

monoxide).

1 = mono- (this is only used for carbon monoxide)

3 = tri- 6 = hexa-

9 = nona-

2 = di-

5 = penta-

8 = octa-

4 = tetra-

7 = hepta-

10 = deca-

Does the formula contain a polyatomic ion?

YES NO

Does the formula begin with a metal that has

more than 1 oxidation number? NO

YES

Does the formula

contain a metal with

only one kind of

oxidation number?

YES NO

Does this formula have a hydrogen symbol (H) on the front? NO YES


Flow Chart for Writing Non-organic Chemical Formulas from IUPAC Chemical Names Begin Here

The first part of the

name is the cation (the positive ion).

The last part of the

name is the anion (the negative ion).

This is

multivalent.

The Roman

numeral is the

charge on the

cation. Ex: iron(II) is Fe2+.

This is monovalent.

Determine the

charge on the cation

from the periodic

table with oxidation

numbers. Ex: calcium is Ca2+.

After balancing the charges, write the balanced formula. Remember: You ONLY use parentheses when you have more than one

polyatomic ion.

Example: 2Na+ + SO4

2– Na2SO4

Balanced charges Balanced formula

First name:

The prefix in front

of the 1st element

name tells you

how many atoms

of that element

there are in the

formula.

Write those

numbers as

subscripts

following the

element symbol.

Second name:

The 2nd

element name will

have an –ide ending.

You have to determine the

identify of that element

without the –ide.

The prefix in front of the 2nd

element name tells you how

many atoms of that element

there are in the formula.

Write that number as a

subscript following the

element symbol.

Write the formula.

Ex: dinitrogen pentaoxide N2O5

Non-metal formula

prefixes/subscripts: 1 = mono- (this is only

used for carbon

monoxide) 2 = di- 3 = tri-

4 = tetra- 5 = penta-

6 = hexa- 7 = hepta-

8 = octa- 9 = nona-

10 = deca-

Roman

numerals/charges:

1+ = I 2+ = II

3+ = III 4+ = IV 5

+ = V 6

+ = VI

7+ = VII 8+ = VIII

9+ = IX 10+ =X

Does the name have an –ic or –ous ending? NO YES

This is an acid. Go to the next page

entitled: ―Flow Chart for Writing Non-

organic Acid Formulas from IUPAC

Chemical Names.‖ This is a covalent compound. Atoms share electrons. This compound is ionic.

Components transfer electrons.

Is there a metal or polyatomic ion

present in the formula? YES NO

Balance the charges! Remember LCM – lowest common multiple.

Look up the

polyatomic

ion on the

Common

Ions chart. Make sure that

you include

the charge on

the ion!

The atom that forms

this mono-atomic

anion is a non-metal.

Determine the

charge on the anion

using the periodic

table with oxidation

numbers. Note: The charge must

be negative!

Is the cation a metal or a polyatomic ion?

A METAL A POLYATOMIC ION

Is the cation name followed by a Roman

numeral? Ex: iron(II). YES NO

Does the last part of the name of this

compound end in -ide? NO YES

Notice: There are no

parentheses in the

formula to the left for

sodium sulfate: Na2SO4.

This is because sodium

sulfate only has one

polyatomic ion.

On the other hand,

ammonium sulfide has 2

ammonium ions and this

formula REQUIRES

parentheses: (NH4)S.


Flow Chart for Writing Non-organic Acid Formulas from IUPAC Chemical Names Begin Here

This is an acid.

This is a binary acid.

Find the non-metal from which the acid

was derived.

This is NOT an acid.

Go to the previous page entitled:

―Flow Chart for Writing Non-

organic Chemical Formulas from

IUPAC Chemical Names.‖

Does the name have an –ic or –ous ending? NO YES

Does the name have a hydro-

prefix? YES NO

Example: hydrosellenic acid The –sellen– root word for this acid name is

taken from a non-metal element. If you examine the element names on the periodic

table you will find that it is spelled just like the

beginning of the element sellenium. Therefore, sellenium is the non-metal from which the acid

was derived.

Determine the non-metal’s negative charge

(using the periodic table with oxidation numbers).

BE CAREFUL! Be sure to include the charge on the polyatomic

ion.

Do NOT confuse the subscript with the charge.

Example: The chlorite ion has a 1 charge. Its

formula is ClO3 . The subscript 3 means that there

are 3 oxygen atoms and NOT a 3 charge.

Example: Sellenium has 4 oxidation

numbers (6+, 4+, and 2–) but the only

negative charge is 2–. Therefore, 2– is the

charge that you MUST use.

Balance the formula with H+.

Example: 2H+ + Se

2– H2Se

The coefficient 2 on the H+ becomes a subscript 2 in

the chemical formula.

Find the polyatomic oxyanion from which

the acid was derived. This polyatomic ion

will have an –ite ending. Polyatomic ion formulas are found on the common

ions chart.

This is an oxyacid. Does the name end

in –ic? NO YES

Balance the polyatomic ion with H+.

Example: 3H+ + PO3

3 H3PO3

Notice: The coefficient 3 on the H+ becomes a

subscript 3 in the chemical formula.

Example: phosphorous acid

The –phosphor– root word for this acid name

is taken from the phosphite ion. The formula

for the phosphite ion is PO33

. BE CAREFUL!

Be sure to include the charge on the polyatomic

ion. Do NOT confuse the subscript with the charge.

Example: The chlorite ion has a 1 charge. Its

formula is ClO3 . The subscript 3 means that there

are 3 oxygen atoms and NOT a 3 charge.

Find the polyatomic oxyanion from

which the acid was derived. This

polyatomic ion will have an –ate ending. Polyatomic ion formulas are found on the

common ions chart.

Balance the polyatomic ion with H+.

Example: 3H+ + PO4

3 H3PO4

Notice: The coefficient 3 on the H+ becomes a

subscript 3 in the chemical formula.

Example: phosphoric acid

The –phosphor– root word for this acid

name is taken from the phosphate ion. The

formula for the phosphate ion is PO43

.

Non-metal formula

prefixes:

1 = mono- (this is only

used for carbon

monoxide) 2 = di- 3 = tri-

4 = tetra- 5 = penta-

6 = hexa- 7 = hepta-

8 = octa- 9 = nona- 10 = deca-

1 = I 2 = II

3= III 4 = IV

5 = V 6 = VI 7 = VII 8 = VIII

9 = IX 10=X

Ro

ma

n n

um

erals

:

Acid Ion -ic -ate

-ous -ite

-ide -ic

Aci

d &

.ion

suff

ixes

:

unit 4 notes chemistry i cp page 1 -...

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