23: Naming Organic Compounds

Naming Organic Compounds/IUPAC Nomenclature:

Alkyl Group: a functional group attached to another hydrocarbon that contains only carbon and hydrogen atoms. Here, we'll only deal with the first two, methyl and ethyl.

Methyl: –CH₃



The blue shading is the main chain of carbons and the pink shading is methyl group. On a skeletal structure, it looks like one extra line protruding from the main chain, to show another carbon to carbon link with the three hydrogens implied on its end. Here is an example of a main chain with the structure of hexane with a methyl group attached to it:



Ethyl: –C₂H₅



Again, the blue is the main chain and the pink is the ethyl group. On a skeletal structure, it looks like an extra line with a kink in it protruding from the main chain. Here is an example of a main chain with the structure of heptane with an ethyl group attached to it:





(Rules for naming taken off of the handout from class)

Naming Alkanes:

1. Find parent chain: longest continuous chain of carbons. It will not always be straight across. In some of the examples, what appears to be an ethyl group is actually the end of the parent chain. The parent chain is named according to how many carbons it has: butane if it has four carbons, pentane if it has five carbons, etc. This will go at the very end of the name when we're done.

2. Find and identify the alkyl groups (methyl or ethyl as far as this class goes).

3. Number the carbons on the chain starting from the side closest to an alkyl group. Each alkyl group will get a prefix of the number associated with the carbon it is bonded to, even if there is more than one of that kind of alkyl group. But instead of writing out “methyl” two times if the alkane has two methyl groups, it is numbered with the appropriate Greek prefix (in this example, it would be di- for two). The numbers associated with the positions of the carbons are written in front, with commas separating them from each other and a dash separating them from words.

4. When both of the alkyl groups we're dealing with are present, number each group and list them in alphabetical order in the name.

Examples:








Naming Cycloalkanes:

Cycloalkanes are saturated hydrocarbons with rings. Their skeletal structures are geometric shapes where the corners represent carbon atoms assumed to be bonded with hydrogens.



1. As far as this class goes, there will either be an ethyl group or a methyl group (or multiples) attached to the ring. If there is only one, there is no need to give it a prefix.

2. When there are two alkyl groups, alphabetize them. The one that comes first alphabetically is given the first spot, and the carbons are numbered in the direction that will give the smallest number for the second substituent (so number towards the shortest path to the second alkyl group).

3. If there are more than two substituents, number in a way that yields the smallest numbers possible in the finished name. If there are identical sets of numbers, use alphabetical priority.

Examples:




Naming Haloalkanes:

Also called alkylhalides/halogenated alkanes. They are alkanes with one or more halogens attached (F, Cl, Br, I).

1.
F = fluoro
Cl = chloro
Br = bromo
I = iodo

2. Halogen and alkyl substituents are considered of equal rank, go by alphabetical priority.

Examples:




Naming Alkenes and Cycloalkenes:

Alkenes and cycloalkenes have one or more carbon double bonds.

1. Replace the suffix -ane with -ene.

2. The parent chain will be the longest continuous chain of carbon atoms that contains both carbons of the double bond.

3. Number the parent chain from the end closest to the double bond.

4. Two numbers will be associated with the double bond, assign it the smaller one. This number will appear as a prefix to the parent chain name, with a dash separating numbers and words.

5. With cycloalkenes, number in the direction that gives the lowest numbers for substituents.

Examples:




Naming Alkynes:

Alkynes have carbon triple bonds. The rules are exactly the same as they are for the alkenes, except the new suffix is -yne.

Example:




Naming Alcohols:

Alcohols have hydroxyl groups: –OH

1. Name the longest parent chain that includes the hydroxyl group. The suffix for alcohols is -anol.

2. Number starting at the end closest to the hydroxyl group. The number associated with the hydroxyl group will appear as a prefix to the parent chain name, with a dash separating numbers and words.

3. Name and number the other substituents the same as before.

Examples:

23: Introduction to Organic Chemistry

Organic Chemistry: The Chemistry of Hydrocarbons and Hydrocarbon Derivatives
Carbon can bond in four different ways. All of these bonds are very strong, especially if the bond is to another carbon.



Definitions:
Hydrocarbon: compound containing only carbon and hydrogen.
Hydrocarbon Derivatives: compounds containing carbon and hydrogen and one or more different elements.
Saturated Hydrocarbon: hydrocarbons that contain only single bonds between the carbon atoms.
Unsaturated Hydrocarbon: hydrocarbons that contain one or more multiple bonds (double or triple bonds) between the carbon atoms.
Cyclic Hydrocarbon: a hydrocarbon in which a chain of carbon atoms has formed a ring.
Acyclic Hydrocarbon: a hydrocarbon that does not contain a ring of carbon atoms.

Alkanes:


These are the molecular and structural formulas for the first six. Following the same pattern, after hexane comes heptane, octane, nonane, and decane.

It is important to remember that carbon atoms form four bonds. Do not give your carbon atom more than four things to carry, because he only has four arms.


Condensed Structural Formula: formulas where the bonds around each carbon atom in the compound are not explicitly written.

Example:
Butane has a chain of four carbons, two of which attach to three hydrogens each, and two of which attach to two hydrogens each.

The condensed structural formula for butane is: CH₃CH₂CH₂CH₃ or CH₃(CH₂)₂CH₃

Think of the CH₃ on each end like the front and back of a train, and the CH₂ in the middle keeps getting added on to as you move to the next molecule in the series, just as you would add extra cars to a train.


Constitutional Isomerism: compounds with the same molecular formula but different structural formulas.



Methane, ethane, and propane all only have one isomer. They only have one structural formula each. Butane has two isomers. Pentane has three.


Remember a straight-chain can bend, but if the bonds on the carbon atoms haven't changed, it's not a new isomer.


Line-Angle Structural Formula (Skeletal Structure):


1. Each intersection, endpoint, and kink is a carbon.
2. The carbon to carbon bonds are shown, as well as the bonds between carbon and anything that's not a hydrogen.
3. Every remaining bond on carbon is assumed to be a hydrogen.
4. When counting the carbons in chain, the very beginning of the chain is where you count the first carbon.

22.7: Crystal Field Theory

Crystal Field Theory:
Remember what the 3d orbitals look like:


Simplify the ligands into negatively charged or slightly negatively charged point particles. A complex ion with six ligands will have them arranged octahedrally around the metal atom (because they will repel each other). The arrangement of these negative charges will lie on the x, y, and z axes, as represented by the red spheres below:


Notice that the arms of the d_z² and d_{x² – y²} orbitals also lie along the axes where the octahedrally arranged ligands (negative point particles) lie (the other orbitals lie between the axes). Those electrons will have a stronger repulsive effect on the ligands, creating a situation where the five 3d orbitals are no longer of the same energy.

These two orbitals that lie on the axes will have higher energy, and the other three will have lower energy.


Where Δ = crystal field splitting = the difference in energy between the two sets of d orbitals on a central metal ion that arises from the interaction of the orbitals with the electric field of the ligands.


High and Low Spin Complexes:

Pairing Energy: the energy required to put two electrons into the same orbital, denoted P. When an orbital already has an electron in it, there needs to be an input of energy to overcome the repulsion these two electrons will have for each other, this is the pairing energy.

Imagine the electron configuration of a metal ion in a complex, specifically the 3d orbitals. There are five to fill. Three have lower energy than the other two, which have higher energy. So the three low energy orbitals get filled first. But say there are four electrons in the configuration, not three, how do we know whether to start pairing the electrons in the low energy orbitals or to go ahead and start filling the empty, higher energy orbitals?


It depends on P and Δ. If P is greater than Δ, this means it requires more energy to pair an unpaired electron in a lower energy orbital than it requires to boost that electron up to the higher energy orbital. So, if P is greater than Δ, that fourth electron goes in the first empty high energy orbital, and it is called a high spin complex:



If Δ is greater than P, this means it requires more energy to boost that electron up to the higher energy orbital than it does to pair an unpaired electron in a lower energy orbital. So, if Δ is greater than P, that fourth electron pairs with the first unpaired electron in the low energy orbital, and it is called a low spin complex:


Here is a complete filling of the 3d orbitals, comparing high spin and low spin:


Paramagnetic substances have at least one unpaired electron in their orbitals. Most of the configurations in the diagram above show paramagnetism. Diamagnetic substances have no unpaired electrons in their orbitals. Diamagnetism is shown in the diagram above at six electrons at low spin, and at ten electrons at both high and low spin.

Spectrochemical Series: an arrangement of ligands according to the relative magnitudes of crystal field splittings they induce in the d orbitals of a metal ion.

As Δ increases, ligands bond more strongly (found on page 957).

Example:
Which one is high spin, [CrBr₆]^3– or [Cr(CN)₆]^3–?

From the spectrochemical series, we see that Br has a relatively small Δ, so it is high spin.
CN has a relatively large Δ, so it is low spin.


Visible Spectra of Transition-Metal Complexes:

Many transition-metal substances are colored, and the color is related to Δ through the wavelength, λ, of light absorbed:

Δ = hυ = hc / λ
or
λ = hc / Δ

where:
Δ = crystal field splitting
λ = wavelength
υ = frequency
h = Planck's Constant = 6.626068 x 10^-34 J · s
c = the speed of light = 3.00 x 10⁸ m / s
Note: With these problems, wavelength might be given in nanometers, but the rest of the units are in meters. To convert: 1 nm = 1 x 10^–9 m

Example:
For a certain substance, if λ = 555 nm, what is Δ?

Convert from nanometers to meters: 555 nm = 555 x 10^–9 m = λ
h = 6.626068 x 10^-34 J · s
c = 3.00 x 10⁸ m / s

Δ = hc / λ
Δ = [(6.626068 x 10^-34 J · s)(3.00 x 10⁸ m / s)] / (555 x 10^–9 m)
Δ = 3.58 x 10^-19J

Example:
For a certain substance, if Δ = 2.00 x 10^–19, what is λ?

λ = hc / Δ
λ = [(6.626068 x 10^-34 J · s)(3.00 x 10⁸ m / s)] / (2.00 x 10^–19 J)
λ = 9.95 x 10^–7 m = 995 nm

22.6: Valence Bond Theory of Complexes

Valence Bond Theory of Complexes:

Remember from Chapter 10 that a covalent bond is usually formed by the overlap of the a bonding orbital from each atom, each containing one electron. The subsequent bond then has two electrons.

Remember also that the bonds in complexes are coordinate covalent bonds, where both electrons are donated from the ligand to the metal ion. With the coordinate covalent bond in a complex, a ligand orbital with two electrons overlaps an empty orbital on the metal.



This bonding requires hybrid orbitals.

Example:

What are the hybridized orbitals on Co in [Co(H₂O)₆]³⁺?

First we need to find the configuration of the metal ion. Start by writing the configuration of the cobalt atom. Here's a helpful table to remember how to do it:



Co: [Ar]4s²3d⁷

If our ion has a charge of +3, and H₂O is a neutral molecule, we know that the charge on the cobalt ion is +3. To reflect this in the configuration, remove 3 electrons. Remember the outer s electrons are removed first:

Co³⁺: [Ar] 3d⁶

Now draw the orbital diagram of the ion. Remember to follow Hund's Rule, where each electron is placed in a separate orbital with the same spin before pairing it with an electron of opposite spin:



In the complex ion [Co(H₂O)₆]³⁺ each water molecule wants to donate one pair of electrons to the metal. In order to do this, it needs an empty orbital. Since there are six H₂O, it needs six empty orbitals. Starting with the first empty orbital, work forward until you get six, and these are the hybridized orbitals:



It is called sp³d² because one orbital comes from s, three come from p, and two come from d.

Example:

What are the hybridized orbitals on Mn in [Mn(NH₃)₆]⁴⁺?

First we need to find the configuration of the metal ion. Start by writing the configuration of the manganese atom.

Mn: [Ar] 4s²3d⁵

If our ion has a charge of +4, and NH₃ is a neutral molecule, we know that the charge on the manganese ion is +4. To reflect this in the configuration, remove 4 electrons. Remember the outer s electrons are removed first:

Mn⁴⁺: [Ar] 3d³

Now draw the orbital diagram of the ion, following Hund's Rule:



In the complex ion [Mn(NH₃)₆]⁴⁺ each NH₃ molecule wants to donate one pair of electrons to the metal. In order to do this, it needs an empty orbital. Since there are six, it needs six empty orbitals:



It is called d²sp³ because two orbitals come from d, one comes from s, and three come from p.

22.5: Structure and Isomerism in Coordination Compounds

Structure and Isomerism in Coordination Compounds:

Constitutional Isomers: isomers that differ in how the atoms are joined together, specifically, in the order that they are bonded to each other.

Stereoisomers: isomers in which the atoms are bonded to each other in the same order but differ in the precise arrangement of the atoms in space.


Constitutional Isomers:

Example: Ion-Exchange Isomer

[Co(NH₃)₅(SO₄)]Br
[Co(NH₃)₅Br]SO₄

In the first isomer, SO₄ is attached to the Cobalt and is part of the complex ion (the cation), with Br as the anion. In the second isomer, Br is attached to the cobalt as part of the complex and SO₄ is acting as the anion.

Example: Coordination Isomer

[Cu(NH₃)₄][PtCl₄]
[Pt(NH₃)₄][CuCl₄]

Here, both the cation and anion are complex ions. In the first isomer, NH₃ is attached to the copper and the Cl^– are attached to the platinum. In the second isomer, they have swapped.

Example: Linkage Isomer

[FeCl₅(NO₂)]^3–
[FeCl₅(ONO)]^3–

Here, the difference is in how the ligand bonds to the metal. In the first isomer, the ligand bonds to the metal through an electron pair on the nitrogen. In the second isomer, the ligand bonds to the metal through an electron pair on one of the oxygen atoms. It's easier to see it:




Stereoisomerism:

Geometric Isomers: isomers in which the atoms are joined to one another in the same way but differ because some atoms occupy different relative positions in space.



In the cis arrangement, two like ligands are placed next to each other. In the trans arrangement, two like ligands are placed across from each other.




(helpful source)

(info on mer/fac isomers)


Chiral vs. Achiral:



Achiral Molecules: these isomers are superimposable mirror images of one another. A molecule and it's mirror image can completely overlap.

Chiral Molecules: If a molecule and a molecule that appears as its mirror image cannot completely overlap, they are chiral molecules. These isomers are non-superimposable mirror images of one another. Also called optical isomers, optically active, or enantiomers.

Enantiomers have identical properties, but are differentiated by their effect on plane-polarized light.



Dextrorotatory: a compound whose solution rotates the plane of polarized light to the right (when looking toward the source of light).

Levorotatory: a compound whose solution rotates the plane of polarized light to the left (when looking toward the source of light).

Racemic Mixture: a mixture of equal amounts of optical isomers. Because the two isomers rotate the plane of polarized light by the same angle in opposite directions, they cancel each other out and have no net effect.

22.4: Naming Coordination Compounds

Naming Coordination Compounds:

Ligands That Are Part of a Complex: (never (rarely?) positively charged)

Ligands With No Charge:
H₂O = aqua
NH₃ = ammine
CO = carbonyl
en = ethylenediamine

Ligands With a Charge of –1:
CN^– = cyano
F^– = fluoro
Br^– = bromo
Cl^– = chloro
I^– = iodo
OH^– = hyrdoxo


Ligands With a Charge of –2:
O^2– = oxo
C₂O₄^2– = oxalato (only other bidentate ligand)
SO₄^2– = sulfato

With the exceptions of en and oxalato, all of these ligands are monodentate (one “bite”).

Rules For Naming:

1. The name of the cation comes first, the anion comes second. The name of the complex consists of two parts written together as one word. Ligands are named first and the metal atom is second.

2. Name and number the ligands in alphabetical order.
Prefixes for Numbering:
1 = mono
2 = di
3 = tri
4 = tetra
5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona/ennea
10 = deca

3. Metal name stays intact if it is a cation. Add -ate to the end otherwise.

Name/Anion Name:
Copper/Cuprate
Gold/Aurate
Iron/Ferrate
Lead/Plumbate
Silver/Argentate
Tin/Stannate

4. Add “ion” if it is one.

Example:

What is the name of the following complex ion?

[Co(NH₃)₆]³⁺

First, name and number the ligands in alphabetical order. There is only one ligand, the NH₃, so no need to worry about alphabetizing here. This ligand is called ammine.

There are six of them, so it is hexaammine.

We know that the NH₃ molecule has no charge, and we see that the ion has a +3 overall charge, which means this must come from the cobalt, so we know the name of the metal is Cobalt(III). And since it is the cation, we can use this name instead of changing it.

The name of our complex ion then is hexaamminecobalt(III) ion.

What is the Coordination Number of this ion?

NH₃ is monodentate, and there are six of them. So C.N. = 6


Example:

[Fe(H₂O)₂(CN)₄]^2–

Name and number the ligands in alphabetical order:

(H₂O)₂ = diaqua
(CN)₄ = tetracyano

Name the metal:

H2O has no charge. CN has a –1 charge, and there are 4 of them. And the total charge on the complex ion is –2. So we have:

Charge on Fe + 0 + (–4) = –2
Charge on Fe = +2

So we know we have Iron(II). And since the complex ion is an anion, it's called ferrate(II).

The name of the complex ion is diaquatetracyanoferrate(II) ion.

All six ligands are monodentate, so C.N. = 6.


Example:

Na₃[Ti(O)₂(OH)F₂(SO₄)]

The name of the cation, Sodium, comes first.

Name and number the ligands in alphabetical order:

F₂ = difluoro
OH = hydroxo (the mono is unnecessary)
O₂ = dioxo
SO₄ = sulfato

Name the metal:

O has a charge of –2 and there are two of them. OH has a charge of –1. F has a charge of –1 and there are two of them. SO4 has a charge of –2. And since Na has a charge of +1, and it takes 3 of them to neutralize the anion, we know the charge on the anion must be –3. So:

Charge on Ti + (–2 x 2) + (–1) + (–1 x 2) + (–2) = –3
Charge on Ti = +6

So it is Titanium(VI). Since it is part of the anion, it is called titinate(VI).

The name of the complex ion is Sodium difluorohydroxodioxosulfatotitinate(VI).

It has six monodentate ligands, so the C.N. = 6.


Example:

[Os(C₂O₄)₂Cl₂]Cl₂

Notice that Cl₂ appears twice. The first time it is part of the complex. The second time, it is the anion. And just like with other ionic compounds, the Cl₂ denotes two Cl^– and not Cl₂ gas. And as the anion, chloride will be named after the complex ion.

Name and number the ligands in alphabetical order:

Cl₂ = dichloro
(C₂O₄)₂ = dioxalato

Name the metal:

C2O4 has a charge of –2 and there are two of them. Cl has a charge of –1 and there are two of them. And we know that it takes two of the Cl^– anion to neutralize the charge on the complex ion, which must be +2. So:

Charge on Osmium + (–2 x 2) + (–1 x 2) = +2
Charge on Osmium = +8

So it is Osmium(VIII). As part of the cation, it gets to keep this name.

So the name of the complex is dichlorodioxalatoosmium(VIII) chloride.


Example:

[Cd(en)₃][PtCl₆]

Here we have a complex ion as both the anion and the cation.

Special case with en:

When the name of the ligand also has a number prefix, as in ethylenediamine, there is a new set of prefixes that needs to be used to name it:
2 en = bis-
3 en = tris-
4 en = tetrakis-

Remember that Cadmium has only one oxidation state, and can only have a charge of +2, and en has no charge. So the charge on the cation is +2, which means the charge on the anion is –2. Cl has a charge of –1 and there are six of them. So:

Charge on Platinum + (–1 x 6) = –2
Charge on Platinum = +4

So, following the rules, the name of our complex compound is trisethylenediaminecadmium(II) hexachloroplatinate(IV).

Now try to write the notation for the complex ion from the name:

Example:

dicarbonyltrioxalatoaurate(IV) ion

When writing it this way, the metal comes first. Aurate(IV) = Au

dicarbonyl = (CO)₂
trioxalato = (C₂O₄)₃

charge on Au(IV) = +4
charge on (CO)₂ = 0 (neutral molecule)
charge on (C2O4)3 = –6
charge on ion = –2

[Au(CO)₂(C₂O₄)₃]^2–

22.3: Formation and Structure of Complexes

Formation and Structure of Complexes:

Transition-metal atoms often function as a Lewis acid in a chemical reaction. Remember that a Lewis acid accepts electron pairs from molecules or ions.

Example:

Complex ions are not held together by ionic bonds, but coordinate covalent bonds. Remember that coordinate covalent bonds are bonds between atoms where both electrons involved in the bond come from only one of the atoms.

In water, the Fe²⁺ ion ultimately bonds to six H₂O, forming the [Fe(H₂O)₆]²⁺ ion. The final charge on the ion can be found by adding up all the charges of its components.

Example:

Fe = +2
H₂O = 0

2 + (0 x 6) = 2

Charge on the ion = +2

Definitions:

Complex Ion: A metal ion with Lewis bases attached to it through coordinate covalent bonds.

Complex/Coordination Compound: a compound consisting either of complex ions and other ions of opposite charge or a neutral complex species (such as the anti-cancer drug in the chapter opening). Has a cation and an anion.

Example:

[FeOH₂]Cl₂ ← means 2 Cl^–, not Cl₂ gas. Two negatively charged chlorine ions neutralize the +2 charge on the complex ion. Everything in the brackets in front of the Cl₂ is the cation.

Ligands: the Lewis bases attached to the metal atom in a complex. Electron pair donors. Can be neutral (such as H₂O or NH₃) or anions (such as Cl^–).

Coordination Number: the total number of bonds the metal atom forms with ligands. The number of lone pairs that are attached to the metal via the coordinate covalent bond.

Example:

In [Fe(H₂O)₆]²⁺, the metal atom, Fe, bonds to each oxygen atom in six water molecules.

Coordination Number (C.N.) = 6

Example:

[Ag(NH₃)₂]⁺ → [H₃N꞉Ag꞉NH₃]⁺

One coordinate covalent bond between the Ag and each N of the two NH₃ molecules.
C.N. = 2

Polydentate Ligands: ligands that can bind with two or more atoms to a metal atom. Think of the word “bite” when you hear dentate. Monodentate means it forms one bond with its atoms (one bite), bidentate means it forms two bonds with its atoms (two bites), etc.

22.1: Periodic Trends in the Transition Elements

Transition Elements:



Transition Elements: those metallic elements that have an incompletely filled d subshell or easily give rise to common ions that have incompletely filled d subshells. Highlighted in the table above. They are the basis for Inorganic Chemistry (every other element except carbon).

Periodic Trends of the Transition Elements:

1. They are all metals.

2. Mostly have high melting points and high boiling points and are hard solids. This is due to metallic bonding.



The negatively charged electrons form an “electron sea” around the positively charged nuclei of the metal atoms and are shared as they move about the sea.

3. They have more than one oxidation state (exceptions: Cd²⁺, Ag⁺, Zn²⁺).

Example:

Fe²⁺, Fe³⁺; Mn²⁺, Mn³⁺, Mn⁴⁺, Mn⁷⁺.

4. Transition metal ion solutions have a color.




Atomic Radii:

For the main group elements, covalent radii decrease in size across a row due to effective nuclear charge: the positive pull of the nucleus on the negative electrons. This is the positive nuclear charge that acts on the covalent electrons to pull them in, diminished (shielded) by any electrons between them.



Going down from the fourth period to the fifth period, the atomic radii increase. Further down to the sixth period, atomic radii are approximately the same as they are in the fifth period.

This similarity is due to lanthanide contraction. The lanthanides are the first row of the inner transition elements (placed at the bottom of the table), from cerium to lutetium. Here, the 4f subshell is filled, and by the time we get back to this row of transition elements with hafnium, the covalent radius is practically the same as the elements in the row above it.