Hydrogen Atom Spectrum

This page introduces the atomic hydrogen emission spectrum, showing how it arises from electron movements between energy levels within the atom. It also looks at how the spectrum can be used to find the ionisation energy of hydrogen. What is an emission spectrum? Observing hydrogen's emission spectrum A hydrogen discharge tube is a slim tube containing hydrogen gas at low pressure with an electrode at each end. If you put a high voltage across this (say, 5000 volts), the tube lights up with a bright pink glow. If the light is passed through a prism or diffraction grating, it is split into its various colours. What you would see is a small part of the hydrogen emission spectrum. Most of the spectrum is invisible to the eye because it is either in the infra-red or the ultra-violet. The photograph shows part of a hydrogen discharge tube on the left, and the three most easily seen lines in the visible part of the spectrum on the right. (Ignore the "smearing" - particularly to the left of the red line. This is caused by flaws in the way the photograph was taken. See note below.)
	Note:  This photograph is by courtesy of Dr Rod Nave of the Department of Physics and Astronomy at Georgia State University, Atlanta. The photograph comes from notes about the hydrogen spectrum in his HyperPhysics pages on the University site. If you are interested in more than an introductory look at the subject, that is a good place to go. Ideally the photo would show three clean spectral lines - dark blue, cyan and red. The red smearing which appears to the left of the red line, and other similar smearing (much more difficult to see) to the left of the other two lines probably comes, according to Dr Nave, from stray reflections in the set-up, or possibly from flaws in the diffraction grating. I have chosen to use this photograph anyway because a) I think it is a stunning image, and b) it is the only one I have ever come across which includes a hydrogen discharge tube and its spectrum in the same image.
Extending hydrogen's emission spectrum into the UV and IR There is a lot more to the hydrogen spectrum than the three lines you can see with the naked eye. It is possible to detect patterns of lines in both the ultra-violet and infra-red regions of the spectrum as well. These fall into a number of "series" of lines named after the person who discovered them. The diagram below shows three of these series, but there are others in the infra-red to the left of the Paschen series shown in the diagram. The diagram is quite complicated, so we will look at it a bit at a time. Look first at the Lyman series on the right of the diagram - this is the most spread out one and easiest to see what is happening.
	Note:  The frequency scale is marked in PHz - that's petaHertz. You are familiar with prefixes like kilo (meaning a thousand or 103 times), and mega (meaning a million or 106 times). Peta means 1015 times. So a value like 3 PHz means 3 x 1015 Hz. If you are worried about "Hertz", it just means "cycles per second".
The Lyman series is a series of lines in the ultra-violet. Notice that the lines get closer and closer together as the frequency increases. Eventually, they get so close together that it becomes impossible to see them as anything other than a continuous spectrum. That's what the shaded bit on the right-hand end of the series suggests. Then at one particular point, known as the series limit, the series stops. If you now look at the Balmer series or the Paschen series, you will see that the pattern is just the same, but the series have become more compact. In the Balmer series, notice the position of the three visible lines from the photograph further up the page. Complicating everything - frequency and wavelength You will often find the hydrogen spectrum drawn using wavelengths of light rather than frequencies. Unfortunately, because of the mathematical relationship between the frequency of light and its wavelength, you get two completely different views of the spectrum if you plot it against frequency or against wavelength. The relationship between frequency and wavelength The mathematical relationship is: Rearranging this gives equations for either wavelength or frequency. What this means is that there is an inverse relationship between the two - a high frequency means a low wavelength and vice versa.
	Note:  You will sometimes find frequency given the much more obvious symbol, f.
Drawing the hydrogen spectrum in terms of wavelength This is what the spectrum looks like if you plot it in terms of wavelength instead of frequency: . . . and just to remind you what the spectrum in terms of frequency looks like: Is this confusing? Well, I find it extremely confusing! So what do you do about it? For the rest of this page I shall only look at the spectrum plotted against frequency, because it is much easier to relate it to what is happening in the atom. Be aware that the spectrum looks different depending on how it is plotted, but, other than that, ignore the wavelength version unless it is obvious that your examiners want it. If you try to learn both versions, you are only going to get them muddled up!
	Note:  Syllabuses probably won't be very helpful about this. You need to look at past papers and mark schemes. If you are working towards a UK-based exam and don't have these things, you can find out how to get hold of them by going to the syllabuses page.
Explaining hydrogen's emission spectrum The Balmer and Rydberg Equations By an amazing bit of mathematical insight, in 1885 Balmer came up with a simple formula for predicting the wavelength of any of the lines in what we now know as the Balmer series. Three years later, Rydberg generalised this so that it was possible to work out the wavelengths of any of the lines in the hydrogen emission spectrum. What Rydberg came up with was: RH is a constant known as the Rydberg constant. n1 and n2 are integers (whole numbers). n2 has to be greater than n1. In other words, if n1 is, say, 2 then n2 can be any whole number between 3 and infinity. The various combinations of numbers that you can slot into this formula let you calculate the wavelength of any of the lines in the hydrogen emission spectrum - and there is close agreement between the wavelengths that you get using this formula and those found by analysing a real spectrum.
	Note:  If you come across a version of Balmer's original equation, it won't look like this. In Balmer's equation, n1 is always 2 - because that gives the wavelengths of the lines in the visible part of the spectrum which is what he was interested in. His original equation was also organised differently. The modern version shows more clearly what is going on.
You can also use a modified version of the Rydberg equation to calculate the frequency of each of the lines. You can work out this version from the previous equation and the formula relating wavelength and frequency further up the page.
	Note:  You may come across versions of the Rydberg equation where the n1 and n2 are the other way around, or they may even be swapped for letters like m and n. Whichever version you use, the bigger number must always be the one at the bottom of the right-hand term - the one you take away. If you get them the wrong way around, it is immediately obvious if you start to do a calculation, because you will end up with a negative answer!
The origin of the hydrogen emission spectrum The lines in the hydrogen emission spectrum form regular patterns and can be represented by a (relatively) simple equation. Each line can be calculated from a combination of simple whole numbers. Why does hydrogen emit light when it is excited by being exposed to a high voltage and what is the significance of those whole numbers? When nothing is exciting it, hydrogen's electron is in the first energy level - the level closest to the nucleus. But if you supply energy to the atom, the electron gets excited into a higher energy level - or even removed from the atom altogether. The high voltage in a discharge tube provides that energy. Hydrogen molecules are first broken up into hydrogen atoms (hence the atomic hydrogen emission spectrum) and electrons are then promoted into higher energy levels. Suppose a particular electron was excited into the third energy level. This would tend to lose energy again by falling back down to a lower level. It could do this in two different ways. It could fall all the way back down to the first level again, or it could fall back to the second level - and then, in a second jump, down to the first level. Tying particular electron jumps to individual lines in the spectrum If an electron falls from the 3-level to the 2-level, it has to lose an amount of energy exactly the same as the energy gap between those two levels. That energy which the electron loses comes out as light (where "light" includes UV and IR as well as visible). Each frequency of light is associated with a particular energy by the equation: The higher the frequency, the higher the energy of the light. If an electron falls from the 3-level to the 2-level, red light is seen. This is the origin of the red line in the hydrogen spectrum. By measuring the frequency of the red light, you can work out its energy. That energy must be exactly the same as the energy gap between the 3-level and the 2-level in the hydrogen atom. The last equation can therefore be re-written as a measure of the energy gap between two electron levels. The greatest possible fall in energy will therefore produce the highest frequency line in the spectrum. The greatest fall will be from the infinity level to the 1-level. (The significance of the infinity level will be made clear later.) The next few diagrams are in two parts - with the energy levels at the top and the spectrum at the bottom. If an electron fell from the 6-level, the fall is a little bit less, and so the frequency will be a little bit lower. (Because of the scale of the diagram, it is impossible to draw in all the jumps involving all the levels between 7 and infinity!) . . . and as you work your way through the other possible jumps to the 1-level, you have accounted for the whole of the Lyman series. The spacings between the lines in the spectrum reflect the way the spacings between the energy levels change. If you do the same thing for jumps down to the 2-level, you end up with the lines in the Balmer series. These energy gaps are all much smaller than in the Lyman series, and so the frequencies produced are also much lower. The Paschen series would be produced by jumps down to the 3-level, but the diagram is going to get very messy if I include those as well - not to mention all the other series with jumps down to the 4-level, the 5-level and so on. The significance of the numbers in the Rydberg equation n1 and n2 in the Rydberg equation are simply the energy levels at either end of the jump producing a particular line in the spectrum. For example, in the Lyman series, n1 is always 1. Electrons are falling to the 1-level to produce lines in the Lyman series. For the Balmer series, n1 is always 2, because electrons are falling to the 2-level. n2 is the level being jumped from. We have already mentioned that the red line is produced by electrons falling from the 3-level to the 2-level. In this case, then, n2 is equal to 3. The significance of the infinity level The infinity level represents the highest possible energy an electron can have as a part of a hydrogen atom. So what happens if the electron exceeds that energy by even the tiniest bit? The electron is no longer a part of the atom. The infinity level represents the point at which ionisation of the atom occurs to form a positively charged ion. Using the spectrum to find hydrogen's ionisation energy When there is no additional energy supplied to it, hydrogen's electron is found at the 1-level. This is known as its ground state. If you supply enough energy to move the electron up to the infinity level, you have ionised the hydrogen. The ionisation energy per electron is therefore a measure of the distance between the 1-level and the infinity level. If you look back at the last few diagrams, you will find that that particular energy jump produces the series limit of the Lyman series.
	Note:  Up to now we have been talking about the energy released when an electron falls from a higher to a lower level. Obviously if a certain amount of energy is released when an electron falls from the infinity level to the 1-level, that same amount will be needed to push the electron from the 1-level up to the infinity level.
If you can determine the frequency of the Lyman series limit, you can use it to calculate the energy needed to move the electron in one atom from the 1-level to the point of ionisation. From that, you can calculate the ionisation energy per mole of atoms. The problem is that the frequency of a series limit is quite difficult to find accurately from a spectrum because the lines are so close together in that region that the spectrum looks continuous. Finding the frequency of the series limit graphically Here is a list of the frequencies of the seven most widely spaced lines in the Lyman series, together with the increase in frequency as you go from one to the next. As the lines get closer together, obviously the increase in frequency gets less. At the series limit, the gap between the lines would be literally zero. That means that if you were to plot the increases in frequency against the actual frequency, you could extrapolate (continue) the curve to the point at which the increase becomes zero. That would be the frequency of the series limit. In fact you can actually plot two graphs from the data in the table above. The frequency difference is related to two frequencies. For example, the figure of 0.457 is found by taking 2.467 away from 2.924. So which of these two values should you plot the 0.457 against? It doesn't matter, as long as you are always consistent - in other words, as long as you always plot the difference against either the higher or the lower figure. At the point you are interested in (where the difference becomes zero), the two frequency numbers are the same. As you will see from the graph below, by plotting both of the possible curves on the same graph, it makes it easier to decide exactly how to extrapolate the curves. Because these are curves, they are much more difficult to extrapolate than if they were straight lines. Both lines point to a series limit at about 3.28 x 1015 Hz.
	Note:  Remember that 3.28 PHz is the same as 3.28 x 1015 Hz. You can use the Rydberg equation to calculate the series limit of the Lyman series as a check on this figure: n1 = 1 for the Lyman series, and n2 = infinity for the series limit. 1/(infinity)2 = zero. That gives a value for the frequency of 3.29 x 1015 Hz - in other words the two values agree to within 0.3%.
So . . . now we can calculate the energy needed to remove a single electron from a hydrogen atom. Remember the equation from higher up the page: We can work out the energy gap between the ground state and the point at which the electron leaves the atom by substituting the value we've got for frequency and looking up the value of Planck's constant from a data book. That gives you the ionisation energy for a single atom. To find the normally quoted ionisation energy, we need to multiply this by the number of atoms in a mole of hydrogen atoms (the Avogadro constant) and then divide by 1000 to convert it into kilojoules.
	Note:  It would be wrong to quote this to more than 3 significant figures. The value for the frequency obtained from the graph is only to that accuracy.
This compares well with the normally quoted value for hydrogen's ionisation energy of 1312 kJ mol-1.

© 2010-2012 The Little Ganesha® All Rights Reserved. Subscribe in a reader

Aufbau principle and Hund's Rule

In the independent-particle model (i.e., the all-electron wave function is written as a Slater determinant) we have in general more spin orbitals than we have electrons. This means that we have to make a choice with which orbitals to construct the wave function. The choice is not difficult, but the procedure is often presented in separated parts, dealing with different cases. Here these parts are put together. We use the Aufbau principle and Hund's rules to make the choice. The Aufbau principle says that you should look at the orbital energies, and take the orbitals with the lowest energies. In almost all cases this is the only rule you need. Suppose you have (spatial) orbitals with increasing energies e0, e1, e2, e3, ..., and you have six electrons. Then you can put two electrons in the orbitals with energies e0, e1, and e2 each. The orbital diagram looks as follows
	The Aufbau principle does not suffice if the highest level in which you put electrons is degenerate and there are various ways to distribute the electrons. There is not problem with one electron in a highest non-empty degenerate level. All possibilities for this one electron yield the same energy, so there is no preference. The same holds if the level would be completely occupied if one extra electron would be added. (The level then has one hole). Hund's rules (sometimes also collective called Hund's rule) tell what to do if there are at least two electrons and at least two holes. Hund's first rule says that one should distribute the electrons as much as possible over the spatial orbitals. The reason is that in this way the electrons stay away from each other as much as possible thus reducing their repulsion. Hund's second rule says that unpaired electrons should be given the same spin. In this way the exchange interaction is maximized. As this interaction has a negative sign in the Fock operator and the expression for the electronic energy, this reduces the energy. The following diagram shows how to apply all three rules.

Alkaline Earth Metals

The elements of Group 2, the Alkaline Earth Metals, are:   symbol electron configuration beryllium Be [He]2s2 magnesium Mg [Ne]3s2 calcium Ca [Ar]4s2 strontium Sr [Kr]5s2 barium Ba [Xe]6s2 radium Ra [Rn]7s2 The last element, radium, is radioactive and will not be considered here. Appearance The Group 2 elements are all metals with a shiny, silvery-white colour. General Reactivity The alkaline earth metals are high in the reactivity series of metals, but not as high as the alkali metals of Group 1. Occurrence and Extraction These elements are all found in the Earth’s crust, but not in the elemental form as they are so reactive. Instead, they are widely distributed in rock structures. The main minerals in which magnesium is found are carnellite, magnesite and dolomite. Calcium is found in chalk, limestone, gypsum and anhydrite. Magnesium is the eighth most abundant element in the Earth’s crust, and calcium is the fifth. Of the elements in this Group only magnesium is produced on a large scale. It is extracted from sea-water by the addition of calcium hydroxide, which precipitates out the less soluble magnesium hydroxide. This hydroxide is then converted to the chloride, which is electrolysed in a Downs cell to extract magnesium metal. Physical Properties The metals of Group 2 are harder and denser than sodium and potassium, and have higher melting points. These properties are due largely to the presence of two valence electrons on each atom, which leads to stronger metallic bonding than occurs in Group 1. Three of these elements give characteristic colours when heated in a flame: Mg brilliant white   Ca brick-red   Sr crimson   Ba apple green Atomic and ionic radii increase smoothly down the Group. The ionic radii are all much smaller than the corresponding atomic radii. This is because the atom contains two electrons in an s level relatively far from the nucleus, and it is these electrons which are removed to form the ion. Remaining electrons are thus in levels closer to the nucleus, and in addition the increased effective nuclear charge attracts the electrons towards the nucleus and decreases the size of the ion. Chemical Properties The chemical properties of Group 2 elements are dominated by the strong reducing power of the metals. The elements become increasingly electropositive on descending the Group. Once started, the reactions with oxygen and chlorine are vigorous: 2Mg(s) + O2(g) ® 2MgO(s) Ca(s) + Cl2(g) ® CaCl2(s) All the metals except beryllium form oxides in air at room temperature which dulls the surface of the metal. Barium is so reactive it is stored under oil. All the metals except beryllium reduce water and dilute acids to hydrogen: Mg(s) + 2H+(aq) ® Mg(aq) + H2(g) Magnesium reacts only slowly with water unless the water is boiling, but calcium reacts rapidly even at room temperature, and forms a cloudy white suspension of sparingly soluble calcium hydroxide. Calcium, strontium and barium can reduce hydrogen gas when heated, forming the hydride: Ca(s) + H2(g) ® CaH2(s) The hot metals are also sufficiently strong reducing agents to reduce nitrogen gas and form nitrides: 3Mg(s) + N2(g) ® Mg3N2(s) Magnesium can reduce, and burn in, carbon dioxide: 2Mg(s) + CO2(g) ® 2MgO(s) + C(s) This means that magnesium fires cannot be extinguished using carbon dioxide fire extinguishers. Oxides The oxides of alkaline earth metals have the general formula MO and are basic. They are normally prepared by heating the hydroxide or carbonate to release carbon dioxide gas. They have high lattice enthalpies and melting points. Peroxides, MO2, are known for all these elements except beryllium, as the Be2+ cation is too small to accommodate the peroxide anion. Hydroxides Calcium, strontium and barium oxides react with water to form hydroxides: CaO(s) + H2O(l) ® Ca(OH)2(s) Calcium hydroxide is known as slaked lime. It is sparingly soluble in water and the resulting mildly alkaline solution is known as lime water which is used to test for the acidic gas carbon dioxide. Halides The Group 2 halides are normally found in the hydrated form. They are all ionic except beryllium chloride. Anhydrous calcium chloride has such a strong affinity for water it is used as a drying agent. Oxidation States and lonisation Energies In all their compounds these metals have an oxidation number of +2 and, with few exceptions, their compounds are ionic. The reason for this can be seen by examination of the electron configuration, which always has two electrons in an outer quantum level. These electrons are relatively easy to remove, but removing the third electron is much more difficult, as it is close to the nucleus and in a filled quantum shell. This results in the formation of M2+. The ionisation energies reflect this electron arrangement. The first two ionisation energies are relatively low, and the third very much higher. Industrial Information Magnesium is the only Group 2 element used on a large scale. It is used in flares, tracer bullets and incendiary bombs as it burns with a brilliant white light. It is also alloyed with aluminium to produce a low-density, strong material used in aircraft. Magnesium oxide has such a high melting point it is used to line furnaces. Further Information For further information look up the individual elements. Data   Atomic Number Relative Atomic Mass Melting Point/K Density/kg m-3 Be 4   9.012   1551   1847.7   Mg 12   24.31   922   1738   Ca 20   40.08   1112   1550   Sr 38   87.62   1042   2540   Ba 56   137.33   1002   3594   Ionisation Energies/kJ mol-1   1st 2nd 3rd Be 899.4 1757.1 14848 Mg 737.7 1450.7 7732.6 Ca 589.7 1145 4910 Sr 549.5 1064.2 4210 Ba 502.8 965.1 3600   Atomic Radius/nm Ionic Radius/nm (M2+) Standard Electrode Potentials/V Be 0.113 0.034 -1.85 Mg 0.160 0.078 -2.36 Ca 0.197 0.106 -2.87 Sr 0.215 0.127 -2.89 Ba 0.217 0.143 -2.90

Chemical Equilibrium

Definition of Chemical Equilibrium

Chemical equilibrium applies to reactions that can occur in both directions. In a reaction such as:
CH₄(g) + H₂O(g) <--> CO(g) + 3H₂(g) The reaction can happen both ways. So after some of the products are created the products begin to react to form the reactants. At the beginning of the reaction, the rate that the reactants are changing into the products is higher than the rate that the products are changing into the reactants. Therefore, the net change is a higher number of products.
Even though the reactants are constantly forming products and vice-versa the amount of reactants and products does become steady. When the net change of the products and reactants is zero the reaction has reached equilibrium. The equilibrium is a dynamic equilibrium. The definition for a dynamic equilibrium is when the amount of products and reactants are constant. (They are not equal but constant. Also, both reactions are still occurring.)

Equilibrium Constant

To determine the amount of each compound that will be present at equilibrium you must know the equilibrium constant. To determine the equilibrium constant you must consider the generic equation:
aA + bB <--> cC + dD The upper case letters are the molar concentrations of the reactants and products. The lower case letters are the coefficients that balance the equation. Use the following equation to determine the equilibrium constant (K_c).
For example, determining the equilibrium constant of the following equation can be accomplished by using the K_c equation.
Using the following equation, calculate the equilibrium constant.
N₂(g) + 3H₂(g) <--> 2NH₃(g)
A one-liter vessel contains 1.60 moles NH₃, .800 moles N₂, and 1.20 moles of H₂. What is the equilibrium constant?

Answer: 1.85

Le Chatelier's Principle

Le Chatelier's principle states that when a system in chemical equilibrium is disturbed by a change of temperature, pressure, or a concentration, the system shifts in equilibrium composition in a way that tends to counteract this change of variable. The three ways that Le Chatelier's principle says you can affect the outcome of the equilibrium are as follows:

• Changing concentrations by adding or removing products or reactants to the reaction vessel.
• Changing partial pressure of gaseous reactants and products.
• Changing the temperature.

These actions change each equilibrium differently, therefore you must determine what needs to happen for the reaction to get back in equilibrium.

Example involving change of concentration:

In the equation
2NO_(g) + O_2(g) <--> 2NO_2(g) If you add more NO_(g) the equilibrium shifts to the right producing more NO_2(g)
If you add more O_2(g) the equilibrium shifts to the right producing more NO_2(g)
If you add more NO_2(g) the equilibrium shifts to the left producing more NO_(g) and O_2(g)

Example involving pressure change:

In the equation
2SO_2(g) + O_2(g) <--> 2SO_3(g), an increase in pressure will cause the reaction to shift in the direction that reduces pressure, that is the side with the fewer number of gas molecules. Therefore an increase in pressure will cause a shift to the right, producing more product. (A decrease in volume is one way of increasing pressure.)

Example involving temperature change:

In the equation
N_2(g) + 3H_2(g) <--> 2NH₃ + 91.8 kJ, an increase in temperature will cause a shift to the left because the reverse reaction uses the excess heat. An increase in forward reaction would produce even more heat since the forward reaction is exothermic. Therefore the shift caused by a change in temperature depends upon whether the reaction is exothermic or endothermic.

IUPAC Nomenclature for organic chemistry

Nomenclature

Naming Organic Compounds

The increasingly large number of organic compounds identified with each passing day, together with the fact that many of these compounds are isomers of other compounds, requires that a systematic nomenclature system be developed. Just as each distinct compound has a unique molecular structure which can be designated by a structural formula, each compound must be given a characteristic and unique name.
As organic chemistry grew and developed, many compounds were given trivial names, which are now commonly used and recognized. Some examples are:

Name	Methane	Butane	Acetone	Toluene	Acetylene	Ethyl Alcohol
Formula	CH4	C4H10	CH3COCH3	CH3C6H5	C2H2	C2H5OH

Such common names often have their origin in the history of the science and the natural sources of specific compounds, but the relationship of these names to each other is arbitrary, and no rational or systematic principles underly their assignments.

The IUPAC Systematic Approach to Nomenclature

A rational nomenclature system should do at least two things. First, it should indicate how the carbon atoms of a given compound are bonded together in a characteristic lattice of chains and rings. Second, it should identify and locate any functional groups present in the compound. Since hydrogen is such a common component of organic compounds, its amount and locations can be assumed from the tetravalency of carbon, and need not be specified in most cases.
The IUPAC nomenclature system is a set of logical rules devised and used by organic chemists to circumvent problems caused by arbitrary nomenclature. Knowing these rules and given a structural formula, one should be able to write a unique name for every distinct compound. Likewise, given a IUPAC name, one should be able to write a structural formula. In general, an IUPAC name will have three essential features:
•  A root or base indicating a major chain or ring of carbon atoms found in the molecular structure.
•  A suffix or other element(s) designating functional groups that may be present in the compound.
•  Names of substituent groups, other than hydrogen, that complete the molecular structure. As an introduction to the IUPAC nomenclature system, we shall first consider compounds that have no specific functional groups. Such compounds are composed only of carbon and hydrogen atoms bonded together by sigma bonds (all carbons are sp³ hybridized).
An excellent presentation of organic nomenclature is provided on a Nomenclature Page. created by Dave Woodcock.

Alkanes

Alkanes

Hydrocarbons having no double or triple bond functional groups are classified as alkanes or cycloalkanes, depending on whether the carbon atoms of the molecule are arranged only in chains or also in rings. Although these hydrocarbons have no functional groups, they constitute the framework on which functional groups are located in other classes of compounds, and provide an ideal starting point for studying and naming organic compounds. The alkanes and cycloalkanes are also members of a larger class of compounds referred to as aliphatic. Simply put, aliphatic compounds are compounds that do not incorporate any aromatic rings in their molecular structure.
The following table lists the IUPAC names assigned to simple continuous-chain alkanes from C-1 to C-10. A common "ane" suffix identifies these compounds as alkanes. Longer chain alkanes are well known, and their names may be found in many reference and text books. The names methane through decane should be memorized, since they constitute the root of many IUPAC names. Fortunately, common numerical prefixes are used in naming chains of five or more carbon atoms.

Examples of Simple Unbranched Alkanes

Name	Molecular Formula	Structural Formula	Isomers		Name	Molecular Formula	Structural Formula	Isomers
methane	CH4	CH4	1		hexane	C6H14	CH3(CH2)4CH3	5
ethane	C2H6	CH3CH3	1		heptane	C7H16	CH3(CH2)5CH3	9
propane	C3H8	CH3CH2CH3	1		octane	C8H18	CH3(CH2)6CH3	18
butane	C4H10	CH3CH2CH2CH3	2		nonane	C9H20	CH3(CH2)7CH3	35
pentane	C5H12	CH3(CH2)3CH3	3		decane	C10H22	CH3(CH2)8CH3	75

Some important behavior trends and terminologies:
  (i)   The formulas and structures of these alkanes increase uniformly by a CH₂ increment.
 (ii)   A uniform variation of this kind in a series of compounds is called homologous.
(iii)   These formulas all fit the C_nH_2n+2 rule. This is also the highest possible H/C ratio for a stable hydrocarbon.
(iv)   Since the H/C ratio in these compounds is at a maximum, we call them saturated (with hydrogen).
Beginning with butane (C₄H₁₀), and becoming more numerous with larger alkanes, we note the existence of alkane isomers. For example, there are five C₆H₁₄ isomers, shown below as abbreviated line formulas (A through E):

Although these distinct compounds all have the same molecular formula, only one (A) can be called hexane. How then are we to name the others?
The IUPAC system requires first that we have names for simple unbranched chains, as noted above, and second that we have names for simple alkyl groups that may be attached to the chains. Examples of some common alkyl groups are given in the following table. Note that the "ane" suffix is replaced by "yl" in naming groups. The symbol R is used to designate a generic (unspecified) alkyl group.

Group	CH3–	C2H5–	CH3CH2CH2–	(CH3)2CH–	CH3CH2CH2CH2–	(CH3)2CHCH2–	CH3CH2CH(CH3)–	(CH3)3C–	R–
Name	Methyl	Ethyl	Propyl	Isopropyl	Butyl	Isobutyl	sec-Butyl	tert-Butyl	Alkyl

IUPAC Rules for Alkane Nomenclature  1.   Find and name the longest continuous carbon chain.  2.   Identify and name groups attached to this chain.  3.   Number the chain consecutively, starting at the end nearest a substituent group.  4.   Designate the location of each substituent group by an appropriate number and name.  5.   Assemble the name, listing groups in alphabetical order.     The prefixes di, tri, tetra etc., used to designate several groups of the same kind, are not considered when alphabetizing.

For the above isomers of hexane the IUPAC names are:   B  2-methylpentane    C  3-methylpentane    D  2,2-dimethylbutane    E  2,3-dimethylbutane Halogen substituents are easily accommodated, using the names: fluoro (F-), chloro (Cl-), bromo (Br-) and iodo (I-). For example, (CH₃)₂CHCH₂CH₂Br would be named 1-bromo-3-methylbutane. If the halogen is bonded to a simple alkyl group an alternative "alkyl halide" name may be used. Thus, C₂H₅Cl may be named chloroethane (no locator number is needed for a two carbon chain) or ethyl chloride.


 

Cycloalkanes

Cycloalkanes

      Cycloalkanes have one or more rings of carbon atoms. The simplest examples of this class consist of a single, unsubstituted carbon ring, and these form a homologous series similar to the unbranched alkanes. The IUPAC names of the first five members of this series are given in the following table. The last (yellow shaded) column gives the general formula for a cycloalkane of any size. If a simple unbranched alkane is converted to a cycloalkane two hydrogen atoms, one from each end of the chain, must be lost. Hence the general formula for a cycloalkane composed of n carbons is C_nH_2n. Although a cycloalkane has two fewer hydrogens than the equivalent alkane, each carbon is bonded to four other atoms so such compounds are still considered to be saturated with hydrogen.

Examples of Simple Cycloalkanes

Name	Cyclopropane	Cyclobutane	Cyclopentane	Cyclohexane	Cycloheptane	Cycloalkane
Molecular Formula	C3H6	C4H8	C5H10	C6H12	C7H14	CnH2n
Structural Formula						(CH2)n
Line Formula

Substituted cycloalkanes are named in a fashion very similar to that used for naming branched alkanes. The chief difference in the rules and procedures occurs in the numbering system. Since all the carbons of a ring are equivalent (a ring has no ends like a chain does), the numbering starts at a substituted ring atom.

IUPAC Rules for Cycloalkane Nomenclature  1.   For a monosubstituted cycloalkane the ring supplies the root name (table above) and the substituent group is named as usual. A location number is unnecessary.  2.   If the alkyl substituent is large and/or complex, the ring may be named as a substituent group on an alkane.  3.   If two different substituents are present on the ring, they are listed in alphabetical order, and the first cited substituent is assigned to carbon #1. The numbering of ring carbons then continues in a direction (clockwise or counter-clockwise) that affords the second substituent the lower possible location number.  4.   If several substituents are present on the ring, they are listed in alphabetical order. Location numbers are assigned to the substituents so that one of them is at carbon #1 and the other locations have the lowest possible numbers, counting in either a clockwise or counter-clockwise direction.  5.   The name is assembled, listing groups in alphabetical order and giving each group (if there are two or more) a location number. The prefixes di, tri, tetra etc., used to designate several groups of the same kind, are not considered when alphabetizing.

For examples of how these rules are used in naming substituted cycloalkanes   .
Small rings, such as three and four membered rings, have significant angle strain resulting from the distortion of the sp³ carbon bond angles from the ideal 109.5º to 60º and 90º respectively. This angle strain often enhances the chemical reactivity of such compounds, leading to ring cleavage products. It is also important to recognize that, with the exception of cyclopropane, cycloalkyl rings are not planar (flat). The three dimensional shapes assumed by the common rings (especially cyclohexane and larger rings) are described and discussed in the Conformational Analysis Section.
Hydrocarbons having more than one ring are common, and are referred to as bicyclic (two rings), tricyclic (three rings) and in general, polycyclic compounds. The molecular formulas of such compounds have H/C ratios that decrease with the number of rings. In general, for a hydrocarbon composed of n carbon atoms associated with m rings the formula is: C_nH_{(2n + 2 - 2m)}. The structural relationship of rings in a polycyclic compound can vary. They may be separate and independent, or they may share one or two common atoms. Some examples of these possible arrangements are shown in the following table.

Examples of Isomeric C₈H₁₄ Bicycloalkanes

Isolated Rings	Spiro Rings	Fused Rings	Bridged Rings
No common atoms	One common atom	One common bond	Two common atoms