What makes alkanes reactive

In general, the shorter the bond length, the greater the bond energy. When a chemical reaction occurs, the atoms in the reactants rearrange their chemical bonds to make products. The new arrangement of bonds does not have the same total energy as the bonds in the reactants. Therefore, when chemical reactions occur, there will always be an accompanying energy change. The energy change of a reaction or the heat of the reaction is called enthalpy. For an exothermic chemical reaction, energy is given off as reactants are converted to products.

For an endothermic chemical reaction, energy is absorbed as reactants are converted to products. In some reactions, the energy of the resulting products is lower than the energy of the reactants. Thus, in the course of the reaction, the excess energy released by product formation will be released to the surrounding environment.

Such reactions are exothermic and can be represented by an energy-level diagram in Figure 7. In most cases, the energy is given off as heat although a few reactions give off energy as light.

In chemical reactions where the products have a higher energy than the reactants, the reactants must absorb energy from their environment to be able to react. These reactions are endothermic and can be represented by an energy-level diagrams like Figure 7. Exothermic reactions release energy, so energy is a product. Endothermic reactions require energy, so energy is a reactant. Exothermic reactions will have a negative overall enthalpy and endothermic reactions will have a positive overall enthalpy.

The energy of a reaction can be treated stoichiometrically within the reaction just like any of the compounds within the reaction. For the combustion of methane, CH 4 , each mole that is burned will release kJ of energy. Thus, it is easy to calculate how much energy is released for any amount of CH 4 that is burned. We can also easily calculate how much CO 2 is produced for each mole of CH 4 that is burned, since one mole of CO 2 is produced for each mole of CH 4 burned.

Complete combustion given sufficient oxygen of any hydrocarbon produces carbon dioxide CO 2 and water H 2 O. It is quite important that you can write properly balanced equations for these reactions, because they often come up as a part of thermochemistry calculations. Some are easier than others. For example, with alkanes, the ones with an even number of carbon atoms are marginally harder than those with an odd number! For example, with propane C 3 H 8 , you can balance the carbons and hydrogens as you write the equation down.

Your first draft would be:. Counting the oxygens leads directly to the final version:. With butane C 4 H 10 , you can again balance the carbons and hydrogens as you write the equation down. Counting the oxygens leads to a slight problem — with 13 on the right-hand side. Having an odd number of oxygens on the product side makes it impossible to balance with the even number on the reactant side.

In cases like this, start trying to balance the equation, by changing the cofactor in front of the alkane to 2. Then rebalance the carbon and hydrogens on the product side. In this case, the oxygen number should come out to a positive number and you should now be able to balance the equation. The hydrocarbons become harder to ignite as the molecules get bigger.

If the liquid is not very volatile, only those molecules on the surface can react with the oxygen. Bigger molecules have greater Van der Waals attractions which makes it more difficult for them to break away from their neighbors and turn to a gas.

Provided the combustion is complete, all the hydrocarbons will burn with a blue flame. However, combustion tends to be less complete as the number of carbon atoms in the molecules rises.

That means that the bigger the hydrocarbon, the more likely you are to get a yellow, smoky flame. Incomplete combustion where there is not enough oxygen present can lead to the formation of carbon or carbon monoxide. As a simple way of thinking about it, the hydrogen in the hydrocarbon gets the first chance at the oxygen to form water in the product, and the carbon gets whatever is left over!

When only carbon is formed, the presence of glowing carbon particles in a flame turns it yellow, and black carbon is often visible in the smoke. If some oxygen can interact with the carbon, but not enough to form carbon dioxide CO 2 , then carbon monoxide CO is produced as a colorless poisonous gas.

Why carbon monoxide is poisonous : Oxygen is carried around the blood by hemoglobin, a protein found in red blood cells. Carbon dioxide also binds with hemoglobin. Hemoglobin carries oxygen from your lungs to every cell in your body, where it drops off the oxygen and picks up carbon dioxide and carries it back to the lungs. Hemoglobin will then exchange the carbon dioxide for the oxygen in the air that you breath in.

When you exhale, you release the carbon dioxide. Carbon monoxide can also bind to hemoglobin. The difference is that carbon monoxide binds irreversibly or very strongly — making that particular molecule of hemoglobin useless for carrying oxygen.

If you breath in enough carbon monoxide you will die from an internal form of suffocation! It is calculated the same way that enthalpy is calculated, however, it is specific for a set of standard conditions. Heat of Combustion is a useful value because it is a constant value for the type of material being burned and can be used to compare the efficiency and utility of different fuel sources used most often in society oil, gasoline, natural gas, diesel, coal, wood, hydrogen, ethanol, etc.

A common conversion to the metric units is:. It consists mainly of alkanes, cycloalkanes, and alkenes of various lengths and some additional minor organic compounds. Alkanes with five or more carbons are liquids and are found as common components of petroleum also called crude oil.

Many cycloalkanes are also found in petroleum products as well including, gasoline, kerosene, diesel, motor oil and many other heavy oils. Natural gas also contains trace levels of nitrogen, carbon dioxide CO 2 , and hydrogen sulfide H 2 S. Components of petroleum are separated by size using a technique called fractional distillation.

The resulting samples include gasoline with alkanes ranging from five to ten carbons in length and kerosene with a mixture of alkanes ranging in carbon length from ten to seventeen. Alkanes with longer carbon chains are found in diesel fuel, fuel oil, petroleum jelly, paraffin wax, motor oils, and the very highest chain length are used in asphalt.

It is estimated that the world uses about 95 million barrels of oil each day! The Environmental Protection Agency estimates that each barrel of oil produces 0. Note that a metric ton is equivalent to 1, kg. That means that 40,, metric tons of CO 2 or 40,,, kg of CO 2 are released into the atmosphere each day just from oil consumption! That is almost 15 billion metric tons of CO 2 per year! This calculation only includes the consumption of crude oil and not other common fuel sources, such as natural gas, coal, wood, and renewable energy sources like ethanol and biodiesel.

In Oregon, it is estimated that during the winter months that it take about 40 BTU per hour per square foot to heat an average, well-insulated home. If you owned a 2, square foot home, how many BTUs would you require to heat your own for 1 month? If you had the choice of heating your house with natural gas, coal lignite , or wood, how many kg of each of these fuels would be required to heat your home for one month?

How much CO2 would be produced by each fuel each winter month? In the presence of heat or light, alkanes can react with halogens to form alkyl halides or haloalkanes. This type of reaction is called a substitution reaction , because the halogen atom is taking the place of or substituting for one of the hydrogen atoms on the alkane structure. It should be noted that not all of the halogens react in the same way with alkanes.

But the reaction goes so fast, that this is the result:. In the presence of a flame, the reactions are rather like the fluorine one — producing a mixture of carbon and the hydrogen halide. The violence of the reaction drops considerably as you go from fluorine to chlorine to bromine.

The interesting reactions happen in the presence of ultra-violet light sunlight will do. These are photochemical reactions that happen at room temperature. In the substitution reaction, a hydrogen atom in the methane is replaced by a chlorine atom.

This can happen multiple times, until all the hydrogens are replaced. Ultimately, the longer the reaction proceeds, the more hydrogens in the alkane are replaced.

The original mixture of a colorless gas CH 4 and a green gas Cl 2 would produce steamy fumes of hydrogen chloride HCl and a mist of organic liquids mixture of the chlorinated methane. All of the organic products are liquid at room temperature with the exception of the chloromethane CH 3 Cl which is a gas. This substitution reaction is an example of a radical reaction, where only one electron is transferred at a time.

The heat or light, initiates the reaction by breaking the bond between the two Cl atoms in the chloride ion. This forms two radicals. A radical is an atom, molecule, or ion that has unpaired valence electrons. Thus, they are very unstable and reactive. In the diagram below, the first step of the halogenation reaction is shown below. This is termed initiation.

Once the radical is initiated, it will attack the alkane, in this case methane CH 4 , and create a new carbon radical. This stage of the reaction is termed propagation , as one radical species creates, or propagates, another radical.

The final stage of a radical reaction is the termination reaction which quenches the radical species present. For the methane — chlorine reaction this is the formation of the chloromethane CH 3 Cl. For example, with propane, you could get one of two isomers:. If chance was the only factor, you would expect to get three times as much of the isomer with the chlorine on the end. There are 6 hydrogens that could get replaced on the end carbon atoms compared with only 2 in the middle. In fact, you get about the same amount of each of the two isomers.

If you use bromine instead of chlorine, the great majority of the product is where the bromine is attached to the center carbon atom. Why does this happen?

It has to do with the stability of the carbon radical intermediate that forms during the reaction. Carbons that have more carbon neighbors will more easily lose a hydrogen and form a carbon radical intermediate. The neighboring carbons, being larger than the neighboring hydrogen atoms, can help stabilize the formation of the carbon radical.

Thus, in the halogenation reaction tertiary carbons will be the most reactive positions, followed by secondary carbons and finally primary carbons. Quaternary carbons are unreactive as they do not have any hydrogens available that can be substituted by a halogen. The reactions of the cycloalkanes are generally just the same as the alkanes, with hydrogen atoms on the cyclic ring structure being replaced by the halogen atom. For example i n the presence of UV light, cyclopropane will undergo substitution reactions with chlorine or bromine just like a non-cyclic alkane.

However, the small ring structures — particularly cyclopropane — also have the ability to react in the dark. In the absence of UV light, cyclopropane can undergo addition reactions in which the ring is broken. For example, with bromine, cyclopropane gives following linear compound. This can still happen in the presence of UV light — but you will get a mixture of the substitution reactions as well.

The ring is broken because cyclopropane suffers badly from ring strain. Why are alkanes sometimes called paraffins? Which halogen reacts most readily with alkanes? Which reacts least readily?

Alkanes do not react with many common chemicals. It should be clear from a review of the two steps that make up the free radical chain reaction for halogenation that the first step hydrogen abstraction is the product determining step. Once a carbon radical is formed, subsequent bonding to a halogen atom in the second step can only occur at the radical site. Since the H-X product is common to all possible reactions, differences in reactivity can only be attributed to differences in C-H bond dissociation energies.

In our previous discussion of bond energy we assumed average values for all bonds of a given kind, but now we see that this is not strictly true. In the case of carbon-hydrogen bonds, there are significant differences, and the specific dissociation energies energy required to break a bond homolytically for various kinds of C-H bonds have been measured. These values are given in the following table. By this reasoning we would expect benzylic and allylic sites to be exceptionally reactive in free radical halogenation, as experiments have shown.

The methyl group of toluene, C 6 H 5 CH 3 , is readily chlorinated or brominated in the presence of free radical initiators usually peroxides , and ethylbenzene is similarly chlorinated at the benzylic location exclusively.

The hydrogens bonded to the aromatic ring referred to as phenyl hydrogens above have relatively high bond dissociation energies and are not substituted. Since carbon-carbon double bonds add chlorine and bromine rapidly in liquid phase solutions, free radical substitution reactions of alkenes by these halogens must be carried out in the gas phase, or by other halogenating reagents. One such reagent is N-bromosuccinimide NBS , shown in the second equation below.

By using NBS as a brominating agent, allylic brominations are readily achieved in the liquid phase. The covalent bond homolyses that define the bond dissociation energies listed above may are described by the general equation:.

Because alkyl radicals are important intermediates in many reactions, this stability relationship will prove to be very useful in future discussions. The enhanced stability of allyl and benzyl radicals may be attributed to resonance stabilization. If you wish to review the principles of resonance Click Here. Formulas for the allyl and benzyl radicals are shown below.

Draw structural formulas for the chief canonical forms contributing to the resonance hybrid in each case. Addition Reactions of Dienes Addition reactions of isolated dienes proceed more or less as expected from the behavior of simple alkenes. Reagents are added through the formation of single bonds to carbon in an addition reaction.

Catalytic Hydrogenation of Alkenes II Alkene hydrogenation is the syn-addition of hydrogen to an alkene, saturating the bond. The alkene reacts with hydrogen gas in the presence of a metal catalyst which allows the reaction to occur quickly.

The energy released in this process, called the heat of hydrogenation, indicates the relative stabily of the double bond in the molecule see Catalytic Hydrogenation. Catalytic Hydrogenation of Alkenes II The carbon-carbon double bond in alkenes react with hydrogen in the presence of a metal catalyst. This attraction does not occur with alkanes because alkane molecules do not have this separation of charge. The net effect is that alkanes have a fairly restricted set of reactions, including the following:.

Cycloalkanes are very similar to the alkanes in reactivity, except for the very small cycloalkanes, especially cyclopropane. Cyclopropane is much more reactive than what is expected because of the bond angles in the ring. Normally, when carbon forms four single bonds, the bond angles are approximately With the electron pairs this close together, there is a significant amount of repulsion between the bonding pairs joining the carbon atoms, making the bonds easier to break.

More on reactivity of alkanes. Skip to main content. Search for:. Properties of alkanes Alkanes are organic compounds that consist entirely of single-bonded carbon and hydrogen atoms and lack any other functional groups.

Isomerism Alkanes with four or more carbon atoms can have more than one arrangement of atoms, so they can form structural isomers. Uses for alkanes The simplest alkane is methane CH 4 , a colorless, odorless gas that is the major component of natural gas.

Example The boiling points of the three isomers of C 5 H 12 are as follows: pentane Solubility Alkanes both normal and cycloalkanes are virtually insoluble in water but dissolve in organic solvents. When a molecular substance dissolves in water, the following must occur: breaking of the intermolecular forces within the substance. In the case of the alkanes, these are the Van der Waals dispersion forces. In water, the primary intermolecular attractions are hydrogen bonds.