Strongest carbon carbon bond

Strongest carbon carbon bond DEFAULT


Learning Objective

  • Describe the types of orbital overlap that occur in single, double, and triple bonds

Key Points

    • Double and triple covalent bonds are stronger than single covalent bonds and they are characterized by the sharing of four or six electrons between atoms, respectively.
    • Double and triple bonds are comprised of sigma bonds between hybridized orbitals, and pi bonds between unhybridized p orbitals. Double and triple bonds offer added stability to compounds, and restrict any rotation around the bond axis.
    • Bond lengths between atoms with multiple bonds are shorter than in those with single bonds.


  • bond strengthDirectly related to the amount of energy required to break the bond between two atoms. The more energy required, the stronger the bond is said to be.
  • bond lengthThe distance between the nuclei of two bonded atoms. It can be experimentally determined.
  • orbital hybridizationThe concept of mixing atomic orbitals to form new hybrid orbitals suitable for the qualitative description of atomic bonding properties and geometries.
  • atomic orbitalsThe physical region in space around the nucleus where an electron has a probability of being.

Double and Triple Covalent Bonds

Covalent bonding occurs when electrons are shared between atoms. Double and triple covalent bonds occur when four or six electrons are shared between two atoms, and they are indicated in Lewis structures by drawing two or three lines connecting one atom to another. It is important to note that only atoms with the need to gain or lose at least two valence electrons through sharing can participate in multiple bonds.

Bonding Concepts


Double and triple bonds can be explained by orbital hybridization, or the ‘mixing’ of atomic orbitals to form new hybrid orbitals. Hybridization describes the bonding situation from a specific atom’s point of view. A combination of s and p orbitals results in the formation of hybrid orbitals. The newly formed hybrid orbitals all have the same energy and have a specific geometrical arrangement in space that agrees with the observed bonding geometry in molecules. Hybrid orbitals are denoted as spx, where s and p denote the orbitals used for the mixing process, and the value of the superscript x ranges from 1-3, depending on how many p orbitals are required to explain the observed bonding.

Pi Bonds

Pi, or [latex]\pi[/latex], bonds occur when there is overlap between unhybridized p orbitals of two adjacent atoms. The overlap does not occur between the nuclei of the atoms, and this is the key difference between sigma and pi bonds. For the bond to form efficiently, there has to be a proper geometrical relationship between the unhybridized p orbitals: they must be on the same plane.

Multiple bonds between atoms always consist of a sigma bond, with any additional bonds being of the π type.

Examples of Pi Bonds

The simplest example of an organic compound with a double bond is ethylene, or ethene, C2H4. The double bond between the two carbon atoms consists of a sigma bond and a π bond.

From the perspective of the carbon atoms, each has three sp2 hybrid orbitals and one unhybridized p orbital. The three sp2 orbitals lie in a single plane at 120-degree angles. As the carbon atoms approach each other, their orbitals overlap and form a bond. Simultaneously, the p orbitals approach each other and form a bond. To maintain this bond, the p orbitals must stay parallel to each other; therefore, rotation is not possible.

A triple bond involves the sharing of six electrons, with a sigma bond and two [latex]\pi[/latex] bonds. The simplest triple-bonded organic compound is acetylene, C2H2. Triple bonds are stronger than double bonds due to the the presence of two [latex]\pi[/latex] bonds rather than one. Each carbon has two sp hybrid orbitals, and one of them overlaps with its corresponding one from the other carbon atom to form an sp-sp sigma bond. The remaining four unhybridized p orbitals overlap with each other and form two [latex]\pi[/latex] bonds. Similar to double bonds, no rotation around the triple bond axis is possible.

Observable Consequences of Multiple Bonds

Bond Strength

Covalent bonds can be classified in terms of the amount of energy that is required to break them. Based on the experimental observation that more energy is needed to break a bond between two oxygen atoms in O2 than two hydrogen atoms in H2, we infer that the oxygen atoms are more tightly bound together. We say that the bond between the two oxygen atoms is stronger than the bond between two hydrogen atoms.

Experiments have shown that double bonds are stronger than single bonds, and triple bonds are stronger than double bonds. Therefore, it would take more energy to break the triple bond in N2 compared to the double bond in O2. Indeed, it takes 497 kcal/mol to break the O2 molecule, while it takes 945 kJ/mol to do the same to the N2 molecule.

Bond Length

Another consequence of the presence of multiple bonds between atoms is the difference in the distance between the nuclei of the bonded atoms. Double bonds have shorter distances than single bonds, and triple bonds are shorter than double bonds.




"We aimed to develop a technology that could degrade polyfluoroalkyl substances (PFAS), one of the most challenging pollutant remediation problems of the present day," said Jaehong Kim, a professor in the department of chemical and environmental engineering at Yale University. "PFAS are widely detected all over the world, from Arctic biota to the human body, and concentrations in contaminated groundwater significantly exceed the regulatory limit in many areas. Currently, there are no energy-efficient methods to destroy these contaminants. Our collaboration with Brookhaven Lab aims to solve this problem by taking advantage of the unique properties of single atom catalysts."

Synthesizing smaller, more efficient catalysts

To optimize the efficiency of catalysts -- substances that initiate or speed-up chemical reactions -- scientists break them down into smaller pieces, all the way down to nanomaterials. And recently, scientists have started to break catalysts down even further, beyond the nanoscale and into single atoms.

"The only part of a catalyst that is reactive is its surface," said Brookhaven scientist Eli Stavitski. "So, if you break a catalyst down into small pieces, you increase its surface area and expose more of the catalyst's reactive properties. But also, when you break catalysts down below 10 nanometers, their electronic properties change dramatically. They suddenly become very reactive. Ultimately, you want to go to the next step, and break catalysts down to individual atoms."

The challenge is that individual atoms don't behave the same as larger catalysts; they don't like to stand alone, and they can cause unwanted side reactions to occur. To use single atom catalysts effectively, scientists must identify the perfect combination of a strong, reactive metal and a stable, complementary environment.

Now, researchers have identified single atoms of platinum as an efficient catalyst for breaking carbon-fluorine bonds. Platinum is an especially strong metal, and it is capable of splitting hydrogen gas into individual hydrogen atoms -- a key step towards breaking the carbon-fluorine bond.

"Our team at Yale recently developed a readily scalable method to synthesize single-atom catalysts in two simple steps," said Kim. "First, we bind metals to anchor sites on a support material, then we photoreduce the metals to single atoms under mild UV-C irradiation. Using this method, our group has been synthesizing a suite of single-atom catalysts involving various metals (platinum, palladium, and cobalt) and supports (silicon carbide, carbon nitride, and titanium dioxide) for numerous catalytic reactions. In this work, we found single platinum atoms loaded onto silicon carbide to be strikingly effective in catalyzing carbon-fluoride bond cleavage and breaking down contaminants like PFAS."

Imaging single atoms

To visualize their new catalyst and assess its performance, the scientists came to two DOE Office of Science User Facilities at Brookhaven Lab -- the Center for Functional Nanomaterials (CFN) and the National Synchrotron Light Source II (NSLS-II). The world-class tools at each facility provided complementary techniques for seeing this incredibly small catalyst.

At CFN, the scientists used an advanced transmission electron microscope (TEM) to get a close-up view of the platinum atoms. By scanning an electron probe over the sample, the scientists were able to visualize discrete platinum atoms on the silicon carbide support.

"This research offers a golden standard for showing how multimodal characterization can contribute to the understanding of fundamental reaction mechanisms of single atom catalysts," said Huolin Xin, a former scientific staff member at CFN and now a professor at University of California.

Compared to the smaller, more focused view of the catalyst that CFN could provide, NSLS-II enabled the researchers to get a broader view of the catalyst and its surrounding environment.

"We have a technique at NSLS-II, called x-ray absorption spectroscopy, that is uniquely sensitive to the state of the catalyst and the environment surrounding it," said Stavitski, who is also a beamline scientist at NSLS-II's Inner-Shell Spectroscopy (ISS) beamline, where the research was conducted.

By shining NSLS-II's ultrabright x-ray light onto the catalyst and using ISS to see how the light interacted with the sample and its environment, the scientists were able to "see" how the single-atom catalyst was built.

The research at ISS was part of NSLS-II's strategic partnership with Yale University, and illustrates how universities and industry can work with Brookhaven Lab to solve their research challenges.

"We are pursuing a number of strategic partnerships to strengthen our connections with nearby institutions and to leverage the tremendous intellectual power and expertise in the northeastern U.S.," said Qun Shen, the NSLS-II Deputy Director for Science. "Yale faculty groups are an excellent example in this regard. We are happy to see this is starting to bear fruit."

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Story Source:

Materials provided by DOE/Brookhaven National Laboratory. Note: Content may be edited for style and length.

Journal Reference:

  1. Dahong Huang, Glen Andrew de Vera, Chiheng Chu, Qianhong Zhu, Eli Stavitski, Jing Mao, Huolin Xin, Jacob A. Spies, Charles A. Schmuttenmaer, Junfeng Niu, Gary L. Haller, Jae-Hong Kim. Single-Atom Pt Catalyst for Effective C–F Bond Activation via Hydrodefluorination. ACS Catalysis, 2018; 9353 DOI: 10.1021/acscatal.8b02660

Cite This Page:

DOE/Brookhaven National Laboratory. "Single atoms break carbon's strongest bond: Single atoms of platinum can break the bond between carbon and fluorine, one of the strongest known chemical bonds." ScienceDaily. ScienceDaily, 2 October 2018. <>.

DOE/Brookhaven National Laboratory. (2018, October 2). Single atoms break carbon's strongest bond: Single atoms of platinum can break the bond between carbon and fluorine, one of the strongest known chemical bonds. ScienceDaily. Retrieved October 14, 2021 from

DOE/Brookhaven National Laboratory. "Single atoms break carbon's strongest bond: Single atoms of platinum can break the bond between carbon and fluorine, one of the strongest known chemical bonds." ScienceDaily. (accessed October 14, 2021).

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Which bond is strongest in a carbon carbon triple bond?

A triple bond involves the sharing of six electrons, with a sigma bond and two [latex]pi[/latex] bonds. The simplest triple-bonded organic compound is acetylene, C2H2. Triple bonds are stronger than double bonds due to the the presence of two [latex]pi[/latex] bonds rather than one.

Click to see full answer.

Simply so, which carbon carbon bond is strongest?

The bond is labeled as "the strongest in organic chemistry," because fluorine forms the strongest single bond to carbon. Carbon–fluorine bonds can have a bond dissociation energy (BDE) of up to 544 kJ/mol. The BDE (strength of the bond) is higher than other carbon–halogen and carbon–hydrogen bonds.

Furthermore, which bond is the strongest which bond is the longest? And finally the single bonds are weaker than the other two. This way, Triple bonds are the shortest. Then comes double bonds. Finally, single bonds are the longest among the three.

Keeping this in consideration, is a carbon carbon bond stronger than a carbon oxygen bond?

For example a carbon oxygen single bond, C-O, is both longer and weaker than a carbon oxygen double bond, C=O.

Can Carbon 3 bonds?

As you can see, there are 3 orbitals outwards at angles of 120 degrees. These orbitals will overlap with each other, so each carbon forms 3 bonds with other carbons to form a hexagonal layer. The carbons form only three bonds because they are sp 2 hybridized (hence the -ene suffix).


Carbon–carbon bond

Covalent bond between two carbon atoms

A carbon–carbon bond is a covalent bond between two carbonatoms.[1] The most common form is the single bond: a bond composed of two electrons, one from each of the two atoms. The carbon–carbon single bond is a sigma bond and is formed between one hybridized orbital from each of the carbon atoms. In ethane, the orbitals are sp3-hybridized orbitals, but single bonds formed between carbon atoms with other hybridizations do occur (e.g. sp2 to sp2). In fact, the carbon atoms in the single bond need not be of the same hybridization. Carbon atoms can also form double bonds in compounds called alkenes or triple bonds in compounds called alkynes. A double bond is formed with an sp2-hybridized orbital and a p-orbital that is not involved in the hybridization. A triple bond is formed with an sp-hybridized orbital and two p-orbitals from each atom. The use of the p-orbitals forms a pi bond.[2]

Chains and branching[edit]

Carbon is one of the few elements that can form long chains of its own atoms, a property called catenation. This coupled with the strength of the carbon–carbon bond gives rise to an enormous number of molecular forms, many of which are important structural elements of life, so carbon compounds have their own field of study: organic chemistry.

Branching is also common in C−C skeletons. Carbon atoms in a molecule are categorized by the number of carbon neighbors they have:

In "structurally complex organic molecules", it is the three-dimensional orientation of the carbon–carbon bonds at quaternary loci which dictates the shape of the molecule.[3] Further, quaternary loci are found in many biologically active small molecules, such as cortisone and morphine.[3]


Carbon–carbon bond-forming reactions are organic reactions in which a new carbon–carbon bond is formed. They are important in the production of many man-made chemicals such as pharmaceuticals and plastics.

Some examples of reactions which form carbon–carbon bonds are aldol reactions, Diels–Alder reactions, the addition of a Grignard reagent to a carbonyl group, a Heck reaction, a Michael reaction and a Wittig reaction.

The directed synthesis of desired three-dimensional structures for tertiary carbons was largely solved during the late 20th century, but the same ability to direct quaternary carbon synthesis did not start to emerge until the first decade of the 21st century.[3]

Bond strengths and lengths[edit]

The carbon-carbon single bond is weaker than C-H, O-H, N-H, H-H, H-Cl, C-F, and many double or triple bonds, and comparable in strength to C-O, Si-O, P-O, and S-H bonds,[4] but is commonly considered as strong.

The values given above represent C-C bond dissociation energies that are commonly encountered; occasionally, outliers may deviate drastically from this range.

Extreme cases[edit]

Long, weak C-C single bonds[edit]

Various extreme cases have been identified where the C-C bond is elongated. In Gomberg's dimer, one C-C bond is rather long at 159.7 picometers. It is this bond that reversibly and readily breaks at room temperature in solution:[6]

Gomberg dimer dissociation.png

In the even more congested molecule hexakis(3,5-di-tert-butylphenyl)ethane, the bond dissociation energy to form the stabilized triarylmethyl radical is only 8 kcal/mol. Also a consequence of its severe steric congestion, hexakis(3,5-di-tert-butylphenyl)ethane has a greatly elongated central bond with a length of 167 pm.[7]

Twisted, weak C-C double bonds[edit]

The structure of tetrakis(dimethylamino)ethylene (TDAE) is highly distorted. The dihedral angle for the two N2C ends is 28º although the C=C distance is normal 135 pm. The nearly isostructural tetraisopropylethylene also has a C=C distance of 135 pm, but its C6 core is planar.[8]

Short, strong C-C triple bonds[edit]

On the opposite extreme, the central carbon–carbon single bond of diacetylene is very strong at 160 kcal/mol, as the single bond joins two carbons of sp hybridization.[9] Carbon–carbon multiple bonds are generally stronger; the double bond of ethylene and triple bond of acetylene have been determined to have bond dissociation energies of 174 and 230 kcal/mol, respectively.[10] A very short triple bond of 115 pm has been observed for the iodonium species [HC≡C–I+Ph][CF3SO3], due to the strongly electron-withdrawing iodonium moiety.[11]

See also[edit]


  1. ^Dembicki, Harry (2016-10-06). Practical Petroleum Geochemistry for Exploration and Production. Elsevier. p. 7. ISBN .
  2. ^Smith, Michael B.; March, Jerry (2007), Advanced Organic Chemistry: Reactions, Mechanisms, and Structure (6th ed.), New York: Wiley-Interscience, ISBN 
  3. ^ abcQuasdorf, Kyle W.; Overman, Larry E. (2014). "Review: Catalytic enantioselective synthesis of quaternary carbon stereocentres". Nature (paper). 516 (7530): 181–191. Bibcode:2014Natur.516..181Q. doi:10.1038/nature14007. PMC 4697831. PMID 25503231.closed access
  4. ^Yu-Ran Luo and Jin-Pei Cheng "Bond Dissociation Energies" in CRC Handbook of Chemistry and Physics, 96th Edition.
  5. ^CRC Handbook of Chemistry and Physics, 88th edition
  6. ^Bochkarev, L. N.; Molosnova, N. E.; Zakharov, L. N.; Fukin, G. K.; Yanovsky, A. I.; Struchkov, Y. T. (1995). "1-Diphenylmethylene-4-(triphenylmethyl)cyclohexa-2,5-diene Benzene Solvate". Acta Crystallographica Section C Crystal Structure Communications. 51 (3): 489–491. doi:10.1107/S0108270194009005.
  7. ^Rösel, Sören; Balestrieri, Ciro; Schreiner, Peter R. (2017). "Sizing the role of London dispersion in the dissociation of all-meta tert-butyl hexaphenylethane". Chemical Science. 8 (1): 405–410. doi:10.1039/c6sc02727j. ISSN 2041-6520. PMC 5365070. PMID 28451185.
  8. ^Bock, Hans; Borrmann, Horst; Havlas, Zdenek; Oberhammer, Heinz; Ruppert, Klaus; Simon, Arndt (1991). "Tetrakis(dimethylamino)ethene: An Extremely Electron-Rich Molecule with Unusual Structure both in the Crystal and in the Gas Phase". Angewandte Chemie International Edition in English. 30 (12): 1678–1681. doi:10.1002/anie.199116781.
  9. ^"NIST Webbook".
  10. ^Blanksby, Stephen J.; Ellison, G. Barney (April 2003). "Bond Dissociation Energies of Organic Molecules". Accounts of Chemical Research. 36 (4): 255–263. CiteSeerX doi:10.1021/ar020230d. ISSN 0001-4842. PMID 12693923.
  11. ^1927-, Streitwieser, Andrew (1992). Introduction to organic chemistry. Heathcock, Clayton H., 1936-, Kosower, Edward M. (4th ed.). Upper Saddle River, N.J.: Prentice Hall. p. 574. ISBN . OCLC 52836313.CS1 maint: numeric names: authors list (link)

Carbon bond carbon strongest

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