Many important compounds in the chemical industry are derivatives of ethylene (C2 H4). Two of th

step 1 Okay, so let me just write out the structure first. You have C double bond C and H, another H he

Many important compounds in the chemical industry are...

Key Concepts

Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms of a molecule and the lone pairs of electrons that may exist. They are essential for understanding how valence electrons are arranged in bonds and non-bonding pairs, allowing us to predict molecular reactivity, geometry, and the presence of multiple bonds.

Hybridization

Hybridization is the concept of mixing atomic orbitals to form new orbitals that are suitable for the pairing of electrons to form chemical bonds in molecules. It explains the observed molecular geometries and bond angles in organic compounds, such as why carbon atoms can adopt sp, sp2, or sp3 configurations, which directly influences molecular shape and reactivity.

Molecular Geometry and Bond Angles

Molecular geometry involves the three-dimensional arrangement of atoms in a molecule, which is determined by the hybridization of the central atoms and the repulsion between electron pairs (bonding and lone pairs). Bond angles are the angles between adjacent bonds, and they provide insight into the structural layout of the molecule. Understanding these helps in predicting molecular behavior and interactions.

Sigma and Pi Bonds

Sigma (?) bonds are the first bonds formed between two atoms, featuring head-on orbital overlap, while pi (?) bonds result from the side-on overlap of unhybridized orbitals when a second bond is present between two atoms. This distinction is critical for understanding the strength, reactivity, and electronic structure of molecules, particularly in systems featuring double or triple bonds.

Molecular Planarity

Molecular planarity refers to the arrangement of atoms in a single flat plane. In many organic molecules, planarity is influenced by the hybridization of the atoms, particularly sp2 hybridization, and by the delocalization of electrons in conjugated systems. This concept is important when considering the chemical and physical properties, such as reactivity and optical characteristics, of the molecule.

Transcript

00:02 Okay, so let me just write out the structure first.

00:07 You have c double bond c and h, another h here, c, triple bond, n.

00:15 So the way to see if your lewis structure is complete is to check whether each atom within the molecule has a completed duplet or octet configuration.

00:26 Let's go with the blue here.

00:29 Hydrogen has two electrons over here, so it has completed its duplicate configuration.

00:37 So has this hydrogen.

00:39 So has this hydrogen.

00:40 Now the carbon here has 1, 2, 3, 4, 5, 6, 7, 8 electrons around it.

00:50 So it has completed its octet configuration.

00:53 And so has the carbon next to it because it has 4 electrons from the double bonds.

00:59 Two electrons from the hydrogen so that adds up to six electrons and then two electrons from the single bond that it shares with the carbon next to it now let's check the carbon that's triple bonded to nitrogen it has one two three four five six electrons from the triple bond two electrons from the single bond with the other carbon so it's completed with its octet configuration and the nitrogen has six electrons, but it's lacking its octet configuration.

01:34 So to complete its octet configuration, we have to add two more electrons, which will constitute a lone pair on the nitrogen that will complete its octet configuration and hence the fluid structure.

01:47 So that is step one done.

01:51 Step two asks you to label the, the angles a, b, and c, which is this one right here.

02:11 So a is 120 degrees because the carbon is sb2 hybridized.

02:20 B is also 120 degrees because the carbon is sp2 hybridized.

02:27 And c is 180 degrees because the carbon here is sp hybridized.

02:34 The third step is writing down the hybridization of each carbon, and this is what we've done when we have been labeling the angles.

02:49 And the last question that we have is, are all these atoms on the same plane or different planes? and the answer is all the atoms are on the same plane.

03:08 And let us delve into that a little bit.

03:11 So these two carbons have to be on the same plane because they are both sp2 trigonal planner atoms or systems.

03:24 And since this carbon is attached to this carbon, they are in the same plane.

03:29 And this carbon here is sp2 hybridized, sorry, sp hybridized, which means it shares a triple bond, which means also that it has to be in the same plane as the nitrogen.

03:42 And so the nitrogen is in the same plane with the carbon, which inherently makes it in the same plane with this original carbon that we started out with.

03:53 So all these atoms will have to be in the same plane.

04:01 Okay, let's go on to our next molecule here...

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