General Chemistry Review for Organic Chemistry Part 3

Summary

In this chemistry review segment, tutor Melissa Maribel explains the concept of formal charges in molecules, demonstrating how to calculate them for atoms in both neutral and charged molecules. She emphasizes the role of formal charges in identifying the most stable Lewis structures and discusses the impact of electronegativity on charge distribution. The video also covers resonance, illustrating how delocalized electrons contribute to a molecule's stability through multiple Lewis structures, and introduces hybridization, detailing the process of determining electron geometry and the resulting hybridization types for central atoms.

Takeaways

• 🔍 Formal charges are the charges assigned to individual atoms in a molecule, which help determine the most stable Lewis structure.
• 📚 The formula for calculating formal charge is: Valence electrons - (Non-bonding electrons) - 0.5 * (Bonding electrons).
• 🧪 In a neutral molecule, the sum of all formal charges should equal the overall charge of the molecule.
• 🌐 For the CO2 molecule, both oxygen atoms have a formal charge of zero, indicating stability.
• 🔄 In charged molecules, formal charges help identify the most stable structure by considering electronegativity and the preference for the most electronegative atom to be negatively charged.
• ⚛️ Resonance occurs when a Lewis structure can be drawn in multiple ways, indicating the delocalization of electrons.
• 🔁 The double-headed arrow is used to represent resonance, showing the movement of electrons between different structures.
• 🌟 The resonance hybrid is the complete Lewis structure that combines all resonance structures, providing a more accurate representation of the molecule's electron distribution.
• 💡 The more resonance structures a molecule has, the more stable it is, as it distributes electrons more evenly.
• 📐 Hybridization is determined by the electron geometry around a central atom, which is found by counting electron groups (bonds + lone pairs).
• 🌀 Different electron geometries correspond to different types of hybridization, such as sp, sp2, and sp3, each with distinct properties and bond angles.

Q & A

What are formal charges and why are they important in chemistry?

Formal charges are the charges assigned to each individual atom within a molecule, based on the Lewis structure. They are important because they help determine the most stable Lewis structure of a molecule and ensure that the sum of the formal charges equals the overall charge of the molecule.

How is the formal charge of an atom calculated?

The formal charge of an atom is calculated using the formula: Formal Charge = (Valence Electrons of the atom) - (Non-bonding electrons/2) - (Bonding electrons/2). This formula helps in assessing the distribution of electrons around an atom.

What is the typical valence electron count for oxygen and how does it relate to formal charges?

Oxygen typically has 6 valence electrons. When calculating formal charges, these are compared with the electrons directly involved in bonding and the lone pairs to determine the charge on the oxygen atom in a molecule.

How does the formal charge formula apply to the oxygen atom in CO2?

For the oxygen atom in CO2, the formula is applied by considering its 6 valence electrons, 2 bonding electrons, and 4 lone electrons. The calculation results in a formal charge of 0 for oxygen in CO2.

What is the significance of the formal charge being zero for a neutral molecule?

A formal charge of zero for all atoms in a neutral molecule suggests that the molecule is as stable as possible, with an even distribution of electrons, which is typically preferred in stable molecular structures.

What is the difference between a neutral and a charged molecule in terms of formal charges?

In a neutral molecule, the sum of the formal charges is zero, indicating an even distribution of electrons. In contrast, a charged molecule has a non-zero sum of formal charges, indicating an uneven distribution of electrons and the presence of a net charge.

Why is the most electronegative atom typically assigned a negative formal charge?

The most electronegative atom is typically assigned a negative formal charge because it has a greater tendency to attract electrons towards itself, thus pulling electron density away from other atoms in the molecule.

What is resonance and how does it relate to the stability of a molecule?

Resonance occurs when a molecule can be represented by multiple Lewis structures, indicating the delocalization of electrons. The more resonance structures a molecule has, the more stable it is, as it represents a distribution of electron density over multiple atoms.

How does the concept of hybridization relate to the electron geometry of an atom?

Hybridization is the process by which atomic orbitals combine to form new hybrid orbitals suitable for bonding. The electron geometry around an atom, determined by the number of electron groups (bonds plus lone pairs), dictates the type of hybridization (e.g., sp, sp2, sp3).

What is the hybridization of carbon in a molecule with four electron groups?

When a carbon atom has four electron groups, it adopts a tetrahedral electron geometry, which corresponds to sp3 hybridization.

How does the concept of resonance hybrid help in understanding the actual structure of a molecule?

A resonance hybrid is a combination of all possible resonance structures, representing the average distribution of electrons in a molecule. It provides a more accurate picture of the molecule's structure by accounting for the delocalization of electrons across multiple atoms.

Outlines

00:00

🔍 Understanding Formal Charges in Organic Chemistry

This paragraph introduces the concept of formal charges in organic chemistry, which are the charges assigned to individual atoms within a molecule to determine the most stable Lewis structure. The tutorial explains how to calculate formal charges for a neutral molecule like CO2 and a charged molecule, using the example of a carbonate ion. It emphasizes the importance of electronegativity in determining the stability of a structure, with more electronegative atoms preferring to have a negative charge. The paragraph also touches on the concept of resonance, where multiple Lewis structures can be drawn for a molecule, and how resonance structures contribute to the molecule's stability through electron delocalization.

05:04

🌐 Exploring Resonance and Hybridization in Organic Molecules

The second paragraph delves deeper into the concept of resonance, illustrating how it is represented by the movement of electrons and the breaking and reforming of double bonds in different structures. It describes the process of combining all resonance structures to form a resonance hybrid, which is a complete Lewis structure that accounts for the delocalization of electrons. The paragraph also covers the topic of hybridization, explaining how to determine the hybridization of central atoms in various molecular structures by identifying electron groups, which include bonds and lone pairs. Examples are provided to demonstrate how different electron geometries, such as tetrahedral, trigonal planar, and linear, correspond to different types of hybridization (sp3, sp2, and sp, respectively).

Keywords

💡Formal charges

Formal charges refer to the charge assigned to each atom in a molecule based on the Lewis structure. It is calculated by taking the number of valence electrons of the atom, adding the non-bonding electrons, and subtracting the bonding electrons. In the video, formal charges are used to determine the most stable Lewis structure for both neutral and charged molecules, such as CO2 and a charged molecule involving nitrogen and oxygen. The concept is crucial for understanding molecular stability and structure.

💡Lewis structure

A Lewis structure is a graphical representation of a molecule's electron distribution, showing how the valence electrons are arranged among the atoms. It is used to understand the bonding and overall charge of a molecule. In the script, the Lewis structure is central to the discussion of formal charges, resonance, and hybridization, as it provides the framework for these concepts.

💡Valence electrons

Valence electrons are the outermost electrons of an atom that are involved in chemical bonding. They are essential in determining the reactivity of an element and its ability to form compounds. In the video, the number of valence electrons is used to calculate formal charges, as seen in the examples of oxygen and carbon in CO2.

💡Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. It plays a significant role in determining the polarity of a bond and the distribution of electrons in a molecule. The script mentions that the most electronegative atoms prefer to be negative, which is a key factor in evaluating the stability of Lewis structures.

💡Resonance

Resonance occurs when a molecule can be represented by two or more Lewis structures, none of which alone can fully describe the bonding situation. It is a way to describe the delocalization of electrons in a molecule. The video script explains resonance with the example of the carbonate ion, where the double bond can move between different oxygen atoms, leading to multiple resonance structures.

💡Delocalization

Delocalization refers to the distribution of electrons over a larger area than would be expected if they were localized in specific bonds. It is a key concept in resonance, where electrons are not confined to a single bond but are spread out over several atoms. The script uses delocalization to explain the movement of electrons in resonance structures.

💡Resonance hybrid

A resonance hybrid is a combination of all possible resonance structures of a molecule, which together provide a more accurate representation of the molecule's electron distribution. It is a concept that emerges from the idea of resonance, as the video script illustrates with the example of the carbonate ion, combining the different resonance structures to form the complete Lewis structure.

💡Hybridization

Hybridization is the concept where atomic orbitals combine to form new hybrid orbitals that are suitable for bonding in molecules. It helps to explain the geometry of molecules. In the script, hybridization is used to determine the electron geometry and the shape of molecules, such as sp3, sp2, and sp hybridizations for carbon in different structures.

💡Electron geometry

Electron geometry is the three-dimensional arrangement of the electron groups surrounding a central atom in a molecule. It is determined by the number of electron groups, which include both bonding pairs and lone pairs. The script discusses electron geometry in the context of hybridization, using it to classify the shape of molecules and their central atoms.

💡Bonding and lone pairs

In the context of Lewis structures, bonding pairs are electrons shared between two atoms, while lone pairs are electrons that belong to a single atom. The distinction between bonding and lone pairs is crucial for calculating formal charges and determining the electron geometry. The script uses these concepts to explain how to calculate formal charges and to identify the electron groups for hybridization.

💡Stability

Stability in chemistry refers to the tendency of a molecule to resist change or decomposition. In the video, the stability of a molecule is related to its formal charges, resonance structures, and hybridization. A more stable molecule will have a lower energy state, and the script discusses how formal charges and resonance can indicate a molecule's stability.

Highlights

Introduction to formal charges in organic chemistry and their importance in determining the most stable Lewis structure.

Explanation of formal charges for a neutral CO2 molecule, demonstrating the calculation for oxygen and carbon atoms.

Formal charge calculation formula application for oxygen in CO2, resulting in a zero formal charge.

Demonstration that both oxygen atoms in CO2 have the same formal charge due to identical bonding environments.

Finding the formal charge of carbon in CO2, which also results in zero, indicating a preference for neutral formal charges in stable molecules.

Transition to charged molecules and the use of formal charges to determine the most stable structure among possible configurations.

Analysis of a charged molecule's formal charges, including nitrogen with a formal charge of -1, and the preference for electronegative atoms to be negative.

Discussion on the electronegativity trend and its impact on the stability of molecular structures, with oxygen being more electronegative than nitrogen.

Introduction to resonance in Lewis structures and its role in determining the most stable configuration.

Illustration of delocalization of electrons in resonance structures and the concept of resonance hybrid.

Explanation of how combining resonance structures results in a more stable resonance hybrid, similar to combining personality traits.

The significance of the number of resonance structures in determining molecular stability.

Introduction to the concept of hybridization and its calculation for central atoms in molecular structures.

Determination of hybridization for carbon in a structure with four electron groups, resulting in sp3 hybridization.

Calculation of hybridization for carbon in a trigonal planar structure with sp2 hybridization due to three electron groups.

Analysis of hybridization for carbon in a linear structure with sp hybridization, considering the triple bond as one electron group.

Encouragement for viewers to stay determined and continue learning organic chemistry, with links to additional resources.

Transcripts

Welcome to part three of the general

chemistry review for organic chemistry.

I'm Melissa Maribel your personal tutor

and let's go over formal charges.

Formal charges are the charges of each

individual atom within the molecule

all the formal charges add up to the overall

charge of the molecule. Formal charges

are a great way to check if we drew the

most stable lewis structure there are

two different cases case one is a

neutral molecule and case two is a

charged molecule here's a neutral

molecule of CO2 and here's the Lewis

structure. Let's practice finding formal

charges for each atom. Here's our formal charge formula.

we'll apply this formula for the oxygen on the left.

Start with the typical valence electrons

for oxygen which is 6

next are the bonding electrons this

means the electrons that are found in

the bonds that are directly touching

that oxygen so we have one, two bonding

electrons. Now count each individual lone

pair or really each electron on that

oxygen. So we have one, two, three, four

remember order of operations or PEMDAS, we will start with parentheses first and

then subtract so our formal charge is

zero. The oxygen on the right would have the same formal charge since it is

exactly the same as the one on the left.

Let's find carbon's formal charge. Carbon has four valence electrons, the electrons

directly touching carbon are 1, 2, 3 & 4

and there are no individual lone pairs so 4 minus 4 gives us 0. A neutral

molecule prefers to have all formal

charges be 0 if it's possible for the structure. Case 2, we have a charged

molecule here are all the possible structures we have. Let's use formal

charges to help us see which one is the

most stable, we'll find the formal charges

for each atom in the structure. Oxygen has 6 valence electrons there are two

bonding electrons and four individual

lone pairs so we will get 0 for our formal charge.

Carbon has four valence electrons there are four bonding electrons and no

individual lone pairs. Carbon's formal

charge is also zero. Nitrogen has five valence electrons there are two bonding

electrons and for individual lone pairs. Nitrogen's formal charge is negative one.

Note: the most electronegative atoms want to be negative and looking back at our

electronegativity trend oxygen is more electronegative than nitrogen so this

structure is not the most stable since oxygen should be negative here are the

formal charges for the next structure. Once again oxygen prefers to be negative

and this one is out since the formal charge is a plus one this last structure

is the most stable due to the most electronegative atom of oxygen being

negative and you will also see that the central atom prefers to have a formal

charge of zero. Next is resonance when Lewis structures can be drawn multiple

ways this is known as resonance. Here's the first lewis structure for a

carbonate ion. To find the next structure we can move the double bond to a

different oxygen what we are doing is moving the electrons to this oxygen and

our double bond is breaking and reforming to a different oxygen

resonance is represented by a double arrow there is one more resonance

structure so we'll move the electrons to this other oxygen and break the double

bond here and reform it on this oxygen. These are all of our resonance

structures. This movement of electrons that we just saw is called

delocalization. In organic chemistry we're going to get a bit deeper into the

concept of resonance it is going to be described as a structure that has the

delocalized electrons which is referring to the ability of moving electrons. If we

were to combine all the resonance structures it would give us a resonance

hybrid think of it this way every person has different personality traits let's

say and in my case I'm academic, dorky and persistent combine all those traits

together and you get me. Same goes for

our carbonate ion by combining all the resonance structures or personality

traits this gives you the complete lewis structure which is known as the

resonance hybrid. The more resonance structures a molecule has the more

stable it typically is. Now for hybridization let's find the

hybridization of each central atom for this structure looking at this first

carbon we will identify the electron geometry to find the electron geometry

we must find the electron groups of this carbon electron groups are bonds plus

lone pairs. We have one, two, three, four bonds and no lone pairs so we have four

electron groups. Using our table this is tetrahedral and tetrahedral has a

hybridization of sp3 so each carbon has a hybridization of sp3. Let's find the

hybridization of this structure we'll find the electron groups and note a

double or triple bond only counts as one bond whenever we are finding the

electron groups so this carbon has one, two, three electron groups making it

trigonal planar for our electron geometry and our hybridization for each

carbon is then sp2. Let's find the hybridization for each carbon in this

structure we'll find the electron groups first and the triple bond only counts as

one bond so we have one, two electron groups making it linear so our

hybridization for each carbon is sp.

If you missed part 1 or part 2 you can find

that right over here and if you're ready for an introduction to organic chemistry

you can also find that right over here

and remember stay determined you can do this!