Acetate (CH3COO-) ion Lewis dot structure, molecular geometry or shape, electron geometry, bond angles, hybridization, charges, polar vs nonpolar
CH3COO– represents the acetate anion, a polyatomic ion, the conjugate base of acetic acid (CH3COOH). Acetate salts are widely used in various biochemical processes.
For instance, acetate buffers efficiently maintain a stable pH. Acetate films are commonly used in developing negatives for film photography.
In this article, you will learn how to draw the Lewis dot structure of CH3COO–, what is its molecular geometry or shape, electron geometry, bond angles, hybridization, resonance structures, formal charges and polarity.
The Lewis structure of CH3COO– comprises a C-atom at the center. It is single covalently bonded to an O-atom on one side and a methyl (CH3) group on the other side. The central C-atom contains another O-atom double-covalently bonded to it at the third end.
There is no lone pair of electrons on the central C-atom however, the terminal O-atoms contain lone pairs respectively.
If you want to draw the Lewis dot structure of the acetate (CH3COO–) ion, then follow the simple steps given below, one by one.
Steps for drawing the Lewis dot structure of CH3COO–
1. Count the total valence electrons present in CH3COO–
The three distinct elements present in CH3COO– are carbon, hydrogen, and oxygen.
Carbon (C) is located in Group IVA (or 14) of the Periodic Table of Elements, containing a total of 4 valence electrons in each atom.
In contrast, oxygen (O) is present in Group VI A (or 16) of the Periodic Table. Thus, it has a total of 6 valence electrons in each oxygen atom.
However, hydrogen (H) lies at the top of the Periodic Table, having a single valence electron only.
Total number of valence electrons in hydrogen = 1
Total number of valence electrons in carbon = 4
Total number of valence electrons in oxygen = 6
The acetate (CH3COO–) anion comprises 2 C-atoms, 2 O-atoms and 3 H-atoms.
An important point to remember is that CH3COO– carries a negative one (-1) charge, which means 1 extra valence electron is added to its Lewis structure.
∴ Therefore, the total valence electrons available for drawing the Lewis dot structure of CH3COO– = 2(4) + 2(6) + 3(1) = 23 + 1 = 24 valence electrons.
2. Find the least electronegative atom and place it at the center
By convention, the least electronegative atom out of all those available is chosen as the central atom while drawing the Lewis structure of a molecule or molecular ion.
The least electronegative atom can easily form covalent bonds with other atoms by sharing its electrons.
Hydrogen (E.N = 2.20) is less electronegative than both carbon (E.N = 2.55) and oxygen (E.N = 3.44).
However, the H-atom is an exception as it cannot be chosen as the central atom in any Lewis structure. It can accommodate a total of 2 valence electrons, forming a single covalent bond with 1 adjacent atom only.
Therefore, we select the second option, i.e., a C-atom as the central atom in CH3COO– Lewis structure.
Any of the two C-atoms can be considered a central atom. In this article, we select Cb as the central atom. So it is placed in the middle while the other C-atom and 2 O-atoms are spread around it, as shown below.
If you are wondering where to place the remaining atoms, i.e., 3 H-atoms, then these are placed next to Ca. This ensures that the carboxylate (COO–) functional group stays intact in the acetate ion (salt of carboxylic acid).
A functional group in organic chemistry is a collection of atoms within molecules that bind together to react in predictable ways.
Therefore, in the above structure, the COO– group atoms are placed together, whereas hydrogen goes outside to form the methyl (CH3) functional group.
3. Connect the outer atoms with the central atom
In this step, the outer atoms are joined to the central atom using single straight lines.
Therefore 2 O-atoms and 1 C-atom are directly joined to Cb. In contrast, 3 H-atoms are joined to Ca, as shown below.
The H-atom can form a single bond with its adjacent atom only.
A straight line represents a single covalent bond, i.e., a bond pair containing 2 electrons.
In the above structure, there are a total of 6 single bonds, i.e., 6(2) = 12 valence electrons are already consumed out of the 24 initially available.
Now let’s see in the next steps where to place the remaining 12 valence electrons in the CH3COO– Lewis dot structure.
4. Complete the octet and/or duplet of the outer atoms
An O-atom needs a total of 8 valence electrons in order to achieve a stable octet electronic configuration.
A C-O bond represents 2 valence electrons already present around each O-atom in the CH3COO– Lewis structure drawn till this step.
Therefore, the remaining 6 valence electrons are placed as 3 lone pairs around each O-atom to complete its octet, as shown below.
In contrast, Ca already has a complete octet possessing a total of 4 single bonds (3 C-H and 1 C-C), i.e., 8 valence electrons surrounding it. Also, all three H-atoms have a complete duplet configuration required for their stability.
So no changes are required in the above Lewis structure with regard to these outer atoms.
5. Complete the octet of the central atom and convert lone pairs into covalent bonds if necessary
Total valence electrons used till step 4 = 6 single bonds + 2(electrons placed around each O-atom, shown as dots) = 6(2) + 2(6) = 24 valence electrons.
Total valence electrons – electrons used till step 4 = 24 – 24 = 0 valence electrons.
As all the valence electrons initially available for drawing the CH3COO– Lewis structure are already consumed so there is no lone pair on the central C-atom.
However, the central C-atom (Cb) still has an incomplete octet with a total of 3 single bonds, i.e., 3(2) = 6 valence electrons surrounding it.
Therefore, a lone pair from one of the two terminal O-atoms is converted into an additional covalent bond between the central C-atom and this O-atom, as shown below.
Now the central C-atom also has a complete octet.
So let’s move ahead and check the stability of the Lewis structure drawn above by applying the formal charge concept.
6. Check the stability of Lewis’s structure using the formal charge concept
The less the formal charge on the atoms of a molecule, the better the stability of its Lewis structure.
The formal charges can be calculated using the formula given below.
Formal charge = [valence electrons- nonbonding electrons- ½ (bonding electrons)].
Now let us use this formula and the Lewis structure obtained in step 5 to determine the formal charges present on the CH3COO– bonded atoms.
For COO– bonded carbon atom
Valence electrons of carbon = 4
Bonding electrons = 2 single bonds + 1 double bond = 2(2) + 4 = 8 electrons
Non-bonding electrons = no lone pair = 0 electrons
Non-bonding electrons = no lone pairs = 0 electrons
Formal charge = 1-0-2/2 = 1-0-1= 1-1 = 0
As per the above calculation, none of the atoms carry a formal charge in the CH3COO– Lewis structure except the C-O single-bonded oxygen atom that carries a formal charge of -1. Therefore, the overall charge present on the acetate anion is -1.
This ensures that we have drawn the acetate ion Lewis structure correctly.
However, you may also note that the following resonance structures are possible for representing the acetate ion.
A C=O double bond can be formed using any of the two oxygen atoms available. Thus, both the above resonance structures are equivalent. The actual structure of the acetate ion is a weighted average of the above two, and it is known as the resonance hybrid.
The continuous movement of lone pairs and pi-bonded electrons from one position to another on a molecule or molecular ion is known as electronic delocalization.
Now that we know everything about the CH3COO– Lewis structure let us move ahead and discuss its molecular and electron geometry.
What are the electron and molecular geometry of CH3COO-?
The molecular geometry or shape of the acetate (CH3COO–) ion w.r.t the COO– bonded carbon atom is trigonal planar while that w.r.t CH3 bonded carbon, it is tetrahedral in shape. In either case, the molecular geometry of CH3COO– w.r.t both the carbon atoms is identical to its ideal electron pair geometry.
The absence of any lone pair of electrons on the C-atom results in the molecular ion having negligible electronic repulsion and, thus, no molecular distortion.
Molecular geometry of CH3COO–
The molecular geometry or shape of the acetate (CH3COO–) ion w.r.t the central C-atom (Cb) is trigonal planar.
The central C-atom is directly bonded to three other atoms. There is no lone pair of electrons on the central C-atom thus, no lone pair-lone pair or lone pair-bond pair electronic repulsions exist in the molecular shape.
In contrast, the shape of the acetate ion w.r.t the methyl-bonded carbon (Ca) is tetrahedral. To a C-atom at the center, three H-atoms and another C-atom are directly attached, like four corners of a tetrahedron, as shown below.
Electron geometry of CH3COO–
The electron geometry of a molecule or molecular ion can be determined by applying the AXN concept of the valence shell electron pair repulsion (VSEPR) theory of chemical bonding.
AXN is a simple formula representing the number of bonded atoms and lone pairs present on the central atom.
The C-atom at the center is surrounded by 3 bond pairs, and it has no lone pair of electrons, making a total of 3 electron density regions. Hence, the ideal electronic geometry of CH3COO– w.r.t Cb is trigonal planar.
The C-atom at the center is surrounded by 4 bond pairs, and it has no lone pair of electrons, making a total of 4 electron density regions. Hence, the ideal electron pair geometry of CH3COO– w.r.t Ca is tetrahedral.
You may confirm the above-mentioned information using the VSEPR chart given below:
Hybridization of CH3COO–
A shortcut to finding the hybridization present in a molecule or molecular ion is using its steric number against the table given below.
The steric number of Ca in CH3COO– is 4 while that of Cb is 3, so the two carbon atoms are sp3 hybridized and sp2 hybridized, respectively.
Steric number
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
Ca uses its four sp3 hybrid orbitals to form C-C and C-H sigma bonds by sp3-sp2 and sp3-s orbital overlap, respectively.
Conversely, Cb uses its three sp2 hybrid orbitals to form the C-C, C-O and C=O sigma bonds in CH3COO–.
However, the unhybridized p-orbital of Cb goes in the formation of the C=O pi bond by p-p overlap.
Refer to the figure drawn below.
The bond angles of CH3COO–
The ideal bond angle in a trigonal planar molecular shape is 120°, while that in a tetrahedral molecule is 109.5°.
Therefore, in the acetate (CH3COO–) ion, each H-C-H bond angle is equal to 109.5° while the –O-C=O bond angle equals 120°.
As per Pauling’s electronegativity scale, a polar covalent bond is formed between two dissimilar atoms having an electronegativity difference between 0.4 to 1.6 units.
The three main types of bonds present in CH3COO– are C-C, C-H and C-O (or C=O) bonds.
The C-C bond is purely non-polar.
The C-H bonds are very weakly polar (almost non-polar as per Pauling’s electronegativity scale), having a small electronegativity difference of 0.35 units between the bonded atoms.
Contrarily, the C-O bonds are strongly polar as per an electronegativity difference of 0.89 units between the bonded atoms.
The two strongly electronegative O-atoms attract the shared electron cloud of the acetate ion to a great extent. This leads to an overall polar acetate anion (net dipole moment µ > 0).
What is the Lewis structure for acetate (CH3COO–) ion?
The acetate (CH3COO–) ion displays a total of 24 valence electrons, i.e., 24/2 = 12 electron pairs.
Out of the 12 electron pairs, there are 7 bond pairs and 5 lone pairs of electrons.
A C-atom at the center is bonded to another C-atom, 2 O-atoms and a CH3-group in CH3COO– Lewis structure.
There is no lone pair of electrons on the central C-atom, while 2 and 3 lone pairs are present respectively on C=O and C-O bonded oxygen atoms.
How many lone pairs does C2H3O2– have?
C2H3O2– represents the condensed molecular formula for the acetate ion (CH3COO–). There are a total of 5 lone pairs in its structure.
These include 2 lone pairs of electrons on the C=O double-bonded oxygen atom, while 3 lone pairs are present on C-O single-bonded oxygen atom.
What are the formal charges present on each of the bonded atoms in CH3COO–?
Zero or no formal charges are present on either of the bonded atoms in CH3COO– except the C-O single bonded oxygen atom that carries a formal charge of -1.
Thus the overall charge present on the acetate anion is also -1.
What is the VSEPR shape of CH3COO–?
The VSEPR shape of CH3COO–w.r.tCH3 bonded C-atom is tetrahedral while that w.r.t COO–bonded C-atom is trigonal planar.
Is the molecular shape of CH3COO– identical to its ideal electron geometry?
Yes. The absence of any lone pair of electrons on either of the C-atoms in CH3COO– implies that no lone pair-lone pair or lone pair-bond pair electronic repulsions exist in the molecular shape.
So no distortion is seen in its shape and/or geometry.
How is the shape of acetate (CH3COO–) anion different from its conjugate acid, acetic acid (CH3COOH)?
The molecular shape of CH3COO– is similar to that of its conjugate acid (CH3COOH) w.r.t both C-atoms, i.e., tetrahedral w.r.t CH3 bonded carbon atom and trigonal planar w.r.t COO– bonded carbon atom.
However, there is an additional possibility in CH3COOH, i.e., the shape of acetic acid w.r.t OH bonded oxygen atom is bent, angular or V-shaped.
How are the shapes of acetone (CH3COCH3) and acetate anion (CH3COO–) the same or different?
The molecular shape of acetone (CH3COCH3) w.r.t the central C-atom is trigonal planar. To a C-atom at the center, 1 O-atom and two methyl (CH3) groups are directly attached. It has no lone pair of electrons.
Likewise, the shapes of both CH3COCH3 and CH3COO–w.r.t methyl-bonded C-atoms are tetrahedral. Thus both CH3COCH3 and CH3COO– possess similar shapes.
The total number of valence electrons available for drawing the acetate (CH3COO–) ion Lewis structure is 24.
The molecular geometry or shape of CH3COO– w.r.t CH3 bonded C-atom is tetrahedral while that w.r.t COO– bonded C-atom is trigonal planar.
The molecular geometry of CH3COO– is identical to its ideal electronic geometry.
The two C-atoms are sp3 hybridized and sp2 hybridized respectively in CH3COO–.
The acetate ion is overall polar as it contains strongly electronegative O-atoms that lead to an asymmetrically spread electron cloud (net µ > 0).
The C-O single bonded oxygen atom carries a -1 formal charge in CH3COO– which is the charge present on the acetate ion overall, thus ensuring that we have drawn its Lewis structure correctly in this article.
Vishal Goyal is the founder of Topblogtenz, a comprehensive resource for students seeking guidance and support in their chemistry studies. He holds a degree in B.Tech (Chemical Engineering) and has four years of experience as a chemistry tutor. The team at Topblogtenz includes experts like experienced researchers, professors, and educators, with the goal of making complex subjects like chemistry accessible and understandable for all. A passion for sharing knowledge and a love for chemistry and science drives the team behind the website. Let's connect through LinkedIn: https://www.linkedin.com/in/vishal-goyal-2926a122b/
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