🧪AP Chemistry
Verified for the 2025 AP Chemistry exam•5 min read•Last Updated on June 18, 2024
Lewis structures can determine properties such as geometry, bond orders, bond lengths, and dipoles for molecules. The Valence-Shell-Electron-Pair-Repulsion (VSEPR) theory can predict molecular geometry by minimizing electron-electron repulsion. It specifically uses the Coulombic repulsion between electrons as a basis for predicting electron arrangement.
You should definitely memorize the table below for the AP Exam. It gives you everything you need to know about VSEPR and will answer a lot of questions that require memorization on the AP. Once you practice, the questions that involve the VSEPR Theory become free points🥳!
Let's go over what each column means:
Sigma (σ) bonds are covalent bonds where electrons are found shared on the internuclear axis. Hybrid orbitals form σ bonds, and they are stronger than π bonds.
Pi (π) bonds are covalent bonds where orbitals are perpendicular📐 to the internuclear axis. Unhybridized orbitals form π bonds.
You don't really have to know these definitions, but be aware of the following:
Count the number of σ bonds and the number of π bonds in the following two structures:
In the molecule on the left, there is 1 triple bond and 2 single bonds. 1 triple bond is made up of 1 σ bond and 2 π bonds, while the single bond is only made up of 1 σ bond. Therefore, in total, there are 3 σ bonds and 2 π bonds in this molecule.
In the molecule on the right, there are 3 double bonds and 9 single bonds. This means this molecule is made up of 12 σ bonds and 3 π bonds.
Hybridization is the idea that atomic orbitals fuse to form newly hybridized orbitals, which in turn, influences molecular geometry and bonding properties. Hybridization is also an expansion of the valence bond theory💥.
There are 5 main hybridizations, 3 of which you'll be tested on: sp3, sp2, sp, sp3d, sp3d2. For these hybridizations, electron orbitals fuse together to fill subshells and go to a lower energy state. It also allowed for things like CH4, since technically the way the electron pairs are organized, 4 sigma bonds would not be possible.
In the above example, carbon's 2p and 2s orbitals fuse into 4 half-filled sp3 orbitals that can make 4 sp3-orbital sigma bonds. The same principle applies to the other hybridizations.
The following questions are from past AP Chemistry exams that were posted online by College Board.
(a) In the box provided, draw a complete Lewis electron-dot diagram for the IF3 molecule.
(b) On the basis of the Lewis electron-dot diagram that you drew in part (a), predict the molecular geometry of the IF3 molecule.
(c) In the SO2 molecule, both of the bonds between sulfur and oxygen have the same length. Explain this observation, supporting your explanation by drawing in the box below a Lewis electron-dot diagram (or diagrams) for the SO2 molecule.
(d) On the basis of your Lewis electron-dot diagram(s) in part (c), identify the hybridization of the sulfur atom in the SO2 molecule.
Use the information in the table below to respond to the statements and questions that follow. Your answers should be in terms of principles of molecular structure.
(a) Draw the complete Lewis electron-dot diagram for ethyne in the appropriate cell in the table above.
(b) Which of the four molecules contains the shortest carbon-to-carbon bond? Explain.
(c) A Lewis electron-dot diagram of a molecule of ethanoic acid is given below. The carbon atoms in the molecule are labeled x and y, respectively.
Identify the geometry of the arrangement of atoms bonded to each of the following.
(i) Carbon x
(ii) Carbon y
For part a, one point is earned for a correct Lewis diagram, such as the one below. This example is done with solely dots, but you can also represent the bonds with lines.
The correct answer for part b is "T-shaped." One point is earned for the correct molecular geometry of the diagram drawn in part a. If you drew the Lewis diagram incorrectly in part a, you may still earn credit in part b as long as your answer is consistent with your drawing.
Two points can be earned in part c. One point is earned with the correct Lewis structure and the other is earned for stating that both sulfur-oxygen bonds are double bonds.
One point is earned in part d for listing the correct hybridization: sp2. If you drew the Lewis structure incorrectly in part c, but your answer in part d was consistent with it, you would get credit in this part.
One point is earned in part a for the correct Lewis structure drawn in the table.
Two points can be earned in part b: one for the correct choice and one for the correct explanation. The following is a sample response that would earn both points: "Ethyne, which contains a triple bond, has the shortest C-to-C bond. The other molecules have single C-to-C bonds, and triple bonds are shorter than single bonds."
Two points can also be earned in part c, one for naming the correct molecular geometry for carbon x and the other for carbon y:
Bond energy is the amount of energy required to break a chemical bond and form neutral isolated atoms.
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Bond energy is the amount of energy required to break a chemical bond and form neutral isolated atoms.
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Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule.
Valence Electrons: These are electrons in an atom's outer shell that can participate in forming chemical bonds with other atoms - just like walls and beams in our house blueprint analogy.
Covalent Bonding: This type of bond happens when two or more elements share electron pairs - similar to sharing living space within our house blueprint analogy.
Resonance Structures: These are several Lewis structures that represent a single ion or molecule but differ only by position of electrons – think different design layouts for same square footage in our house blueprint analogy.
Molecular geometry refers to the three-dimensional arrangement or shape formed by atoms within a molecule.
VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts the shape of individual molecules based on the concept that electron pairs surrounding an atom tend to repel each other.
Bond Angle: The angle formed between three atoms across at least two bonds.
Tetrahedral Geometry: A type of molecular geometry where a central atom is bonded to four other atoms or groups, forming a shape similar to a pyramid with triangular faces.
Coulombic repulsion is the repulsive force between two positive or two negative charges.
Coulomb's Law: This law describes how the force between two charged objects depends on their charge and distance apart, just like how strongly or weakly our magnets might push against each other.
Ionic Bonding: This type of bond forms when one atom gives up one or more electrons to another atom - think of it as a tug-of-war where one side lets go of the rope (electron).
Electrostatic Attraction: This is the attraction between positive and negative charges, like opposite poles of a magnet being drawn together.
Hybridization is a concept in molecular chemistry that describes the combination of atomic orbitals within an atom to form new hybrid orbitals. These new orbitals have different shapes, energies, and orientations than the original atomic orbitals.
Orbital: An orbital is a region around an atom where there's high probability to find an electron.
sp3 Hybridization: This occurs when one s orbital mixes with three p orbitals forming four equivalent sp3 hybridized orbitals used in methane formation.
Sigma Bonds (σ): These are covalent bonds formed from direct overlap between atomic orbitals leading to electron sharing on the same axis as the bond.
Pi bonds are covalent chemical bonds where two lobes of one involved electron orbital overlap two lobes of the other involved electron orbital perpendicular to or on either side of the line drawn from the nucleus of one atom to another.
Double Bond: A chemical bond in which two pairs of electrons are shared between two atoms.
Triple Bond: A chemical bond in which three pairs (six total) electrons are shared between two atoms.
Hybridization: The concept explains how atomic orbitals fuse to form newly hybridized orbitals, which influences molecular geometry and bonding properties.
Bond energy is the amount of energy required to break a chemical bond and form neutral isolated atoms.
Activation Energy: This is the minimum amount of energy that reactants must have in order for a chemical reaction to occur.
Endothermic Reaction: A type of reaction where heat is absorbed from surroundings, similar to how breaking bonds requires an input of energy.
Exothermic Reaction: A type of reaction where heat is released into surroundings, akin to how forming bonds releases energy.
Bond length is defined as the average distance between two bonded atoms in a molecule.
Atomic Radius: This refers to half the distance between nuclei of two identical atoms bonded together - it influences bond length.
Single/Double/Triple Bonds: These refer to the number of shared electron pairs between atoms. More shared pairs mean shorter bond length and stronger bonds.
VSEPR Theory: This theory predicts molecular shape based on minimizing repulsion of shared and unshared electrons, which affects bond lengths.
The valence bond theory is a model that attempts to explain chemical bonding by considering that overlapping atomic orbitals create bonds, where electrons are shared between atoms.
Covalent Bonding: This is when pairs of electrons are shared by atoms, which is what valence bond theory primarily explains.
Sigma Bonds (σ): These are covalent bonds formed from direct overlap between atomic orbitals leading to electron sharing on the same axis as the bond.
Pi Bonds (π): These are covalent bonds formed by sideways overlapping of atomic orbitals with electron sharing above and below the bond axis.
A sigma bond is a type of covalent bond that forms when two atomic orbitals overlap along the axis connecting the two bonding nuclei.
Covalent Bond: A chemical bond formed by the sharing of one or more electrons, especially pairs of electrons, between atoms.
Atomic Orbital: A region in an atom where there is a high probability of finding electrons.
Molecular Orbital Theory: A method for determining molecular structure in which electrons are not assigned to individual bonds between atoms, but are treated as moving under the influence of the nuclei in the whole molecule.