Chemistry Lecture Review: Chemical Bonds, Molecular Geometry, & Intermolecular Forces
Chemical Bonds
Atoms in a compound are held together by chemical bonds.
Ionic bonds: Formed when valence electrons are transferred from one atom to another.
Covalent bonds: Formed when valence electrons are shared between atoms.
Atomic Charge and Ion Formation
Charge Formula: \text{Charge} = \text{Number of protons} - \text{Number of electrons} .
Metals (left side of periodic table):
Tend to lose electrons.
Form positively charged ions called cations (e.g., a metal losing two electrons becomes a cation with a +2 charge).
Nonmetals (right side of periodic table):
Tend to gain electrons.
Form negatively charged ions called anions (e.g., a nonmetal gaining one electron becomes an anion with a -1 charge).
Calculating Charge for Nonmetal Anions (Simplified Rule): \text{Charge} = \text{Group number} - 8 .
Example: Sulfur (S) is in Group 6A. 6 - 8 = -2 . So, the symbol for a sulfur ion is S^{2-} .
Naming Ions
Cations of Transition Metals with Variable Charges:
A Roman numeral is used in parentheses to indicate the charge when a metal can form more than one type of cation.
Example: Fe^{2+} is an iron(II) ion; Fe^{3+} is an iron(III) ion.
No Roman numeral is included if the metal only has one ionic form (e.g., silver ion, Ag^{+}; zinc ion, Zn^{2+}). These are examples of transition metals that exist in only one ionic form.
Anions: Often end in "-ide" (e.g., O is oxygen, O^{2-} is oxide).
Polyatomic Ions
A polyatomic ion is an ion derived from a molecule (a group of atoms), rather than a single atom.
Each polyatomic ion has a unique name, chemical formula, and overall charge.
Example: The carbonate ion, CO_{3}^{2-} , is a polyatomic anion composed of one carbon atom and three oxygen atoms, with a collective -2 charge.
Naming Covalent Compounds
Prefixes are used to indicate the number of each type of atom in the molecule.
Example: N_{2}O is named dinitrogen monoxide.
The prefix 'di-' before nitrogen indicates two nitrogen (N) atoms.
The prefix 'mono-' before oxygen indicates one oxygen (O) atom.
Comparing Ionic and Covalent Compounds
Ionic Compounds:
Formed by electron transfer, resulting in charged ions.
When components combine (e.g., sodium metal and chlorine gas), the resulting compound has vastly different reactivity, color, and general properties compared to the individual elements.
Have a net charge equal to zero, which determines the ratio of cations to anions.
This ratio is reflected in the subscripts of the compound's formula unit.
Exist as highly ordered structures known as lattices (e.g., crystal lattice).
Covalent Compounds:
Formed by sharing valence electrons between atoms.
Can exist as solids, liquids, or gases at room temperature.
Are composed of discrete molecules with the same composition, specified by their molecular formula.
Molecular formula: specifies the number of each type of atom in a molecule.
Binary covalent compounds: Simplest covalent compounds, composed of molecules with only two different types of atoms.
Diatomic Elements:
Certain elements exist naturally as molecules composed of two identical nonmetal atoms.
These include: Hydrogen (H{2}), Nitrogen (N{2}), Oxygen (O{2}), Fluorine (F{2}), Chlorine (Cl{2}), Bromine (Br{2}), and Iodine (I_{2}).
Describing Molecular Structure
Lewis Dot Structure:
A representation that accounts for all valence electrons present in a molecule.
Shows the arrangement of atoms, the types of bonds (single, double, triple), and nonbonding (lone pair) electrons.
This information is crucial for determining molecular geometry.
Octet Rule:
Bonds typically form between atoms to achieve a filled valence energy level, usually with eight electrons.
Exception: Hydrogen (H) only forms one bond to obtain a filled n=1 energy level with two electrons.
Ball-and-Stick Models:
Represent atoms as spheres ("balls") and covalent bonds as connecting lines ("sticks").
Designed to depict molecular geometry and bond angles visually.
Molecular Geometry
Molecular geometry is the three-dimensional shape of a molecule, defined by the relative positions of the atoms within it.
The Lewis dot structure provides all the necessary information to determine the molecular geometry.
Steps to Determine Molecular Geometry:
Determine Electron Geometry: This is the relative position of the groups of electrons on the central atom.
A "group of electrons" refers to a single bond, a double bond, a triple bond, or a nonbonding pair of electrons (lone pair).
Determine Molecular Geometry: This is determined from the electron geometry by evaluating the number of bonding versus nonbonding electron groups on the central atom.
VSEPR (Valence Shell Electron Pair Repulsion) Theory:
Predicts that electron groups (bonding pairs and lone pairs) around a central atom will arrange themselves to minimize repulsion, thus maximizing the distance between them.
Common Electron Geometries based on Electron Groups/Domains/Regions:
Two groups: Linear (separated by 180^{\circ}).
Three groups: Trigonal planar.
Four groups: Tetrahedral.
Role of Lone Pairs:
Lone pairs (nonbonding regions) are smaller than bonded atoms but still exert repulsive forces, influencing the final molecular shape.
Examples of molecular geometry derivations:
If the electron geometry is trigonal planar and there is one lone pair, the molecular geometry is bent.
If the electron geometry is tetrahedral and there is one lone pair, the molecular geometry is trigonal pyramidal.
If the electron geometry is tetrahedral and there are two lone pairs, the molecular geometry is bent.
Bond Angle: The angle between the central atom and any two atoms bonded to it.
Representing Molecular Geometries on Paper
Straight line: Represents a bond lying in the plane of the paper.
Hashed lines (or dashed lines): Represents a bond projecting into the paper, away from the viewer.
Wedge: Represents a bond projecting out of the paper, towards the viewer.
Molecular Polarity: Bond Dipoles and Electronegativity
Molecular polarity: Determines how molecules interact with each other at the molecular level.
Electronegativity: A measure of an atom's ability to draw electrons toward its nucleus when it is part of a covalent bond.
It is an intrinsic property determined by an atom's electron arrangement.
The most electronegative element is Fluorine (F), with an electronegativity value of 4.0.
Nonmetals are generally more electronegative than metals.
Valence electrons (those in the outermost energy level) determine electronegativity.
Smaller atoms (where valence electrons are closer to the positively charged nucleus) have greater electronegativity.
Within a group on the periodic table, electronegativity decreases from top to bottom (as atomic size increases).
Bond Polarity: The difference in electronegativity between two atoms forming a covalent bond determines whether the bond is polar or nonpolar.
Polar Covalent Bond (or Polar Bond):
Forms when electrons are shared unequally due to a significant electronegativity difference.
The more electronegative atom attracts electron density more strongly, creating a partial negative charge ( \delta^{-} ).
The less electronegative atom obtains a partial positive charge ( \delta^{+} ).
This separation of partial charges is called a dipole.
Example: An O-H bond is very polar because oxygen is much more electronegative than hydrogen.
Nonpolar Bond:
Forms when electrons are shared equally between two atoms.
Electron density is, on average, evenly distributed.
Do not have partial charges or bond dipoles.
Found in diatomic elements (e.g., H{2}, O{2}) and in C-C and C-H bonds of hydrocarbons (compounds containing only carbon and hydrogen).
Electronegativity Difference and Bond Type Classification:
Difference < 0.5: Nonpolar covalent bond.
Difference between 0.5 and 1.9: Polar covalent bond.
Difference \ge 2.0 : Ionic bond.
Overall Molecular Polarity:
An entire molecule can be polar or nonpolar, depending on its molecular shape and the polarity of its bonds.
Polar molecules: Have an overall charge separation, meaning an uneven distribution of electrons, resulting in distinct positive and negative poles (e.g., \delta^{+} and \delta^{-} ).
Nonpolar molecules: Have an even distribution of electrons, without distinct positive and negative poles.
Polar and nonpolar molecules exhibit very different physical and chemical properties.
Indicators of Molecular Polarity:
Uneven electron distribution or the presence of lone pairs on the central atom often leads to a polar molecule (e.g., a molecule with a lone pair region will have an electron-rich, negatively charged area).
To determine molecular polarity, the Lewis structure and molecular geometry are essential.
Intermolecular Forces of Attraction (IMFs)
Intramolecular forces of attraction: The strong forces within a molecule that hold atoms together (e.g., ionic bonds, covalent bonds).
Intermolecular forces of attraction: The forces of attraction or repulsion that act between neighboring particles (atoms, molecules, or ions).
These attractions occur in the solid and liquid phases; in the gas phase, molecules are too far apart to interact significantly.
IMFs are much weaker than intramolecular (covalent) forces, typically only 5\% to 10\% the strength of an average covalent bond.
IMFs determine the macroscopic physical properties of a compound, such as boiling point, melting point, and solubility.
The strength of IMFs can be inferred from boiling points; higher boiling points indicate stronger intermolecular interactions.
Three Main Types of Intermolecular Forces (in pure compounds):
Dispersion Forces (London Dispersion Forces)
Dipole-Dipole Forces
Hydrogen Bonding Forces
1. Dispersion Forces (London Dispersion Forces)
The weakest type of intermolecular force of attraction.
They are the only intermolecular force present in nonpolar compounds and noble gas elements.
Nonpolar molecules do not have permanent molecular dipoles.
Caused by momentary shifts in electron distribution, creating a temporary, instantaneous separation of charge known as a temporary molecular dipole.
Since electrons are constantly in motion and distributed throughout a molecule, there are many opportunities for these temporary dipoles to form.
Effect of Shape and Size:
Larger molecules (with more electrons and greater surface area) have more opportunities for temporary dipoles and thus exhibit stronger dispersion forces.
Example: Alkanes show an increase in boiling point as their molar mass (and hence size) increases, due to stronger dispersion forces.
Compact molecular shapes lead to weaker dispersion forces due to less surface area for interaction, compared to extended shapes with the same molar mass (constitutional isomers).
2. Dipole-Dipole Forces
Occur between the opposite partial charges of permanent dipoles in polar molecules.
These interactions are more permanent and generally stronger than dispersion forces for molecules of comparable size.
Example: Formaldehyde (H_{2}CO) has a permanent dipole due to the polar C=O bond. The partial positively charged end of one formaldehyde molecule aligns with the partial negatively charged end of another molecule, and so forth throughout the sample.
3. Hydrogen Bonding Forces
The strongest type of intermolecular force of attraction between molecules.
This is a special, strong type of dipole-dipole force.
Occurs in molecules that possess one of three highly polar covalent bonds:
Hydrogen-Fluorine (H-F)
Hydrogen-Oxygen (H-O)
Hydrogen-Nitrogen (H-N)
These bonds are the most polar covalent bonds because they involve hydrogen and highly electronegative nonmetal atoms (F, O, or N) that are at opposite ends of the electronegativity scale.
Hydrogen bonding involves an attraction between the partial positive pole (the hydrogen atom, \delta^{+} ) of one molecule and the partial negative pole (the oxygen, nitrogen, or fluorine atom, \delta^{-} ) of another molecule.
Represented by a dashed line between the hydrogen atom of an O-H or N-H group on one molecule and an O, N, or F atom on an adjacent molecule.
Biological Significance: Hydrogen bonding plays crucial roles in the structures and functions of many biological molecules, including proteins, carbohydrates, DNA, and RNA
CHAPTER 3
A compound is composed of two or more atoms joined by chemical bonds. Ionic bonds are formed when valence electrons are transferred from one atom to another. A covalent bond is formed when valence electrons are shared between atoms. The loss of one or more valence electrons from a metal creates a positively charged ion, a cation. The gain of one or more valence electrons by a nonmetal creates a negatively charged ion, an anion. The strong attraction between oppositely charged ions is known as an ionic bond. Ionic compounds have very different physical and chemical properties than the elements from which they are derived. Cations have a positive charge because they have fewer electrons than protons. Anions have a negative charge because they have a greater number of electrons than protons. Groups 1A, 2A, and 3A metals can lose their one, two, and three valence electrons to become ions with the same charge as their group number: + 1 , + 2 , and + 3 , with the symbols 𝑀 + ,  𝑀 2 + , and 𝑀 3 + , where 𝑀 is the symbol of the atom from which the ion is derived. Many transition metals and group 4A metals exist in more than one ionic form. To name these ions, a Roman numeral corresponding to the magnitude of the charge is enclosed in parentheses following the cation name. Groups 5A–7A nonmetals gain one to three electrons to achieve a filled valence energy level. They become anions with a - 3 ,  - 2 , or - 1 charge, written 𝑋 3 - ,  𝑋 2 - ,  𝑋 - . Anions are named by changing the ending of the element name to -ide, followed by “ion.” Ionic compounds are represented by a formula unit, which is the lowest whole-number ratio of cations to anions that results in an electrically neutral compound. To write a formula unit, the cation is written first, followed by the anion, excluding the superscripts (charges), and inserting subscripts following the symbols of the ions to represent their ratio. To name an ionic compound, write the name of the cation followed by the name of the anion. For transition metal cations that can exist in multiple forms, the charge on the ion is indicated by a Roman numeral enclosed in parentheses following the symbol. 3.2 Ionic Compounds Containing Polyatomic Ions A polyatomic ion is an ion derived from a molecule that has gained or lost valence electrons. A polyatomic ion has a unique name, formula, and charge. Writing the formula unit for an ionic compound containing a polyatomic ion follows the same process as with monatomic ions; however, parentheses must be placed around the polyatomic ion when a subscript follows it. 3.3 Covalent Compounds Covalent compounds are composed of molecules, which are discrete entities composed of two or more different atoms held together by covalent bonds. The diatomic elements are molecules. A single bond forms when two atoms share two valence electrons, typically one electron contributed by each atom. A molecular formula provides the number and type of atoms in a molecule by listing the atomic symbols, followed by subscripts. 3.4 Writing Lewis Dot Structures of Covalent Compounds A chemical structure shows the arrangement of atoms and covalent bonds between atoms in a molecule. When writing a Lewis dot structure, bonds are formed such that atoms in period 2 and higher have 8 electrons, the octet rule. Nonbonding electrons are valence electrons of an atom that are not shared; they are depicted as pairs of dots in a Lewis dot structure. In a Lewis dot structure, a shared pair of electrons is represented by a single bond; two shared pairs of electrons, a double bond; and three shared pairs of electrons, a triple bond. Key Words Binary covalent compound: A covalent compound composed of two different types of atoms. Chemical bond: The interaction between valence electrons that holds ions or atoms together in an ionic compound and a covalent compound. Chemical structure: A notation that conveys the arrangement of the atoms and the bonds in a molecule. Compound: A substance composed of two or more different types of atoms. Covalent bond: Valence electrons shared between two nonmetal atoms in a molecule. A single bond refers to two shared electrons; a double bond, four shared electrons; and a triple bond, six shared electrons. Covalent compound: A substance composed of molecules that have two or more different types of atoms. Double bond: The sharing of two pairs of electrons, four valence electrons, between two atoms. A double bond is denoted as two parallel lines in a Lewis dot structure, where each line represents a pair of bonding electrons. Electrolytes: Ions dissolved in water. They are able to conduct electricity. Formula unit: The cations and anions and their whole-number ratio (subscripts) that define an ionic compound. Indicated by atomic symbols (without superscripts) and subscripts. Ionic bond: The strong attraction between the opposite charges of a cation and an anion in an ionic compound. Ionic compound: A compound formed when a metal transfers valence electrons to a nonmetal to form a cation and an anion, which are held together by the attraction of their opposite charges. Includes ionic compounds composed of polyatomic ions. Lewis dot structure: A notation showing the arrangement of atoms and the bonds between atoms in a molecule, in addition to the nonbonding pairs of electrons on atoms. Covalent bonds are drawn as lines, representing shared pairs of electrons, and nonbonding electrons are drawn as pairs of dots. Lewis dot symbol: A way of representing an atom and its valence electrons by writing the atomic symbol surrounded on up to four sides by its valence electrons, shown as dots. Molecular formula: The composition of a covalent compound given by the atomic symbol of each atom in the molecule, followed by a subscript indicating the number of each atom, and listed in alphabetical order. Molecule: Two or more nonmetal atoms that share electrons—form covalent bonds. Monatomic anion: A negatively charged ion resulting from the gain of one or more electrons by a nonmetal atom. Monatomic cation: A positively charged ion resulting from the loss of one or more electrons from a metal atom. Monatomic ions: Cations and anions derived from a metal or nonmetal atom, such as Na + and Cl - . Nonbonding electrons: The unshared valence electrons on an atom in a molecule. Denoted as a pair of dots on the atom in a Lewis dot structure. Octet rule: The tendency for most atoms when part of a molecule to form one to four bonds, so as to attain eight valence electrons—a filled valence energy level. Polyatomic ion: An ion formed from the loss or gain of electrons from a molecule, resulting in a positive or negative charge localized or delocalized over several atoms with covalent bonds. Salt: Another term for an ionic compound. Single bond: A pair of valence electrons shared between two atoms, typically one electron contributed by each atom. A single bond is denoted as a line, representing a pair of electrons, in a Lewis dot structure. Triple bond: Two atoms that share three pairs of electrons; six valence electrons. A triple bond is denoted as three parallel lines, each representing a pair of bonding electrons, in a Lewis dot structure.
CHAPTER 4.
A molecular geometry describes the relative position of the atoms in a molecule. The electron geometry for a molecule is determined by the number of electron groups on the central atom and VSEPR theory. Molecular geometry is determined from electron geometry based on the number of bonding versus nonbonding groups on the central atom. A molecule with a linear electron geometry also has a linear molecular geometry. The two molecular geometries that arise from a trigonal planar electron geometry are trigonal planar (three bonding groups, no nonbonding pairs of electrons), and bent (two bonding groups, one nonbonding pair of electrons). The three molecular geometries that arise from a tetrahedral electron geometry are tetrahedral, trigonal pyramidal, and bent. Bond angles are determined by a molecule’s electron geometry: tetrahedral, 109.5 ° ; trigonal planar, 120 ° ; or linear, 180 ° . The shapes of larger molecules are described by assigning one of the five basic molecular geometries to each atom center, separately. 4.2 Bond Dipoles and Molecular Polarity Electronegativity is a measure of an atom’s ability to draw electrons toward its nucleus when part of a covalent bond. Electronegativity increases within a group from bottom (larger atom) to top (smaller atom) and increases across a period from left to right. A polar covalent bond has a separation of charge. A partial negative charge ( 𝛿 - ) on the more electronegative atom and a partial positive charge ( 𝛿 + ) on the less electronegative atom. A polar covalent bond can also be represented by a dipole arrow. Bonds between atoms that have the same or very similar electronegativities ( EN < 0.5 ) are nonpolar: they have no charge separation and do not contain a bond dipole. Nonpolar covalent bonds are formed between atoms that have equal or similar electronegativity values; and polar covalent bonds are formed between atoms that have different electronegativities. A molecule is nonpolar if (1) all the bonds are nonpolar or (2) the molecular geometry is symmetrical (the central atom has no nonbonding electrons) AND all the bonds are identical. A molecule is polar when it has one polar bond; two or more polar bonds and a molecular geometry that is unsymmetrical; or it has two or more polar bonds that are not identical and a symmetrical molecular geometry. 4.3 Intermolecular Forces of Attraction in a Compound Intermolecular forces of attraction occur between molecules, which is a much weaker force of attraction than between atoms in a covalent bond. The three types of intermolecular forces of attraction are dispersion, dipole–dipole, and hydrogen bonding. Dispersion forces are the only forces of attraction in elements and nonpolar compounds. Dispersion forces result from induced temporary dipoles. Dipole–dipole intermolecular forces of attraction are found in polar molecules that have a permanent dipole. They, are stronger than dispersion forces. Hydrogen bonding, the strongest intermolecular force of attraction, occurs between polar molecules that contain O — H or N — H bonds; it also occurs in H — F . Key Words Ball-and-stick model: A three-dimensional representation of a molecule in which colored spheres represent atoms and gray or white connectors represent covalent bonds. They are used to visualize the shape of a molecule by clearly showing the bond angles between atoms. Bent molecular geometry: One of the five basic molecular geometries. A molecular geometry that arises from either a tetrahedral electron geometry when there are two nonbonding pairs of electrons and two bonding groups of electrons; and also from a trigonal planar electron geometry when there is one nonbonding pair of electrons and two bonding groups of electrons. The bond angle is 109.5 ° when derived from a tetrahedral electron geometry, and 120 ° when derived from a trigonal planar electron geometry. It is not a symmetrical molecular geometry. Bond angle: The angle generated by any three atoms, but for simple molecules, refers to the angle between the central atom and two atoms bonded to it. Bond dipole: The separation of partial charges, 𝛿 + and 𝛿 - , in a polar covalent bond. Bond dipoles exist in covalent bonds between atoms that have electronegativity differences ranging from 0.5   to   2.0 . Dipole arrow: An arrow with a hatch mark opposite the head of the arrow. It is placed parallel to a polar bond to designate a bond dipole. The head of the arrow points toward the more electronegative atom, and the hatch mark lies next to the less electronegative (electropositive) atom. Dipole–dipole forces: The intermolecular forces of attraction between molecules that have a permanent dipole, but do not contain the stronger N — H ,   O — H ,   or   H — F bond dipoles. Dispersion forces: The weakest intermolecular forces of attraction. They occur between the partial charges in induced temporary dipoles in all compounds. They are the only intermolecular forces of attraction in nonpolar compounds. Electron density diagram: A space-filling model that includes color to show regions of high electron density ( 𝛿 - ) and low electron density ( 𝛿 + ) in a molecule. Electron geometry: The relative position of the groups of electrons (bonding and nonbonding) around a central atom. Electron geometry determines bond angles around an atom center and is one factor in determining the molecular geometry of a molecule. Electron Group: As applied to VSEPR theory, an electron group is a single bond, a double bond, a triple bond, or a nonbonding pair of electrons. Electronegativity: A measure of an atom’s ability to attract electrons toward itself in a molecule. Hydrocarbons: Compounds containing only C — C and C — H bonds. Hydrogen bonding: The strongest type of intermolecular forces of attraction between molecules. Hydrogen bonding occurs in polar molecules that have one or more H — O or H — N bonds, as well as the compound H — F . Intermolecular forces of attraction: The forces of attraction between molecules, which arise from the partial positive charge on one molecule and the partial negative charge on another molecule. Partial charges may be temporary, as in dispersion forces, or permanent, as in dipole–dipole and hydrogen bonding intermolecular forces of attraction. Linear electron geometry: The three-dimensional shape of two groups of electrons on a central atom. The groups of electrons point in opposite directions, 180 ° apart. Linear molecular geometry: One of the five basic molecular geometries. Characterized by a central atom that has two groups of bonding electrons and no nonbonding electrons. It is a symmetrical molecular geometry. Molecular dipole: The dipole, or charge separation, in a polar molecule. Molecular geometry: The relative position of the atoms in a simple molecule (a central atom with bonds to two, three, or four atoms). Molecular geometry is determined from the electron geometry by evaluating the relative number of bonding and nonbonding groups on the central atom. Molecular polarity: Charge separation within a molecule; two poles, one that is partially positive and one that is partially negative. Nonpolar covalent bond: A covalent bond between two atoms with the same or similar electronegativities. Defined as an electronegativity difference of less than 0.5. Nonpolar molecule: A molecule which has an even distribution of electrons; no separation of charge. A molecule is nonpolar when the bond dipoles in the molecule cancel or when there are no bond dipoles in the molecule. Partial negative charge, 𝛿 - : A charge less than - 1 on an atom or part of a molecule. Partial positive charge, 𝛿 + : A charge less than + 1 on an atom or part of a molecule. Polar covalent bond: A covalent bond in which there is a separation of charge, that is, a dipole. A polar covalent bond exists when two atoms with significantly different electronegativity values share electrons. Defined as an electronegativity difference between 0.5 and 2.0 . Polar molecule: A molecule that has a separation of partial charges. A molecule is polar when it has one or more bond dipoles that do not cancel. Space-filling model: A three-dimensional representation of a molecule that illustrates the relative amount of space occupied by the atoms in the molecule. Large colored spheres represent the atoms and the relative amount of space that they occupy. Temporary molecular dipole: The momentary separation of charge that occurs in nonpolar molecules and elements. Tetrahedral electron geometry: The three dimensional shape created by of four groups of electrons on a central atom. The groups of electrons point to the four corners of a tetrahedron, 109.5 ° apart. Tetrahedral molecular geometry: One of the five basic molecular geometries. Characterized by a central atom that has four groups of bonding electrons and no nonbonding electrons. A symmetrical molecular geometry. Trigonal planar electron geometry: The three-dimensional shape created by three groups of electrons on a central atom. The groups of electrons point to the three corners of an equilateral triangle, 120 ° apart. Trigonal planar molecular geometry: One of the five basic molecular geometries. Characterized by a central atom that has three groups of bonding electrons and no nonbonding electrons. A symmetrical molecular geometry. Trigonal pyramidal molecular geometry: One of the five basic molecular geometries. Characterized by a central atom that has three groups of bonding electrons and one nonbonding pair of electrons. Bond angles are approximately 109.5 ° . It is not a symmetrical molecular geometry. Valence shell electron pair repulsion (VSEPR): A theory used to predict the molecular geometry of molecules. VSEPR is based on the principle that electrons, both bonding and nonbonding, are repulsive and therefore, adopt a geometry where they are as far apart from each other as geometrically possible.