Chemical bonds and interactions are among the most fundamental concepts in chemistry that you will be tested on during the MCAT. The properties and chemical behavior of substances are heavily dependent on the type of bonds and forces that hold them together. For instance, the strength of intermolecular forces determines the boiling and melting points of substances, while the type of atomic bonds determines the reactivity and stability of molecules.
These concepts are particularly relevant for those aspiring to enter the medical field, as general chemistry constitutes a significant portion of the MCAT section on the chemical and physical foundations of biology. Therefore, a thorough understanding of these concepts is necessary to excel in the entrance exam. Using our guide, you can start your prep and include all the important concepts of this topic in your . This article comprehensively explains all the key concepts regarding bonds and interactions that you will need to start .
We have divided this broader topic into three simpler sections and covered everything from the basic terminologies to the complex theoretical questions you may get to see in MCAT.
Disclaimer: MCAT is a registered trademark of AAMC. BeMo and AAMC do not endorse or affiliate with one another.
Listen to the blog!
Part I describes the major terminologies that you will have to understand before moving on to the complex concepts regarding molecular interactions.
As per Bohr’s atomic structure, atoms have an inner nucleus containing neutrons and protons. While electrons revolve in fixed paths (orbits) around the nucleus due to the electrostatic force of attraction of protons. The electrons residing in the outermost shell or orbit are called valence electrons.
The valence electrons determine whether the atom will form a bond or not. And if it does, with how many atoms will it combine? For instance, if an atom has one electron in its outermost, it will need one more electron to complete its duplet.
Want to know how to prepare for the MCAT for non-science majors? Watch this video:
For an element to be stable, its atoms must have two or eight electrons in their valence shell. However, except noble gases, all elements have incomplete valence shells. Thus, they form chemical bonds to complete their shells and achieve stability.
A chemical bond is formed when either an electron is shared, or transferred among atoms. This results in a strong force of attraction, which holds the atoms together.
Lone pair and Bond Pair
The valence electron pair that takes part in chemical bonding is called bond pair. Whereas, the electron pair that doesn’t participate in chemical bonding is called the lone pair.
Electronegativity is the ability of an atom to attract shared pair of electrons when forming bonds. This property is dependent on atomic number, so it has an increasing trend as we move from left to right in the . However, along the groups, as the size increases, the electronegativity of elements decreases due to a decrease in nuclear charge.
Since atomic and molecular structures are abstract concepts, they cannot be physically seen. So, to understand these complex molecular structures on paper, we draw a Lewis structure that represents the bond and non-bonding electrons in a molecule. It gives an overview of the valence electrons, the type of bonds, and the orientation of atoms in a molecule.
Another term, that is extensively used when talking about chemical bonds, is the formal charge. It describes the relative charge of an atom based on how the electrons are shared among atoms in a chemical bond. It doesn’t consider the electronegativities of atoms. Instead, factors in the valence electrons.
To calculate the formal charge, we subtract the sum of lone pair electrons and no. of bond pairs from the total valence electrons, which an atom possesses at its ground state. Mathematically,
FC = V - N - B/2
Here, V is the valence electrons, N is non-bonding/lone pair electrons, and B is the bonding electrons. For instance, the formal charge on carbon in carbon dioxide is 0 since there are 4 valence electrons and 4 bonded electron pairs (8 electrons).
FCcarbon = 4 - 0 - 8/2 = 0
***Correct answers for all Quick Questions are located at the end of this article.
Each element has orbitals (subshells) in which the electrons reside. And there are four different types of orbitals, s, p, d and f. Each has a different energy, orientation and electron-carrying capability.
Before bond formation, some orbitals may combine and form hybrid orbitals, which possess similar shapes and energy. This process of orbital intermixing is known as hybridization.
In general chemistry, we study three main types of hybridization: sp, sp2, and sp3. The detail on orbitals involved, structure and examples is illustrated in the following table.
The Types of Molecular Forces
Two types of forces hold the atoms within a substance, intramolecular forces and intermolecular forces. The forces that are responsible for holding the individual atoms together within a molecule are called intramolecular forces or chemical bonds. Whereas, the forces that hold a group of molecules together are intermolecular forces.
In this section, we will explore the major intramolecular forces i.e. the chemical bonds that are formed among similar to different atoms.
Ionic bond takes place due to the transfer of an electron from one atom to another due to a great difference in the electronegativities of two atoms. The threshold for the formation of an ionic bond is an electronegativity difference of 1.7.
Ionic bonds are usually formed between metals and non-metals. Since metals have excessive free electrons, they are readily available for electron donation and easily form cations (+ve ion). Whereas, non-metals are electron-deficient, have high electronegativities and readily form an anion (-ve ion). When these two different nature substances react, a new substance, with an entirely different structure and properties, is formed.
One common example of an ionic compound is table salt, NaCl. Independently, sodium is a reactive metal while chlorine is a reactive gas. However, their ionic bond results in a crystalline-structured table salt that we consume on daily basis.
Ionic compounds do not have discrete molecules; their chemical formula just shows the ratio of constituent elements. They have a three-dimensional structure with ionic bonds among the constituent atoms of the crystal lattice. Additionally, they are soluble in all polar solvents and water.
Figure 1-Crystal structure of NaCl ©CK12.org
A covalent bond is formed due to the sharing of electrons among individual atoms. These bonds are formed between similar atoms or the ones with relatively small electronegativity differences. Non-metals usually form covalent bonds.
Sigma (σ) and Pi (π) Bonds
The terms, sigma and pi bonds, are extensively used when speaking of covalent bonds. A sigma bond is formed by the head-to-head overlapping of simple or hybrid orbitals. In this case, the electron density is concentrated along an imaginary line connecting the nuclei of the atoms. The sigma bond does allow for rotation along the bond axis.
On the other hand, a pi bond is formed as a result of the lateral overlapping of p orbitals. In this case, the electron density is concentrated above and below the bond axis. Pi bonds are typically weaker than sigma bonds and do not allow for rotation around the bond axis.
Figure 2-sigma and pi-bonds in Ethene ©LibreTextsChemistry
Single, Double and Triple- Covalent bonds
Based on the pair of electrons shared in bonding, there are three types of covalent bonds: single, double and triple. Triple bonds have the highest bond strength while single bonds have the least.
When a single pair of electrons is shared among atoms, then that is a single covalent bond. A single covalent bond includes a sigma bond, formed as a result of sp3 hybridization. An example of a single covalent bond is hydrogen gas (H2).
If two or three pairs of electrons are shared, then such bonds are called double and triple bonds-respectively. A double covalent bond is a result of sp2 hybridization and comprises one sigma and one pi bond. Whereas, a triple covalent bond involves sp hybridization and has one sigma and two pi bonds.
The atoms of a molecule can form multiple types of bonds at a single time to fill its valence shell. For instance, in ethene, carbon simultaneously forms a double covalent bond with another carbon atom and a pair single of covalent bonds with hydrogen atoms, as shown in Figure 2.
Polar and Non-polar Covalent Bond
When a covalent bond is formed between two dissimilar atoms, a partial polarity is created among the two atoms due to their electronegativity difference. Such a covalent bond is known as a polar covalent bond.
A good example is water H2O. Oxygen has an electronegativity of 3.4, and hydrogen has an electronegativity of 2.2. As a result, oxygen forms a partially negative pole, while hydrogen acts as a positive pole.
In contrast, when a covalent bond is formed between two similar atoms or among atoms having an electronegativity difference of less than 0.4, such a covalent bond is termed a non-polar covalent bond. All gas molecules, H2, O2 and F2, are non-polar.
An important fact regarding bonds is that no chemical bond is completely ionic or covalent in nature. We call it ionic or covalent because one characteristic is dominant over the other.
Want to learn how to prepare for MCAT Chemistry? Check out this infographic:
Coordinate Covalent Bond
A coordinate covalent bond or dative covalent bond is a special type of covalent in which a pair of electrons is transferred from one atom to another. The atom donating the electron pair is called the donor atom, and the atom receiving the electron pair is called the acceptor atom. The donor atom usually has a lone pair of electrons that it can use to form the bond. Transition metals usually form dative covalent bonds with other atoms.
Drawing Lewis structure for Covalent Compounds
The bonds within a molecule of a covalent compound are shown through Lewis structures. It shows valence electrons as well as the bonds formed by the individual atoms inside a molecule.
To draw the Lewis structure of a molecule, first, you need to determine the valence electrons of individual elements constituting the molecule. For that, you can use the Periodic table. The group number of the element indicates its valence electrons.
For instance, consider a compound, silicone tetrafluoride (SiF4). Silicone is in group VI, indicating that it has 6 valence electrons. Whereas, Florine has 7 valence electrons as it resides in group VII of the periodic table.
Next, to draw its Lewis structure, we need a central atom. It is the one with the least electronegativity. In this case, that would be silicone. So, our Lewis structure would look like figure 3.
Figure 3-Lewis Structure of SiF4 ©Study.com
The straight lines in this Lewis structure represent single covalent bonds while dots represent the lone pairs. In the case of double and triple bonds, two and three lines are drawn, respectively.
In chemistry, resonance is a phenomenon that occurs when a molecule or ion can be described by one or more Lewis structures. Though each Lewis structure represents a possible arrangement of atoms and electrons in the molecule, none of them fully represents the true distribution of electrons in the molecule.
In all resonance structures, the position of atoms remains the same, but the location of electrons is different in each structure. Therefore, the position of bonds may also alternate among atoms.
Figure 4-Resonance Structures of Ozone ©LibreTextsChemistry
The above figure shows the two resonance structures of ozone. The double arrowhead doesn’t mean that the ozone structure fluctuates between these two states. Instead, it indicates that the actual Lewis structure of ozone will be the average of these two resonance structures.
Remember that a good resonance structure of a compound is the one which bears the least formal charge.
Valence Shell Electron Pair Repulsion (VSEPR) is a theoretical model used to predict the shapes of molecules based on the arrangement of electron pairs around the central atom. This theory assumes that electron pairs in the valence shell of the central atom repel each other and arrange themselves in a way that maximizes their distance from each other.
The molecular structure of a compound can be determined using the VSPER chart (see figure 5) after calculating the lone pairs and bond pairs of the central atom. This chart also gives brief information on bond angles.
Figure 5-VSEPR Chart, Developed by PSIBERG
To use this chart, we consider a covalent compound, methane (CH4). In this case, the central atom is carbon with a valency of 4. Carbon has 4 bond pairs and 0 lone pairs. So, according to the VSEPR table, its structure is tetrahedral and has a bond angle of 109.5
Generally, students are asked questions related to the molecular structure of different covalent bonds. So, you need to thoroughly go through this chart when preparing the chemistry section of your MCAT.
This part explains the intermolecular forces that form the physical characteristics, such as physical state, melting and boiling points, of a substance.
One important point regarding intermolecular forces is that they all are electrostatic in nature i.e. they exist due to the formation of positive and negative character in the independent molecules.
London dispersion Forces / Van der Waal Forces
London dispersion forces or Vander Waals’ forces exist in all types of molecules. However, they are prominent in non-polar molecules.
London dispersion forces are also called instantaneous dipole-induced forces based on the principle of how these forces work. In any molecule, when the electronic cloud is disturbed at any given moment, it results in the formation of partial positive and negative poles in the molecule. These instantaneous dipoles form temporary bonds with their neighbouring dipoles. We call these intermolecular forces London dispersion forces or van der Waal forces.
LDFs do exist in polar molecules. However, being weaker compared to other intermolecular forces, they are often ignored. Though they are the weakest intermolecular force, their strength tends to increase in large-size molecules. As the size increases, the valence electrons move farther from the nucleus, and hence the electronic cloud can be easily polarized through an external influence.
The common examples of LDFs are halogens: fluorine, chlorine, bromine and iodine. The first three halogens exist as gases, while iodine is a solid at room temperature. The reason is its large size and therefore strong Vander Waal forces.
As the name suggests, these forces exist in molecules with permanent dipoles i.e. polar molecules. In this case, when the molecule itself has one positive end and a negative end, it attracts the opposite ends of the neighboring atoms.
All polar molecules have a dipole-dipole interaction. This intermolecular force is stronger compared to LDF but weaker in comparison to hydrogen bonds. These forces are found usually in liquids since the molecules have to be closer for interaction.
Hydrogen is a special case of dipole-dipole interaction, in which hydrogen involvement is a requirement. One good example of hydrogen bonding is water (H2O).
In water, oxygen being more electronegative than hydrogen gets a partial negative charge. This negatively charged oxygen attracts the partially positive hydrogen of the neighbouring molecule. In this way, a chain of hydrogen bonds is formed among all molecules.
Hydrogen bonds are the strongest intermolecular forces and only exist in a substance containing hydrogen, oxygen, nitrogen or fluorine.
We hope that this comprehensive guide has helped you understand the key concepts regarding bonds and interactions. With some quick questions, we have tried to give you an idea of how the questions may be asked in the MCAT. After reviewing this guide, we recommend you solve multiple sample assessments provided by the AAMC.
Answers to Quick Questions
- C (Hint: V=4, N=2, and B=6)
- D (A double covalent bond is involved)
- C (Refer to VESPR chart )
- C (Polar molecules don’t dissolve in polar molecules)
- D (van der Waal forces in Ar < Dipole Forces< HB)