Intermolecular Properties Notes - AP Chemistry -New Syllabus 2024-2025
AP Chemistry – Intermolecular Properties Notes
AP Chemistry – Intermolecular Properties Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
Explain the relationship between the chemical structures of molecules and the relative strength of their intermolecular forces when:
- The molecules are of the same chemical species
- The molecules are of two different chemical species.
Key Concepts:
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London Dispersion Forces
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Dipole-Induced Dipole Interactions
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Dipole-Dipole Interactions
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Ion-Dipole Interactions
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Molecular Dipole Moment
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Hydrogen Bonding
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Interactions in Large Biomolecules
3.1.A.1 London Dispersion Forces:
1. Intermolecular Forces:
Intermolecular forces (IMFs) are the forces of attraction or repulsion that exist between molecules. These forces are much weaker than the intramolecular forces (covalent or ionic bonds) that hold atoms together within a molecule, but they play a crucial role in determining the physical properties of substances, such as boiling and melting points, vapor pressure, and solubility.
i. London Dispersion Forces (LDF):
a. Description: These are the weakest type of intermolecular force and are present in all molecules, whether polar or nonpolar. London dispersion forces arise due to temporary fluctuations in the electron distribution around atoms or molecules, which create temporary dipoles. These dipoles induce corresponding dipoles in neighboring molecules, leading to weak attractions.
b. Strength: Weakest, but significant in larger atoms or molecules with more electrons (e.g., larger nonpolar molecules).
c. Examples: These forces are significant in noble gases (e.g., argon), and in nonpolar molecules like methane (CH₄), and iodine (I₂).
d. Key Point: The strength of London dispersion forces increases with the size and shape of the molecule because larger molecules have more electrons that can cause more significant dipoles.
ii. Dipole-Dipole Interactions:
a. Description: These forces occur between polar molecules. In a polar molecule, one end has a partial positive charge, and the other end has a partial negative charge. The positive end of one molecule is attracted to the negative end of another molecule.
b. Strength: Stronger than London dispersion forces but still weaker than ionic or covalent bonds.
c. Examples: Hydrogen chloride (HCl), acetone (CH₃COCH₃), and water (H₂O).
d. Key Point: Polar molecules experience dipole-dipole interactions, which can lead to higher boiling points compared to nonpolar molecules of similar size.
iii. Hydrogen Bonding:
a. Description: A specific type of dipole-dipole interaction that occurs when a hydrogen atom covalently bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine) is attracted to a lone pair of electrons on another electronegative atom. It is a particularly strong dipole-dipole interaction.
b. Strength: Stronger than regular dipole-dipole interactions but still weaker than covalent bonds.
c. Examples: Water (H₂O), ammonia (NH₃), and hydrogen fluoride (HF).
d. Key Point: Hydrogen bonds are responsible for the high boiling points of water and the unique properties of substances like DNA and proteins.
iv. Ion-Dipole Forces:
a. Description: These forces occur between an ion and the partial charge on a polar molecule. The ion attracts the oppositely charged end of the polar molecule.
b. Strength: Stronger than dipole-dipole interactions, typically seen when ionic compounds dissolve in polar solvents.
c. Examples: The attraction between sodium ions (Na⁺) and water molecules in an aqueous solution.
d. Key Point: Ion-dipole interactions are important for the solubility of salts in water.
v. Ion-Induced Dipole and Dipole-Induced Dipole Forces:
a. Description: These are weaker forces that occur when an ion or a polar molecule induces a dipole in a neighboring nonpolar molecule. In ion-induced dipole interactions, the ion creates a dipole in a nearby nonpolar molecule. In dipole-induced dipole interactions, a polar molecule induces a dipole in a nonpolar molecule.
b. Strength: Weaker than dipole-dipole or ion-dipole interactions.
c. Examples: The interaction of an ion with a nonpolar molecule, such as an ion inducing a dipole in an oxygen molecule (O₂).
d. Key Point: These interactions are important in some specialized systems, like in the behavior of nonpolar gases in the presence of ions.
2. Coulombic Interactions:
Coulombic interactions are forces that arise from the attraction between charged particles. In the context of temporary dipoles, Coulombic interactions are responsible for the forces that occur due to fluctuations in the electron distribution within atoms and molecules.
i. Formation of Temporary Dipoles through Electron Fluctuations:
a. Electron Movement: Electrons in atoms or molecules are constantly in motion due to their wave-like behavior. At any given moment, the distribution of electrons may not be perfectly symmetrical, creating an uneven charge distribution around the atom or molecule.
b. Temporary Dipole Creation: This temporary uneven distribution of electrons creates a fleeting or “instantaneous” dipole. A dipole occurs when one part of the atom or molecule becomes slightly negatively charged (where electrons are more concentrated), and another part becomes slightly positively charged (where electrons are less concentrated).
c. Induced Dipoles in Neighboring Molecules: When this temporary dipole exists, its fluctuating charge can influence nearby molecules. The partial positive charge of the temporary dipole will attract the electrons in neighboring molecules, while the partial negative charge will repel them. This results in the induction of a dipole in the nearby molecule, which is called an induced dipole.
d. Coulombic Attraction: The temporary dipole and the induced dipole experience a Coulombic attraction, as opposite charges attract. This interaction between the temporary dipole and the induced dipole is a form of London dispersion force, which is a type of Coulombic interaction. The strength of this interaction depends on the polarizability of the molecules involved (how easily their electron cloud can be distorted).
3. Polarizability:
Polarizability refers to the ability of a molecule’s electron cloud to be distorted by external electric fields or nearby charges, such as a temporary dipole in neighboring molecules. The more easily the electron cloud can be distorted, the more polarizable the molecule is. Polarizability plays a significant role in the strength of London dispersion forces, which are a type of intermolecular force arising from temporary dipoles that induce dipoles in neighboring molecules.
i. Molecular Size:
a. Larger molecules tend to have more electrons and larger atomic radii, which make their electron clouds more spread out and less tightly held by the nucleus. This increased size means that it’s easier to induce a temporary dipole, making the molecule more polarizable.
b. Effect on dispersion forces: As the size of the molecule increases, the London dispersion forces become stronger. Larger molecules have greater electron clouds that can fluctuate more easily, leading to stronger temporary dipoles and enhanced interactions with neighboring molecules.
c. Examples: For example, iodine (I₂), a large molecule, exhibits stronger dispersion forces than fluorine (F₂), which is much smaller.
ii. Electron Count:
a. The number of electrons in a molecule directly influences its polarizability. Molecules with more electrons generally have a more easily distorted electron cloud because the electrons are farther from the nucleus and experience less effective nuclear charge (i.e., less attraction from the nucleus).
b. Effect on dispersion forces: Molecules with a higher electron count are more polarizable and, therefore, experience stronger London dispersion forces.
c. Examples: Heavy noble gases like xenon (Xe) and krypton (Kr) exhibit stronger dispersion forces than lighter noble gases like helium (He) and neon (Ne) due to their greater number of electrons.
iii. Pi Bonding:
a. Pi bonds are formed from the overlap of p-orbitals and are typically found in double and triple bonds (as in alkenes, alkynes, and aromatic compounds). These bonds create regions of electron density that are more exposed to the surrounding environment, making the molecule more susceptible to distortion.
b. Effect on dispersion forces: Molecules with pi bonds (especially aromatic rings) tend to be more polarizable because the electron density in the pi system is less tightly bound and more easily influenced by external interactions. This makes these molecules exhibit stronger dispersion forces.
c. Examples: Benzene (C₆H₆), with its aromatic pi bonding, has stronger dispersion forces than a similarly sized alkane like hexane (C₆H₁₄) due to the polarizability introduced by the delocalized electrons in the pi bonds.
4. Van der Waals Forces:
Van der Waals forces is a general term that refers to the intermolecular forces that exist between molecules, including London dispersion forces, dipole-dipole interactions, and hydrogen bonding. These forces are weaker than the covalent or ionic bonds that hold atoms together within a molecule, but they still play a crucial role in determining the physical properties of substances, such as boiling points, melting points, and solubility.
Relationship Between London Dispersion Forces and Van der Waals Forces:
i. London Dispersion Forces as a Subtype of Van der Waals Forces:
a. London dispersion forces (also known as instantaneous dipole-induced dipole interactions) are a subset of Van der Waals forces. They arise due to temporary fluctuations in the electron distribution within molecules, creating temporary dipoles. These temporary dipoles can then induce dipoles in neighboring molecules, leading to an attraction between them.
b. In the broader context of Van der Waals forces, London dispersion forces are the primary attractive force between nonpolar molecules or atoms. These forces are always present, regardless of whether the molecules are polar or nonpolar, but their strength is most significant in nonpolar molecules.
ii. Other Components of Van der Waals Forces: While London dispersion forces are the dominant type of Van der Waals forces in nonpolar molecules, Van der Waals forces also encompass:
a. Dipole-Dipole Interactions: These occur between polar molecules, where the positive end of one polar molecule is attracted to the negative end of another polar molecule.
b. Hydrogen Bonding: This is a special case of dipole-dipole interactions where a hydrogen atom covalently bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine) is attracted to a lone pair of electrons on another electronegative atom.
c. Dipole-Induced Dipole Forces: These occur when a polar molecule induces a dipole in a neighboring nonpolar molecule.
ii. Overall Van der Waals Interactions: The overall Van der Waals force acting between two molecules or atoms is the result of the sum of all these individual interactions, including:
a. London dispersion forces (important for all molecules, especially nonpolar ones),
b. Dipole-dipole interactions (important for polar molecules),
c. Hydrogen bonding (a stronger form of dipole-dipole interaction),
d. Ion-dipole and ion-induced dipole forces (important in ionic solutions).
3.1.A.2 Dipole Interactions:
1. Dipole-Induced Dipole Interactions:
i. What are they?
- These interactions occur between a polar molecule and a nonpolar molecule.
- The polar molecule creates a temporary dipole in the nonpolar molecule by interacting with its electron cloud. The positive end of the polar molecule attracts electrons in the nonpolar molecule, while the negative end repels them, inducing a dipole.
ii. Points to Keep in Mind:
- Always attractive: The interaction is always attractive because opposite charges (positive on the polar molecule and negative on the induced dipole) attract each other.
- Strength depends on two factors:
- Magnitude of the dipole of the polar molecule: The stronger the dipole moment of the polar molecule, the greater its ability to induce a dipole in the nonpolar molecule.
- Polarizability of the nonpolar molecule: The more easily the nonpolar molecule’s electron cloud can be distorted, the stronger the induced dipole will be. Larger molecules with more electrons tend to be more polarizable.
iii. Example:
- Water (H₂O), a polar molecule, can induce a dipole in a nearby oxygen molecule (O₂), a nonpolar molecule, leading to a dipole-induced dipole interaction.
2. Dipole-Dipole Interactions:
i. What are they?
- These occur between two polar molecules. The positive end of one molecule attracts the negative end of another molecule, creating an attractive force.
ii. Key Points:
- Depend on dipole magnitudes: The strength of the dipole-dipole interaction depends on the magnitude of the dipoles. Larger dipoles lead to stronger interactions.
- Orientation matters: The relative orientation of the molecules also affects the interaction. If the molecules are aligned such that opposite charges face each other, the attraction will be stronger. If the molecules are misaligned, the attraction will be weaker.
- Stronger than London dispersion forces: Dipole-dipole interactions are generally stronger than the London dispersion forces between nonpolar molecules of similar size because dipoles interact more specifically and strongly.
iii. Example:
- In hydrogen chloride (HCl), the positive hydrogen end of one molecule interacts with the negative chloride end of another, creating dipole-dipole interactions.
3. Ion-Dipole Forces:
i. What are they?
- These interactions occur between ions (either cations or anions) and polar molecules. The charge of the ion interacts with the opposite partial charge of the polar molecule.
ii. Key Points:
- Stronger than dipole-dipole interactions: Ion-dipole forces tend to be stronger than dipole-dipole interactions because the ion has a full charge, which generates a much stronger attraction compared to the partial charges in dipoles.
- Significance in solutions: Ion-dipole forces are especially important in aqueous solutions or other polar solvents, where ions interact with solvent molecules.
iii. Example:
- In an aqueous solution of sodium chloride (NaCl), the sodium ions (Na⁺) are attracted to the partially negative oxygen atoms of water molecules, while the chloride ions (Cl⁻) are attracted to the partially positive hydrogen atoms of water molecules.
3.1.A.3 Strength and Orientation of Dipole and Ion-Dipole Forces:
1. Molecular Dipoles and Partial Charges:
i. Partial Charges (δ+ and δ-):
- Arise from differences in electronegativity (χ) between atoms in a covalent bond.
- More electronegative atom gains a partial negative charge (δ-), and the less electronegative atom gains a partial positive charge (δ+).
Example: In H₂O:
- Oxygen (δ-) and Hydrogen (δ+)
ii. Dipole Moment (μ):
- A dipole moment occurs when there is an unequal distribution of electron density in a molecule, resulting in a separation of charges.
- Formula: μ = δ × r, where:
- μ = dipole moment
- δ = magnitude of the partial charge
- r = distance between charges
Example: In H₂O, the dipole moment points from the hydrogen atoms (δ+) to the oxygen atom (δ-).
iii. Role of Dipole Moments:
- Dipole-Dipole Interactions: Polar molecules interact through attractive forces between their dipoles.
- Hydrogen Bonding: Occurs when a hydrogen atom bonded to an electronegative atom (like O, N, or F) interacts with a lone pair on another electronegative atom.
Example: In H₂O, hydrogen bonding occurs between the δ+ hydrogen of one molecule and the δ- oxygen of another.
iv. Nonpolar Molecules:
- No dipole moment arises when molecules are symmetric or have equal electronegativity between atoms (e.g., CO₂).
- Example: CO₂ (linear molecule, dipoles cancel out) has no overall dipole moment.
2. Dipole-Dipole Interactions:
Dipole-Dipole Interactions occur between polar molecules due to the attractive forces between the positive end (δ+) of one molecule and the negative end (δ-) of another. The strength of these interactions is influenced by the magnitude of the dipoles and their orientation relative to each other.
i. Magnitude of the Dipoles (δ and μ):
a. The strength of dipole-dipole interactions increases with the magnitude of the dipoles. A larger dipole moment (μ) means a greater charge separation and stronger attraction between molecules.
b. The dipole moment (μ) is calculated as:
where:
- μ = dipole moment
- δ = partial charge
- r = distance between charges (bond length).
- Stronger dipoles (larger δ) lead to stronger interactions. For example, molecules like hydrogen chloride (HCl) have a dipole due to the difference in electronegativity between hydrogen (δ+) and chlorine (δ-).
2. Orientation of the Dipoles:
The relative orientation of the dipoles is crucial in determining the strength of the interaction. The dipoles will interact most strongly when the positive end of one dipole is aligned with the negative end of the other dipole.
Most favorable orientation: When the molecules are aligned so that opposite charges (δ+ and δ-) are in close proximity, the attraction is maximized.
Less favorable orientation: If the dipoles are misaligned (e.g., with like charges facing each other), the interaction is weaker, and repulsive forces might even occur.
* Distance Between Molecules:
- The strength of dipole-dipole interactions also depends on the distance between the interacting molecules. As the distance between dipoles increases, the strength of the interaction decreases rapidly (following an inverse cube law).
where F is the force of interaction and r is the distance between dipoles.
iii. Examples:
Hydrogen Chloride (HCl): In HCl, the chlorine atom (δ-) attracts the hydrogen (δ+), resulting in dipole-dipole interactions between molecules. The stronger the dipole (due to a larger electronegativity difference), the stronger the interaction.
Water (H₂O): Water has a strong dipole moment due to its bent structure, leading to strong dipole-dipole interactions between water molecules. The partial positive hydrogen of one molecule is attracted to the partial negative oxygen of another.
3. Ion-Dipole Interactions:
Ion-Dipole Interactions occur when an ion (charged species) interacts with a polar molecule (molecule with a permanent dipole). These interactions are often stronger than dipole-dipole interactions and play a crucial role in the behavior of ions in polar solvents, such as water.
i. Mechanism of Attraction:
a. Attraction Between Opposite Charges:
- Ions carry a full charge (either positive or negative), while polar molecules have partial charges (δ+ and δ-).
- The positive ion (cation) is attracted to the partial negative charge (δ-) of the polar molecule, and the negative ion (anion) is attracted to the partial positive charge (δ+) of the polar molecule.
b. Orientation of the Dipole:
- The orientation of the dipole is critical in ion-dipole interactions. A polar molecule aligns such that the negative dipole end faces a positive ion (cation) and the positive dipole end faces a negative ion (anion).
- This alignment maximizes the attractive force, as opposite charges attract.
c. Strength of the Interaction:
- Ion charge: The greater the charge of the ion, the stronger the attraction between the ion and the polar molecule. A higher charge increases the electrostatic attraction between the ion and the opposite partial charge of the polar molecule.
- Ion size: Smaller ions with the same charge have a stronger ion-dipole interaction due to their higher charge density (charge-to-size ratio), which makes them more effective at polarizing the surrounding molecule.
- Dipole moment of the molecule: The larger the dipole moment (δ), the stronger the ion-dipole interaction, as a larger dipole can induce a stronger polarization of the ion’s charge.
ii. Example:
- Dissolution of Salt in Water:
- When NaCl (sodium chloride) dissolves in water, the Na⁺ cation interacts with the partial negative oxygen (δ-) of water molecules, while the Cl⁻ anion interacts with the partial positive hydrogen (δ+) of water molecules.
- The small size and high charge density of Na⁺ leads to a strong ion-dipole interaction with water molecules, facilitating its dissolution.
- Hydration of Ions:
- When ions like Na⁺ or Cl⁻ dissolve in water, they are surrounded by water molecules arranged in a way that maximizes ion-dipole interactions, creating a “hydration shell” around each ion.
iii. Formula for the Ion-Dipole Interaction:
The strength of the ion-dipole interaction can be roughly described by an equation like:
where:
- Q is the charge of the ion,
- μ is the dipole moment of the polar molecule,
- r is the distance between the ion and the center of the dipole.
4. Orientation Dependence:
Orientation Dependence plays a significant role in determining the strength of dipole-dipole interactions and ion-dipole interactions. The relative positioning of the dipoles or the ion and dipole affects the magnitude of the attractive forces between them.
i. Dipole-Dipole Interactions:
a. Dipole-dipole interactions occur between polar molecules. The positive end of one molecule (δ+) attracts the negative end (δ-) of another molecule, and vice versa.
b. Orientation Matters:
- Maximal Attraction: The interaction is strongest when the positive end (δ+) of one molecule is aligned with the negative end (δ-) of the other molecule. This is the most favorable alignment because opposite charges attract.
- Minimal Attraction or Repulsion: If the molecules are aligned such that like charges face each other (δ+ to δ+ or δ- to δ-), the interaction weakens and may even become repulsive.
Example:
- In hydrogen chloride (HCl), the δ+ hydrogen of one HCl molecule will attract the δ- chlorine of another HCl molecule, resulting in strong dipole-dipole attraction when aligned properly.
- If two HCl molecules were aligned such that δ+ was facing δ+ (or δ- to δ-), the repulsive forces would dominate.
ii. Ion-Dipole Interactions:
a. Ion-dipole interactions occur between an ion and a polar molecule. The ion interacts with the partial charges of the polar molecule, and its orientation relative to the ion significantly affects the strength of the interaction.
b. Orientation Matters:
- Maximal Attraction: For the cation (positive ion), the negative end (δ-) of the polar molecule should face the ion, and for the anion (negative ion), the positive end (δ+) of the polar molecule should face the ion.
- Suboptimal Orientation: If the ion and dipole are misaligned, the interaction becomes weaker as like charges (cation-δ+ or anion-δ-) are positioned too closely, leading to less effective attraction.
Example:
- In hydration of Na⁺ in water, the oxygen (δ-) atoms of water molecules surround the Na⁺ cation. The alignment of the water molecules maximizes the ion-dipole interaction, resulting in a strong attraction.
- Similarly, Cl⁻ is surrounded by the hydrogen (δ+) ends of the water molecules. If the water molecules were oriented incorrectly (e.g., δ- to Cl⁻), the attraction would be weaker.
iii. Effect of Distance and Angle:
- The strength of both dipole-dipole and ion-dipole interactions also depends on distance and angle. A more favorable angle of alignment (where opposite charges face each other) maximizes the interaction.
- Distance: As the distance between the molecules (or ion and dipole) increases, the interaction strength decreases. For example, in dipole-dipole interactions, the force falls off as 1/r³ (inverse cube law), and in ion-dipole interactions, the force decreases with increasing distance from the ion to the dipole.y.
3.1.A.4 Hydrogen Bonding: Strong Intermolecular Interaction Between H and N, O, F:
1. Intermolecular Forces:
Intermolecular forces (IMFs) are interactions between molecules that influence physical properties like boiling points and solubility
i. Hydrogen Bonding: Strong IMF between hydrogen atoms bonded to electronegative atoms (N, O, F) and lone pairs on other electronegative atoms. Example: Water.
ii. Dipole-Dipole Interactions: Attraction between the positive and negative ends of polar molecules. Example: Hydrogen chloride (HCl).
iii. Van der Waals Forces (London Dispersion): Weak forces caused by temporary dipoles in all molecules. Stronger in larger molecules. Example: Methane (CH₄).
iv. Ion-Dipole Interactions: Attraction between an ion and a polar molecule. Important in ionic solutions. Example: Sodium ions in water.
v. Dipole-Induced Dipole Interactions: Occur when a polar molecule induces a temporary dipole in a nonpolar molecule. Example: Polar HCl with nonpolar argon.
vi. Strength Order:
Hydrogen bonding > Ion-dipole > Dipole-dipole > Dipole-induced dipole > Van der Waals. These forces determine a substance’s properties like boiling and melting points.
2. Electronegativity and Dipoles:
Electronegativity is the tendency of an atom to attract shared electrons in a covalent bond. The larger the difference in electronegativity between two atoms, the more unevenly the electrons are shared, creating a dipole—a separation of charge.
i. Electronegativity Differences and Dipoles:
a. Nitrogen (N), Oxygen (O), Fluorine (F):
- High Electronegativity: N, O, and F are highly electronegative (N: 3.04, O: 3.44, F: 3.98 on the Pauling scale). These atoms attract electrons strongly in covalent bonds.
- Resulting Dipole: When bonded with hydrogen (electronegativity of 2.20), N, O, or F will pull electron density towards themselves, creating a partial negative charge (δ-) on the electronegative atom and a partial positive charge (δ+) on the hydrogen atom.
b. Hydrogen:
- Lower Electronegativity: Hydrogen is less electronegative (2.20) compared to N, O, and F. When bonded to these elements, it has a partial positive charge due to its weaker attraction for the shared electrons.
ii. Examples of Dipoles:
a. NH₃ (Ammonia): Nitrogen pulls electron density towards itself, creating a dipole where nitrogen is δ- and hydrogen is δ+.
b. H₂O (Water): Oxygen’s high electronegativity creates a dipole where oxygen is δ- and hydrogen is δ+.
c. HF (Hydrogen fluoride): Fluorine, being the most electronegative element, strongly attracts electrons, creating a large dipole with δ- on F and δ+ on H.
3. Hydrogen Bonding Mechanism:
Hydrogen bonding occurs when a hydrogen atom, covalently bonded to a highly electronegative atom (N, O, or F), is attracted to the lone pair of electrons on another electronegative atom.
i. Polar Bond: The electronegative atom (N, O, or F) pulls electron density from hydrogen, creating a partial positive charge (δ+) on hydrogen and a partial negative charge (δ-) on the electronegative atom.
ii. Attraction to Lone Pair: The δ+ hydrogen is attracted to the lone pair of electrons on another electronegative atom, forming a hydrogen bond.
iii. Examples:
- Water (H₂O): Hydrogen bonds between the hydrogen of one molecule and the lone pair on oxygen of another.
- Ammonia (NH₃): Hydrogen bonds between hydrogen and lone pairs on nitrogen.
- Hydrogen Fluoride (HF): Strong hydrogen bonds due to fluorine’s high electronegativity.
4. Role in Water and Biological Molecules:
i. High Boiling and Melting Points: Hydrogen bonds between water molecules require significant energy to break, resulting in water’s relatively high boiling and melting points.
ii. Cohesion and Surface Tension: Water molecules are attracted to each other through hydrogen bonding, creating high cohesion and surface tension, which allows insects to “walk” on water.
iii. Solvent Properties: Water’s hydrogen bonding allows it to dissolve many polar substances (like salts and sugars), making it an excellent solvent.
iv. High Specific Heat and Heat of Vaporization: Hydrogen bonds in water absorb a lot of heat before breaking, which helps regulate temperature in living organisms and environments.
v. Hydrogen Bonding in Biological Molecules:
a. DNA: Hydrogen bonds between complementary nitrogenous bases (adenine-thymine, cytosine-guanine) stabilize the double helix structure of DNA.
b. Proteins: Hydrogen bonds contribute to the secondary and tertiary structures of proteins (e.g., alpha helices and beta sheets), influencing protein folding and stability.
c. Enzyme-Substrate Interactions: Hydrogen bonding between enzymes and substrates helps stabilize the enzyme-substrate complex, crucial for catalysis.
5. Factors Affecting Strength:
i. Electronegativity of the Atoms:
- Electronegativity is the ability of an atom to attract electrons in a covalent bond. A larger difference in electronegativity between the hydrogen atom and the electronegative atom (e.g., N, O, F) leads to a stronger hydrogen bond.
- Example: Fluorine (3.98) is the most electronegative element. When bonded to hydrogen (2.20), the large electronegativity difference creates a strong dipole, resulting in a very strong hydrogen bond. On the other hand, oxygen (3.44) creates a weaker hydrogen bond with hydrogen in water because the electronegativity difference is smaller.
ii. Atom Size:
- The size of the atoms involved in the hydrogen bond significantly affects its strength. Smaller atoms like fluorine, oxygen, and nitrogen allow closer proximity of the hydrogen bond donor and acceptor, resulting in a stronger bond.
- Larger atoms (e.g., sulfur or iodine) have more diffuse electron clouds, making the hydrogen bond weaker due to greater distances between the hydrogen atom and the electronegative atom. The larger the atoms, the weaker the interaction because the donor and acceptor cannot approach each other as closely.
iii. Temperature:
- Higher temperature provides more kinetic energy, which can disrupt intermolecular forces, including hydrogen bonds. As the temperature rises, molecules move more vigorously, which can weaken or break hydrogen bonds.
- Lower temperatures reduce molecular motion, allowing hydrogen bonds to form more easily and remain intact. This is why water, for example, remains liquid at temperatures below 100°C due to the presence of hydrogen bonds, but as temperature increases, these bonds start breaking, leading to water vapor (gas phase).
iv. Presence of Solvents:
- Polar solvents, such as water, can compete with hydrogen bonds for interaction with the hydrogen-bond donor or acceptor. Water, being a polar solvent itself, can disrupt hydrogen bonding between other molecules by forming its own hydrogen bonds with the donor or acceptor molecules. This reduces the strength of hydrogen bonding between the solute molecules.
- Nonpolar solvents, however, generally do not disrupt hydrogen bonding much because they don’t engage in hydrogen bonding themselves. The hydrogen bonds between solute molecules remain largely intact in nonpolar solvents.
v. Bond Geometry:
- The geometry of the hydrogen bond plays a significant role in determining its strength. The hydrogen bond is strongest when the donor hydrogen atom, the acceptor atom (e.g., N, O, F), and the lone pair of the acceptor atom are aligned in a straight line (linear geometry). This alignment allows for the optimal overlap of electron clouds, maximizing the strength of the bond.
- Nonlinear or bent geometries (such as when the hydrogen bond is at an angle) result in weaker hydrogen bonds because the overlap between the donor and acceptor electron clouds is reduced, making the bond less effective.
3.1.A.5 Noncovalent Interactions in Large Biomolecules:
1. Types of Noncovalent Interactions:
| Interaction Type | Definition | Characteristics | Example |
|---|---|---|---|
| Hydrogen Bonds | Attraction between a hydrogen atom bonded to an electronegative atom (N, O, F) and a lone pair on another electronegative atom. | Moderate strength, critical for biomolecular structure, directional. | Water (H₂O), DNA base pairing, protein folding. |
| Ionic Interactions | Attraction between positively and negatively charged ions or groups. | Stronger than hydrogen bonds, disrupted in polar solvents, involve full charge separation. | Salt (Na⁺ and Cl⁻), amino acid side chains in proteins. |
| Van der Waals Forces | Weak, short-range interactions due to transient dipoles caused by electron fluctuations. | Very weak, significant in large molecules, arise from instantaneous or induced dipoles. | Stabilizing protein structure, nonpolar molecule interactions (e.g., between hydrocarbons). |
| Hydrophobic Interactions | Nonpolar molecules or regions cluster to avoid water (or other polar solvents). | Driven by entropy, important for protein folding, membrane formation. | Protein folding (hydrophobic core), lipid bilayers. |
2. Intramolecular and Intermolecular Interactions:
i. Intramolecular Interactions:
- Definition: Interactions that occur within a single molecule, where different parts of the same biomolecule interact with each other.
- Role: They help to stabilize the 3D structure of a biomolecule, enabling it to perform its biological function.
- Examples:
- Protein folding: Intramolecular hydrogen bonds, ionic interactions, and hydrophobic interactions between different regions of the same protein help it fold into a functional 3D shape.
- DNA structure: Within a single strand of DNA, hydrogen bonds between nitrogenous bases (A-T, G-C) hold the base pairs together. The overall DNA structure is stabilized by interactions like hydrogen bonds and hydrophobic stacking between the bases.
- Enzyme active sites: Enzymes have specific intramolecular interactions within their active sites that help them bind substrates and catalyze reactions. These interactions often involve hydrogen bonds, van der Waals forces, and ionic interactions.
ii. Intermolecular Interactions:
- Definition: Interactions that occur between different molecules or different parts of molecules (such as different biomolecules) in a biological system.
- Role: They are crucial for the function of biomolecular complexes, signal transduction, molecular recognition, and cellular interactions.
- Examples:
- Protein-protein interactions: When two or more proteins interact, such as in enzyme-substrate binding or the formation of protein complexes, intermolecular interactions (hydrogen bonds, ionic interactions, van der Waals forces) are responsible for the specificity and strength of these interactions.
- DNA-protein interactions: Transcription factors bind to specific sequences on DNA through hydrogen bonds and ionic interactions, enabling gene expression regulation.
- Antigen-antibody interactions: The binding of an antibody to its specific antigen is driven by intermolecular forces, including hydrogen bonds, ionic interactions, and van der Waals forces, leading to immune responses.
- Receptor-ligand binding: The interaction between cell receptors and signaling molecules (ligands) is governed by intermolecular forces, allowing for processes like hormone signaling and neurotransmitter reception.
iii.Key Differences:
- Intramolecular interactions occur within a single molecule and are crucial for maintaining the structure and function of that molecule.
- Intermolecular interactions occur between different molecules and are essential for biological processes such as signaling, molecular recognition, and complex formation.
Old Content
Intramolecular vs Intermolecular
- Intramolecular bonding: what holds atoms together within the molecule
- Covalent and ionic
- Intermolecular forces: forces between (rather than within) molecules
Intermolecular Forces
- Weaker interactions that occur between molecules (3 types)
- Hydrogen bonding (strongest)
- Dipole-dipole forces (intermediate)
- London-dispersion forces (weakest)
- Intramolecular forces are a lot stronger than intermolecular forces
- The weaker the forces, the more likely the substance is to exist as a gas → bcuz particles are able to move far apart since they are not held together very strongly
- Molecules orient themselves to maximize the ⊕,⊝ interactions and to minimize ⊕,⊕ and ⊝,⊝ interactions
Dipole-Dipole Forces
- Dipole-Dipole Attraction: Molecules with dipole moments (polar molecules) can attract each other electrostatically by lining up so that the positive and negative ends are close to each other
- Can be attractive or repulsive
- The strength of the interactions is directly related to the magnitude of the dipole
- The more polar the molecule is (bigger electronegativity difference) → stronger the D-D attraction
- Force becomes weaker as the distance between the dipoles increases (larger atoms)
- Molecules with D-D have atoms held more tightly together → higher MP and BP → more likely to be a liquid at room temp
- The more polar the molecule is (bigger electronegativity difference) → stronger the D-D attraction
Hydrogen “Bonding”
- Hydrogen bonds: A very strong type of dipole-dipole attraction that occurs when hydrogen is covalently bonded (intramolecular) to N, O, or F
- → not HB bcuz H is covalently bonded to Carbon
- It is very strong due to the large difference in electronegativity and hydrogen’s small size that allows dipoles to be close → need more energy to overcome bonds → higher MP and BP
London Dispersion Forces
- ALL covalent compounds experience LDF (aka van der waal forces)
- Do not use the word “has”
- Is the only force present in noble gas atoms and nonpolar molecules
- Forces also exist between polar molecules but not as strong as D-D
- In non-polar molecules the electron charge is usually evenly distributed but it is possible that at a particular moment in time, the electrons might not be evenly distributed (e- are always moving in their orbitals) → distorts electron cloud → separation of charge creates instantaneous dipoles in molecule → induces another instantaneous dipole in another molecule → short/weak attraction between molecules
- Polarizability: how easy the electron cloud of an atom can be induced into a dipole
- Elements with more electrons (greater molar mass) = more polarizable molecule = electron cloud can be more easily induced into a dipole = stronger LDF = higher MP/BP
- Force is usually weak but can be significant in large atoms/molecules (Ex: F2, Cl2, Br2)
Ion-Dipole Forces
- Attractive force between an ion and the oppositely charged end of a polar molecule → strongest force
- Ex: NaCl and water
Forces Between Polar and Nonpolar Molecules
- Dipole-induced dipole interaction: a permanent dipole on a polar molecule can induce a dipole on a neighboring nonpolar molecule
- Thus polar molecules and nonpolar molecules exhibit an attraction for one another
- Explains why oxygen and CO2 can be dissolved (slightly) in water
- Thus polar molecules and nonpolar molecules exhibit an attraction for one another
- Factors that affect strength → nature of both molecules
- The larger the magnitude of the dipole in a polar molecule → the better able it is to induce a dipole in a neighboring molecule.
- Nonpolar molecules with a greater number of electrons have an increased polarizability (LDF) → dipole more easily induced
Factors Affected by IMF’s
Properties that increase with stronger InterMF’s → bcuz molecules experience stronger attractions for each other and it takes more energy to overcome that attraction
- Melting/Freezing and boiling points
- Surface tension
- Heat of vaporization
- Viscosity
Properties that decrease with stronger InterMF’s
- Vapor pressure
- Volatility: how easily a liquid evaporates
Comparing IMF’s
- If the size of molecules is similar → H-bonding > D-D > LDF
- If sizes are comparable → LDFs are similar bcuz have electron clouds of similar size and polarizability
- If the molecules are of different size → more complex → very imp. for AP exam
- LDF’s can become more significant than D-D or H-B when the size of molecules is bigger
- Occurs when the difference in magnitude of LDFs in A exceed the magnitude of all the other interactions in B, giving A more IMFs total
- LDF’s can become more significant than D-D or H-B when the size of molecules is bigger
- More linear compounds (less branched) = greater surface area for contact between molecules = stronger InterMFs
- Ex: for questions involving isomers