Molecular Forces and States of Matter Practice Flashcards

Arrangement and Movement of Matter States

• Solids

  • Particle Arrangement: Particles are tightly packed together, often in a regular, repeating pattern.

  • Particle Movement: Particles vibrate in fixed positions. They do not move in relation to one another, meaning they stay put but undergo constant shaking and vibration. They cannot shake past one another.

• Liquids

  • Particle Arrangement: Particles are close together but lack a regular arrangement. Generally, they are spaced out slightly further than particles in a solid.

  • Particle Movement: Particles can move past each other but remain essentially in contact at all times. While they can flow, they are always touching at least one other particle or molecule in the sample.

• Gases

  • Particle Arrangement: Particles are positioned far apart with no regular arrangement or pattern whatsoever.

  • Particle Movement: Particles move independently of each other. Interaction occurs only during collisions. When collisions happen, it is assumed that particles hit and bounce off with no attractive forces holding them together.

Intermolecular Forces vs. Kinetic Energy

• Relative Extents of Forces: The phase in which a substance exists and its specific physical properties depend on the competition between intermolecular forces and kinetic energy.

  • Intermolecular Forces (IMFs): Defined as various forces of attraction that exist between atoms and molecules of a substance due to electrostatic phenomena. Their role is to hold particles close together.

  • Kinetic Energy: The energy associated with movement. Its role is to move particles further apart. This energy increases directly with temperature.

• Phase Transitions and Temperature

  • As temperature increases, kinetic energy increases, causing a substance to transition from solid to liquid to gas.

  • As temperature increases, the relative influence of attractive forces decreases because kinetic energy overcomes them.

  • As temperature decreases, the influence of intermolecular forces increases. Particles move less, become more closely packed, and eventually transition from gas to liquid to solid.

• Influencing Variables: Both temperature and pressure affect the balance between intermolecular forces and kinetic energy.

Intramolecular vs. Intermolecular Forces

• Intramolecular Forces

  • These are forces within the molecule itself that keep the molecule together (i.e., chemical bonds).

  • They occur from atom to atom (e.g., the covalent bond between $H$ and $Cl$ in a hydrogen chloride molecule).

  • These are generally much stronger than intermolecular forces.

• Intermolecular Forces (IMFs)

  • These are forces of attraction between separate and distinct molecules (e.g., Molecule A attracted to Molecule B).

  • They determine many of the physical properties of a substance.

  • These are represented in diagrams by dotted lines to distinguish them from the solid lines used for intramolecular bonds.

  • These are generally weaker than the bonds within a molecule.

Van der Waals Forces: Dispersion Forces

• Overview of Van der Waals Forces: These are the attractive forces between neutral atoms and molecules. They include three specific types: dispersion forces, dipole-dipole attractions, and hydrogen bonding.

• Dispersion Forces (London Forces)

  • Historical Context: First explained in 1928 by the scientist Fritz London.

  • Definition: Weak attractive forces resulting from the constant movement of electrons, leading to rapidly fluctuating temporary dipoles.

  • Applicability: These exist between all particles, regardless of size, mass, or polarity (both polar and nonpolar molecules), because all atoms/molecules contain electrons.

• Mechanism of Dispersion Forces

  • Instantaneous Dipole: Electrons are in constant motion. At any given fraction of a second, electrons might be distributed asymmetrically (more on one side than the other). This creates a very fast, temporary separation of charge.

  • Induced Dipole: When an atom develops an instantaneous dipole, it distorts the electron cloud of a neighboring atom. The neighbor's electrons respond to the temporary charge, creating their own temporary dipole.

  • Persistence: While these dipoles are temporary and switch "on and off" as electrons move, the process repeats continuously across a sample, maintaining an indefinite attraction.

• Polarizability

  • Definition: A measure of how easy or difficult it is for an external electrostatic charge to distort a molecule's electron cloud or charge distribution.

  • Factors Increasing Strength: Dispersion forces increase with higher proximity and larger atomic or molecular size.

  • High Polarizability: Molecules with large molar masses have more electrons. More electrons lead to a higher probability of asymmetric distribution and larger separations of charge. These are "easily distorted" and result in stronger IMFs.

  • Low Polarizability: Smaller molecules with lower molar masses have fewer electrons and less space for them to spread out. They are "difficult to distort," leading to smaller dispersion forces.

  • Physical Consequences: Stronger dispersion forces lead to higher melting points and higher boiling points because more energy is required to break the attractions between molecules.

Trends in Dispersion Forces and Boiling Points

• Halogen Series (Group 17) Data

  • Fluorine ($F_2$): Molar Mass: roughly $38.00\,g/mol$. Atomic Radius: $72\,pm$.

  • Chlorine ($Cl_2$): Molar Mass: roughly $70.90\,g/mol$. Atomic Radius: $99\,pm$.

  • Bromine ($Br_2$): Molar Mass: roughly $159.80\,g/mol$. Atomic Radius: $114\,pm$.

  • Iodine ($I_2$): Molar Mass: roughly $253.80\,g/mol$. Atomic Radius: $133\,pm$.

  • Astatine ($At_2$): Molar Mass: roughly $420\,g/mol$. Atomic Radius: $150\,pm$.

  • Observations: As you move down the column, molar mass and atomic radius increase. Consequently, polarizability increases, leading to stronger IMFs and progressively higher melting and boiling points.

• Group 14 Hydrides Comparison (Lowest to Highest Boiling Point)

  1. Methane ($CH_4$): Lowest molar mass (approx. $16\,g/mol$); weakest dispersion forces; lowest boiling point.

  2. Silicon tetrahydride ($SiH_4$): Molar mass approx. $32\,g/mol$.

  3. Germanium tetrahydride ($GeH_4$): Molar mass approx. $77\,g/mol$.

  4. Tin(IV) hydride ($SnH_4$): Highest molar mass (approx. $123\,g/mol$); strongest dispersion forces; highest boiling point.

Dipole-Dipole Attractions

• Dipole Definition: A permanent separation of charge due to the unequal distribution of electrons. In polar molecules, one atom has a greater electronegativity (attraction for electrons) than the other, resulting in a persistent partial negative charge ($\delta^-$) on the "electron-hogging" atom and a partial positive charge ($\delta^+$) on the other.

• Attraction Mechanism: The electrostatic force between the partially positive end of one polar molecule and the partially negative end of another.

  • Directionality: Attraction can occur in one direction (end-to-end) or multiple directions (side-by-side).

  • Scope: In a real sample, these attractions are felt between millions of molecules in close proximity, not just between pairs.

• Strength Comparison: Dipole-dipole attractions are generally stronger than dispersion forces because the dipoles are permanent rather than instantaneous and fluctuating.

• Case Study: Hydrogen Chloride ($HCl$) vs. Fluorine ($F_2$)

  • Similarities: Both have two atoms, one bond, and $18$ electrons. They have approximately the same molecular mass.

  • Differences: $F_2$ is nonpolar; $HCl$ is polar.

  • Boiling Points: $HCl = 188\,K$; $F_2 = 85\,K$.

  • Conclusion: The higher boiling point of $HCl$ indicates stronger IMFs due to dipole-dipole attractions, whereas $F_2$ only experiences weaker dispersion forces.

• Case Study: Nitrogen ($N_2$) vs. Carbon Monoxide ($CO$)

  • Structures: Both have triple bonds and two lone pairs. Both have a molar mass of approx. $28\,g/mol$.

  • Nature of Bonds: $N_2$ is nonpolar because both atoms have equal electronegativity. $CO$ is polar because oxygen is more electronegative than carbon, creating a permanent dipole.

  • Conclusion: Even though they have similar London forces (due to similar mass), $CO$ has a higher boiling point because it also possesses dipole-dipole attractions.

Hydrogen Bonding

• Definition: A particularly strong type of dipole-dipole attraction that occurs when exceptionally strong dipoles attract. It is NOT the covalent bond within the molecule, but the attraction between molecules.

• Criteria for Hydrogen Bonding: A molecule must contain an intramolecular covalent bond between hydrogen ($H$) and one of the three most electronegative atoms:

  1. Fluorine ($F$)

  2. Oxygen ($O$)

  3. Nitrogen ($N$)

• Mechanism: The highly electronegative atom ($F$, $O$, or $N$) pulls electrons away from the small hydrogen atom, creating a highly concentrated partial positive charge on the hydrogen and a partial negative charge on the electronegative atom. The hydrogen bond is the attraction between that partial positive $H$ of one molecule and the partial negative $F$, $O$, or $N$ of a neighboring molecule.

• The Importance of Water ($H_2O$)

  • Water molecules have two $O-H$ bonds.

  • Each molecule can engage in hydrogen bonding in multiple directions (interacting with three or four neighboring molecules).

  • This strong network of hydrogen bonds gives water unique properties essential for life.

• Boiling Point Anomalies: When plotting boiling points of hydrides in Groups 15, 16, and 17, the trends generally increase with molar mass (e.g., $H_2S \rightarrow H_2Se \rightarrow H_2Te$). However, the lightest members—$H_2O$, $HF$, and $NH_3$—exhibit drastically higher boiling points than predicted (e.g., water at $100\,^\circ C$). This "jump" is explained by the presence of strong hydrogen bonding, which is absent in the heavier compounds of those groups.

Comprehensive Molecule Comparison

• Experiment: Comparing Dimethyl Ether, Ethanol, and Propane

  • Dimethyl Ether ($CH_3OCH_3$): Molar mass $46\,g/mole$. Slightly polar (bent shape), but lacks $O-H$ bonds. IMFs: Dispersion and dipole-dipole.

  • Ethanol ($CH_3CH_2OH$): Molar mass $46\,g/mole$. Highly polar with an $O-H$ group. IMFs: Dispersion, dipole-dipole, and strong hydrogen bonding.

  • Propane ($CH_3CH_2CH_3$): Molar mass $44\,g/mole$. Symmetrical hydrocarbon; nonpolar. IMFs: Only dispersion forces.

• Matching Boiling Points

  • Propane: $-42.1\,^\circ C$ (Lowest boiling point due to only having dispersion forces).

  • Dimethyl Ether: $-24.8\,^\circ C$ (Intermediate boiling point; polar but no hydrogen bonding).

  • Ethanol: $78.4\,^\circ C$ (Highest boiling point due to strong hydrogen bonding).

Summary Hierarchy of Bond Strengths

• General Rule: Intramolecular forces (covalent bonds) are stronger than intermolecular forces (Van der Waals forces).

• Van der Waals Strength Hierarchy (Strongest to Weakest):

  1. Hydrogen Bonding

  2. Dipole-Dipole Attractions

  3. Dispersion (London) Forces