phases of matter and IMFs - dec 4

A chemical system exists in a specific physical state due to:
1. Thermal Energy: How much thermal energy does the system have? (i.e., high or low temperature) —> higher thermal energy = greater ability to overcome IMFs, resulting in a liquid state for water at room temperature. Additionally, the unique hydrogen bonding in water contributes to its relatively high boiling point despite its molecular weight.
2. Intermolecular Forces (IMFs): How strong are the intermolecular forces holding that system together? attractive forces must be overcome for a substance to change phase
Behavior Based on Thermal Energy:
- High thermal energy = gas —> IMFs are overcome
- Low thermal energy = liquid or solid —> IMFs are not overcome
Phase Relationship: At a given temperature (constant thermal energy present), species with:
- Very strong IMFs are likely to be solids
- Medium strength IMFs yield liquids
- Weak IMFs result in gases

Thermal Energy: Low relative to IMF strength
Density: High (g/mL), particles are closely packed
Shape & Volume: Fixed shape and volume
Kinetic Energy (KE): Low, very low freedom of motion
Structure: Can be crystalline (ordered) or amorphous (random)
Compressibility: Almost perfectly incompressible —> no space between particles

Thermal Energy: Moderate relative to IMF strength
Density: Medium (g/mL)
Shape: Indefinite shape
Volume: Fixed volume
Kinetic Energy (KE): More KE, more freedom of motion - rotational, translational, and vibrational motion
Compressibility: Very low compressibility because particles are close together

Thermal Energy: High relative to IMF strength
Density: Low (g/L) —> around 1000x smaller than liquids or solids
Shape & Volume: Variable shape and variable volume - particles move randomly + rapidly to fill the container that they are in
Kinetic Energy (KE): Very high, most freedom of motion
Compressibility: Very high compressibility

IMFs are not the same as chemical bonds. —> intermolecular rather than intramolecular (within the molecule itself)
Phase Changes occur when IMFs are broken or overcome:
- Heat a solid to melt it (“fusion”)
- Heat a liquid to boil it (“vaporization”)
- Cool a gas to condense it to a liquid
- Cool a liquid to freeze it to a solid
- Also, one can lower pressure to vaporize a liquid or increase pressure to condense a gas!
Key Point: No chemical bonds are broken when phase changes occur! For example, when turning liquid water to gas, no O-H covalent bonds are broken, the hydrogen bonds between water molecules are broken instead
Strength Comparison: The strength of primary chemical bonds (ionic, metallic, covalent) is much greater than that of IMFs.
Reaction Examples:
- For water: $H<em>2O(l) \rightarrow H</em>2O(g) \quad \Delta H_{rxn} = 40.7 \, kJ \text{ at } 100^{\circ}C$ [phase change]
- Breaking all bonds: $H<em>2O(g) \rightarrow 2H(g) + O(g) \quad \Delta H</em>{rxn} \sim 1000 \, kJ \text{ at } 100^{\circ}C$

All bond formation results from Coulombic interactions:
- Maximizing attractive forces between electrons and protons
- Minimizing repulsive forces between like charges [proton/proton and electron/electron]
- Optimizing ion attraction/repulsion in solids
Coulomb's Law: Describes how forces are directly proportional to the charges of particles and inversely proportional to the distance between them.
Types of Bonds:
- Primary chemical bonding involves particles with whole, permanent charges across short distances = strong Coulombic forces.
- IMFs arise from partial/temporary buildup of charges across longer distances = weaker Coulombic forces.

Chemical Species: Present in all species, both polar and nonpolar
Origin: Caused by temporary/instantaneous fluctuations in the location of electrons in orbitals. One side becomes slightly negative while the other side becomes slightly positive, causing attractive forces between a positive and a negative end of neighboring molecules, leading to the formation of temporary dipoles.
Relative Strength: The weakest of all IMFs.
Comparative Strength Factors: Depends on molar mass and shape of the molecule. —> large molecules with more electrons are more polarizable, electron clouds can be distorted more easily which means they are more likely to form instantaneous dipoles. Additionally, greater surface area means stronger LDFs because there is more contact available.

Chemical Species: Present in polar molecules exclusively.
Origin: Results from a permanent uneven distribution of electrons within molecules.
Relative Strength: Medium strength compared to other types.
Comparative Strength Factors: Related to size, shape, and electronegative atoms leading to a “dipole moment”, measured in Debye (D). For example:
- 1.05 D
- 1.224 D
- 1.42 D

Chemical Species: Found in molecules with hydrogen bonded directly to nitrogen (N), oxygen (O), or fluorine (F) atoms.
Origin: Results from the strong dipole due to the high electronegativity of the aforementioned atoms.
Relative Strength: The strongest IMF for a pure species. The greater number of sites available for hydrogen bonding the stronger the IMFs are.
- ethanol has a higher boiling point than dimethyl ether because it has an -OH group which means hydrogen bonding occurs
Comparative Strength Factors: Generally assessed by the number of –OH and –NH2 groups in an organic molecule. (Note: H-F is very polar but typically exists only by itself)

Chemical Species: Exists in ionic compounds dissolved in polar solvents.
Origin: Strong attraction of the partial charges in polar molecules to the permanent, whole charges on ions.
Relative Strength: The strongest overall type of IMF.
Comparative Strength Factors: Depends on the dipole moment of the polar solvent and charge density of ions (largest charge/smallest size).

Thermal Energy Requirement: A stronger IMF requires more thermal energy to overcome, leading to greater stability of the substance.
Stability Implication: More stable substances exhibit higher melting and boiling point temperatures.
Chemical Compatibility: Species that rely on strong IMFs are generally chemically incompatible with those that can only use weaker ones.
- Example: Short-chain hydrocarbons (which exhibit only dispersion forces) do not interact with water (which has hydrogen bonding) because such interaction would decrease the overall attractive forces among water molecules.
- General Rule: “Like dissolves like”.
Strength Comparison: When comparing substances:
- Ionic/metallic/covalent bonds remain significantly stronger than IMFs.
- Example: Salt has an extremely high melting point, whereas metals have very high melting points; covalent molecule melting points are comparatively lower.
Boiling Point Factors: When determining which of several species possesses the highest boiling point:
- The species with the strongest possible IMF generally has the highest boiling point. If all species can utilize the same type of IMF, refer to specific guidelines for that type to identify the strongest.