Liquid phase

Hydrogen bonding

In the liquid phase, molecules are packed close together. Because of that, intermolecular forces strongly affect properties like boiling point and viscosity. One especially important intermolecular force is hydrogen bonding, a strong, specific type of dipole-dipole interaction.

In hydrogen bonding, a partially positive hydrogen bond donor (a hydrogen atom covalently bonded to a highly electronegative atom such as fluorine, oxygen, or nitrogen) is attracted to a partially negative hydrogen bond acceptor on a nearby molecule. Because hydrogen bonds take extra energy to break, substances with hydrogen bonding tend to have higher boiling points.

Hydrogen bond strength depends on bond polarity, with:

$H - F$ bonds being the strongest
- followed by $H - O$ bonds
- then $H - N$ bonds.

For example, ethers can’t form hydrogen bonds with one another because they don’t have a hydrogen atom attached to F, O, or N. Without that hydrogen, they lack the donor needed for hydrogen bonding.

Dipole interactions

Dipole-dipole interactions occur between polar molecules. The molecules tend to align so that the positive end of one molecule is near the negative end of another, creating an attractive force. This attraction raises the boiling point, although it’s generally weaker than hydrogen bonding. Stronger molecular polarity leads to stronger dipole-dipole interactions.

In contrast, ion-dipole interactions involve a charged ion interacting with a polar molecule. Because an ion carries a full charge (not just a partial charge), ion-dipole interactions are stronger than dipole-dipole forces. Their strength increases when:

the ion has a larger magnitude of charge, and
the dipole is more polar.

Van der Waals’ forces (London dispersion forces)

Van der Waals’ forces, also called London dispersion forces, are dispersion forces present in all molecules. They matter most in non-polar molecules, where other intermolecular attractions (like dipole-dipole interactions) are weak or absent.

These forces come from temporary dipoles:

Induced dipoles form when a polar molecule causes a nearby non-polar molecule to temporarily develop a dipole.
Instantaneous dipoles form when electron density in a non-polar molecule fluctuates momentarily, creating a brief dipole.

Dispersion forces get stronger as molecular size increases. For example, decane ( $C_{10} H_{22}$ ) has stronger dispersion forces than ethane ( $C_{2} H_{6}$ ).

Changes in phase

Interpreting phase changes in a phase diagram

A phase diagram shows which phase - solid, liquid, or gas - is stable at different temperatures and pressures.

In the solid state, molecules vibrate around fixed positions. Solids are nearly incompressible and don’t flow.
In the liquid state, molecules are still close together because of intermolecular forces, but they can move past one another. Liquids flow, take the shape of their container, and remain largely incompressible.
In the gas state, molecules move randomly and are far apart. Gases are easily compressible and expand to fill their container.

The diagram also includes phase boundaries, where two phases coexist in equilibrium:

The solid-liquid boundary represents equilibrium between solids and liquids.
The solid-gas boundary represents equilibrium between solids and gases.
The liquid-gas boundary represents equilibrium between liquids and gases.

The triple point is the specific combination of temperature and pressure at which all three phases exist simultaneously The critical point marks the conditions where the liquid and gas phases become indistinguishable The critical temperature is the highest temperature at which a liquid can exist.

Notably, water’s phase diagram is unusual because its solid-liquid boundary slopes to the left. This reflects that liquid water is denser than ice, so increasing pressure can convert ice into liquid water.

Common mnemonic to remember the gas region: “gas comes out this way.”

Freezing, melting, boiling, and sublimation

Freezing point is the temperature at which a liquid starts to turn into a solid.
Melting point is the temperature at which a solid begins to become a liquid. These occur along the solid-liquid boundary.
The boiling point is the temperature at which a liquid transitions into a gas
The condensation point is the temperature at which a gas starts to form a liquid. Both occur along the liquid-gas boundary.
Sublimation describes the direct change from a solid to a gas, occurring under conditions found along the solid-gas boundary.
Deposition describes the direct change from a gas to a solid

Colligative properties are properties of a solution that depend only on the number of solute particles present, not on what those particles are. Adding solute particles helps stabilize the liquid phase:

it becomes harder for the liquid to boil (so the boiling point increases), and
it becomes harder for the liquid to freeze (so the freezing point decreases).

Another way to say this is that lowering a solution’s vapor pressure corresponds to raising its boiling point.

A key idea in colligative properties is the van’t Hoff factor (i), which counts the total number of particles a solute produces in solution. For example, glucose does not dissociate, so its van’t Hoff factor is 1. In contrast, $N a Cl$ separates into $N a^{+}$ and $C l^{-}$ in solution, so it has a van’t Hoff’t factor of 2.

Vapor pressure lowering is described by Raoult’s law, which states that the vapor pressure of a solution equals the mole fraction of the solvent $(χ_{solvent})$ times the vapor pressure of the pure solvent $(P_{solvent}^{\circ})$ :

$P = χ_{solvent} \cdot P_{solvent}^{\circ} .$

The decrease in vapor pressure $(Δ P)$ can be written as:

$Δ P = χ_{solute} \cdot P_{solvent}^{\circ},$

where $χ_{solute}$ is the mole fraction of the solute.

Boiling point elevation is quantified by:

$Δ T_{b} = k_{b} \cdot m \cdot i$

where $Δ T_{b}$ is the increase in boiling point, $k_{b}$ is the molal boiling point constant, $m$ is the molality (moles of solute per kilogram of solvent), and $i$ is the van’t Hoff factor.

Similarly, freezing point depression is described by:

$Δ T_{f} = - k_{f} \cdot m \cdot i$

where $Δ T_{f}$ is the decrease in freezing point (the negative sign indicates a reduction), $k_{f}$ is the molal freezing point constant, $m$ is the molality, and $i$ is the van’t Hoff factor.

Osmotic pressure ( $π$ ) is another colligative property. It’s the pressure required to stop solvent from moving across a semi-permeable membrane from a region of low solute concentration to a region of high solute concentration. It is given by:

$π = MRT \cdot i$

where M is the molarity of the solution (in mol/L), R is the ideal gas constant, T is the absolute temperature in Kelvin, and $i$ is the van’t Hoff factor. Osmotic pressure determines the direction and extent of osmosis.

Differentiating types of mixtures:

A solution is a homogeneous mixture in which substances are combined at the molecular level and remain uniformly dispersed.
Colloids consist of solute aggregates that are extremely small and remain mixed unless separated by processes like centrifugation.
A suspension is a heterogeneous mixture where the particles are large enough that they do not stay evenly distributed over time.
A common example of a colloid is milk, and vigorous shaking of water and oil can produce an emulsion, which is also classified as a colloid.

Hydrogen bonding

Strong dipole-dipole interaction; requires H bonded to F, O, or N
Increases boiling point and viscosity
Strength order: H-F > H-O > H-N

Dipole interactions

Dipole-dipole: between polar molecules; weaker than hydrogen bonds
Ion-dipole: between ions and polar molecules; stronger than dipole-dipole
- Strength increases with higher ion charge and greater dipole polarity

Van der Waals’ forces (London dispersion forces)

Present in all molecules; dominant in non-polar molecules
Arise from induced and instantaneous dipoles
Strength increases with molecular size

Interpreting phase changes in a phase diagram

Phase diagram: shows stable phase (solid, liquid, gas) at given T and P
Key points:
- Triple point: all three phases coexist
- Critical point: liquid and gas phases indistinguishable
- Water: solid-liquid boundary slopes left (liquid denser than ice)

Freezing, melting, boiling, and sublimation

Freezing/melting: occur at solid-liquid boundary
Boiling/condensation: occur at liquid-gas boundary
Sublimation/deposition: direct solid-gas transitions

Colligative properties

Depend on number, not type, of solute particles
- Boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure
van’t Hoff factor ( $i$ ): number of particles solute forms in solution

Vapor pressure lowering (Raoult’s law)

$P = χ_{solvent} \cdot P_{solvent}^{\circ}$
$Δ P = χ_{solute} \cdot P_{solvent}^{\circ}$

Boiling point elevation

$Δ T_{b} = k_{b} \cdot m \cdot i$
- $k_{b}$ : boiling point constant, $m$ : molality, $i$ : van’t Hoff factor

Freezing point depression

$Δ T_{f} = - k_{f} \cdot m \cdot i$
- $k_{f}$ : freezing point constant

Osmotic pressure

$π = MRT \cdot i$
- $M$ : molarity, $R$ : gas constant, $T$ : temperature (K), $i$ : van’t Hoff factor

Differentiating types of mixtures

Solution: homogeneous at molecular level
Colloid: small aggregates, stable unless centrifuged (e.g., milk, emulsions)
Suspension: heterogeneous, large particles settle over time

Liquid phase

Hydrogen bonding

Hydrogen bond strength depends on bond polarity, with:

$H - F$ bonds being the strongest
- followed by $H - O$ bonds
- then $H - N$ bonds.

Dipole interactions

the ion has a larger magnitude of charge, and
the dipole is more polar.

Van der Waals’ forces (London dispersion forces)

These forces come from temporary dipoles:

Induced dipoles form when a polar molecule causes a nearby non-polar molecule to temporarily develop a dipole.
Instantaneous dipoles form when electron density in a non-polar molecule fluctuates momentarily, creating a brief dipole.

Dispersion forces get stronger as molecular size increases. For example, decane ( $C_{10} H_{22}$ ) has stronger dispersion forces than ethane ( $C_{2} H_{6}$ ).

Changes in phase

Interpreting phase changes in a phase diagram

A phase diagram shows which phase - solid, liquid, or gas - is stable at different temperatures and pressures.

In the solid state, molecules vibrate around fixed positions. Solids are nearly incompressible and don’t flow.
In the liquid state, molecules are still close together because of intermolecular forces, but they can move past one another. Liquids flow, take the shape of their container, and remain largely incompressible.
In the gas state, molecules move randomly and are far apart. Gases are easily compressible and expand to fill their container.

The diagram also includes phase boundaries, where two phases coexist in equilibrium:

The solid-liquid boundary represents equilibrium between solids and liquids.
The solid-gas boundary represents equilibrium between solids and gases.
The liquid-gas boundary represents equilibrium between liquids and gases.

The triple point is the specific combination of temperature and pressure at which all three phases exist simultaneously The critical point marks the conditions where the liquid and gas phases become indistinguishable The critical temperature is the highest temperature at which a liquid can exist.

Common mnemonic to remember the gas region: “gas comes out this way.”

Freezing, melting, boiling, and sublimation

Freezing point is the temperature at which a liquid starts to turn into a solid.
Melting point is the temperature at which a solid begins to become a liquid. These occur along the solid-liquid boundary.
The boiling point is the temperature at which a liquid transitions into a gas
The condensation point is the temperature at which a gas starts to form a liquid. Both occur along the liquid-gas boundary.
Sublimation describes the direct change from a solid to a gas, occurring under conditions found along the solid-gas boundary.
Deposition describes the direct change from a gas to a solid

it becomes harder for the liquid to boil (so the boiling point increases), and
it becomes harder for the liquid to freeze (so the freezing point decreases).

Another way to say this is that lowering a solution’s vapor pressure corresponds to raising its boiling point.

$P = χ_{solvent} \cdot P_{solvent}^{\circ} .$

The decrease in vapor pressure $(Δ P)$ can be written as:

$Δ P = χ_{solute} \cdot P_{solvent}^{\circ},$

where $χ_{solute}$ is the mole fraction of the solute.

Boiling point elevation is quantified by:

$Δ T_{b} = k_{b} \cdot m \cdot i$

Similarly, freezing point depression is described by:

$Δ T_{f} = - k_{f} \cdot m \cdot i$

$π = MRT \cdot i$

Differentiating types of mixtures:

A solution is a homogeneous mixture in which substances are combined at the molecular level and remain uniformly dispersed.
Colloids consist of solute aggregates that are extremely small and remain mixed unless separated by processes like centrifugation.
A suspension is a heterogeneous mixture where the particles are large enough that they do not stay evenly distributed over time.
A common example of a colloid is milk, and vigorous shaking of water and oil can produce an emulsion, which is also classified as a colloid.

Hydrogen bonding

Strong dipole-dipole interaction; requires H bonded to F, O, or N
Increases boiling point and viscosity
Strength order: H-F > H-O > H-N

Dipole interactions

Dipole-dipole: between polar molecules; weaker than hydrogen bonds
Ion-dipole: between ions and polar molecules; stronger than dipole-dipole
- Strength increases with higher ion charge and greater dipole polarity

Van der Waals’ forces (London dispersion forces)

Present in all molecules; dominant in non-polar molecules
Arise from induced and instantaneous dipoles
Strength increases with molecular size

Interpreting phase changes in a phase diagram

Phase diagram: shows stable phase (solid, liquid, gas) at given T and P
Key points:
- Triple point: all three phases coexist
- Critical point: liquid and gas phases indistinguishable
- Water: solid-liquid boundary slopes left (liquid denser than ice)

Freezing, melting, boiling, and sublimation

Freezing/melting: occur at solid-liquid boundary
Boiling/condensation: occur at liquid-gas boundary
Sublimation/deposition: direct solid-gas transitions

Colligative properties

Depend on number, not type, of solute particles
- Boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure
van’t Hoff factor ( $i$ ): number of particles solute forms in solution

Vapor pressure lowering (Raoult’s law)

$P = χ_{solvent} \cdot P_{solvent}^{\circ}$
$Δ P = χ_{solute} \cdot P_{solvent}^{\circ}$

Boiling point elevation

$Δ T_{b} = k_{b} \cdot m \cdot i$
- $k_{b}$ : boiling point constant, $m$ : molality, $i$ : van’t Hoff factor

Freezing point depression

$Δ T_{f} = - k_{f} \cdot m \cdot i$
- $k_{f}$ : freezing point constant

Osmotic pressure

$π = MRT \cdot i$
- $M$ : molarity, $R$ : gas constant, $T$ : temperature (K), $i$ : van’t Hoff factor

Differentiating types of mixtures

Solution: homogeneous at molecular level
Colloid: small aggregates, stable unless centrifuged (e.g., milk, emulsions)
Suspension: heterogeneous, large particles settle over time

Liquid phase

Hydrogen bonding

Hydrogen bond strength depends on bond polarity, with:

$H - F$ bonds being the strongest
- followed by $H - O$ bonds
- then $H - N$ bonds.

Dipole interactions

the ion has a larger magnitude of charge, and
the dipole is more polar.

Van der Waals’ forces (London dispersion forces)

These forces come from temporary dipoles:

Induced dipoles form when a polar molecule causes a nearby non-polar molecule to temporarily develop a dipole.
Instantaneous dipoles form when electron density in a non-polar molecule fluctuates momentarily, creating a brief dipole.

Dispersion forces get stronger as molecular size increases. For example, decane ( $C_{10} H_{22}$ ) has stronger dispersion forces than ethane ( $C_{2} H_{6}$ ).

Changes in phase

Interpreting phase changes in a phase diagram

A phase diagram shows which phase - solid, liquid, or gas - is stable at different temperatures and pressures.

In the solid state, molecules vibrate around fixed positions. Solids are nearly incompressible and don’t flow.
In the liquid state, molecules are still close together because of intermolecular forces, but they can move past one another. Liquids flow, take the shape of their container, and remain largely incompressible.
In the gas state, molecules move randomly and are far apart. Gases are easily compressible and expand to fill their container.

The diagram also includes phase boundaries, where two phases coexist in equilibrium:

The solid-liquid boundary represents equilibrium between solids and liquids.
The solid-gas boundary represents equilibrium between solids and gases.
The liquid-gas boundary represents equilibrium between liquids and gases.

The triple point is the specific combination of temperature and pressure at which all three phases exist simultaneously The critical point marks the conditions where the liquid and gas phases become indistinguishable The critical temperature is the highest temperature at which a liquid can exist.

Common mnemonic to remember the gas region: “gas comes out this way.”

Freezing, melting, boiling, and sublimation

Freezing point is the temperature at which a liquid starts to turn into a solid.
Melting point is the temperature at which a solid begins to become a liquid. These occur along the solid-liquid boundary.
The boiling point is the temperature at which a liquid transitions into a gas
The condensation point is the temperature at which a gas starts to form a liquid. Both occur along the liquid-gas boundary.
Sublimation describes the direct change from a solid to a gas, occurring under conditions found along the solid-gas boundary.
Deposition describes the direct change from a gas to a solid

it becomes harder for the liquid to boil (so the boiling point increases), and
it becomes harder for the liquid to freeze (so the freezing point decreases).

Another way to say this is that lowering a solution’s vapor pressure corresponds to raising its boiling point.

$P = χ_{solvent} \cdot P_{solvent}^{\circ} .$

The decrease in vapor pressure $(Δ P)$ can be written as:

$Δ P = χ_{solute} \cdot P_{solvent}^{\circ},$

where $χ_{solute}$ is the mole fraction of the solute.

Boiling point elevation is quantified by:

$Δ T_{b} = k_{b} \cdot m \cdot i$

Similarly, freezing point depression is described by:

$Δ T_{f} = - k_{f} \cdot m \cdot i$

$π = MRT \cdot i$

Differentiating types of mixtures:

A solution is a homogeneous mixture in which substances are combined at the molecular level and remain uniformly dispersed.
Colloids consist of solute aggregates that are extremely small and remain mixed unless separated by processes like centrifugation.
A suspension is a heterogeneous mixture where the particles are large enough that they do not stay evenly distributed over time.
A common example of a colloid is milk, and vigorous shaking of water and oil can produce an emulsion, which is also classified as a colloid.

Hydrogen bonding

Strong dipole-dipole interaction; requires H bonded to F, O, or N
Increases boiling point and viscosity
Strength order: H-F > H-O > H-N

Dipole interactions

Dipole-dipole: between polar molecules; weaker than hydrogen bonds
Ion-dipole: between ions and polar molecules; stronger than dipole-dipole
- Strength increases with higher ion charge and greater dipole polarity

Van der Waals’ forces (London dispersion forces)

Present in all molecules; dominant in non-polar molecules
Arise from induced and instantaneous dipoles
Strength increases with molecular size

Interpreting phase changes in a phase diagram

Phase diagram: shows stable phase (solid, liquid, gas) at given T and P
Key points:
- Triple point: all three phases coexist
- Critical point: liquid and gas phases indistinguishable
- Water: solid-liquid boundary slopes left (liquid denser than ice)

Freezing, melting, boiling, and sublimation

Freezing/melting: occur at solid-liquid boundary
Boiling/condensation: occur at liquid-gas boundary
Sublimation/deposition: direct solid-gas transitions

Colligative properties

Depend on number, not type, of solute particles
- Boiling point elevation, freezing point depression, vapor pressure lowering, osmotic pressure
van’t Hoff factor ( $i$ ): number of particles solute forms in solution

Vapor pressure lowering (Raoult’s law)

$P = χ_{solvent} \cdot P_{solvent}^{\circ}$
$Δ P = χ_{solute} \cdot P_{solvent}^{\circ}$

Boiling point elevation

$Δ T_{b} = k_{b} \cdot m \cdot i$
- $k_{b}$ : boiling point constant, $m$ : molality, $i$ : van’t Hoff factor

Freezing point depression

$Δ T_{f} = - k_{f} \cdot m \cdot i$
- $k_{f}$ : freezing point constant

Osmotic pressure

$π = MRT \cdot i$
- $M$ : molarity, $R$ : gas constant, $T$ : temperature (K), $i$ : van’t Hoff factor

Differentiating types of mixtures

Solution: homogeneous at molecular level
Colloid: small aggregates, stable unless centrifuged (e.g., milk, emulsions)
Suspension: heterogeneous, large particles settle over time