exam 1 chem - patterns

Intermolecular Forces (IMF)

Attractions between molecules that determine properties such as boiling point and vapor pressure.

Strong IMF

Leads to high boiling points, low vapor pressure, high viscosity, and high surface tension.

Weak IMF

Results in low boiling points, high vapor pressure, low viscosity, and low surface tension.

Hydrogen Bonding

Special type of dipole-dipole interaction involving H bonded to N, O, or F.

Dispersion Forces

Present in all molecules; strength increases with the size of electron cloud.

Polarity

Determined by bond polarity and molecular geometry; affects intermolecular forces.

Vapor Pressure

Pressure exerted by a vapor in equilibrium with its liquid or solid phase.

Boiling Point

Temperature at which vapor pressure equals atmospheric pressure.

Colligative Properties

Properties dependent on the number of dissolved particles, not their identity.

Saturated Solution

Contains maximum amount of dissolved solute at equilibrium.

Supersaturated Solution

Contains more dissolved solute than is possible at equilibrium.

Osmosis

Movement of solvent from low solute concentration to high solute concentration.

Van't Hoff Factor (i)

Number of particles the solute produces in solution; affects colligative properties.

Critical Point

Temperature and pressure at which gas and liquid phases become indistinguishable.

Endothermic Process

Process that absorbs energy, such as melting and vaporization.

Exothermic Process

Process that releases energy, such as condensation and freezing.

Phase Change

Transition between solid, liquid, and gas phases, involving energy absorption or release.

Ideal Gas Law

Equation of state for a gas: PV=nRT; relates pressure, volume, and temperature.

Raoult's Law

States that the vapor pressure of a solvent in a solution is directly proportional to the mole fraction of the solvent.

Molarity (M)

Concentration expressed as moles of solute per liter of solution.

Molality (m)

Concentration expressed as moles of solute per kilogram of solvent.

Crystalline Solids

Solids with organized, repeating structures.

Amorphous Solids

Solids that lack a regular three-dimensional arrangement.

Atmospheric Pressure

Weight of the atmosphere above a given point; affects boiling point.

Viscosity

Resistance of a fluid to flow; increases with stronger intermolecular attractions.

Surface Tension

Energy required to increase the surface area of a liquid; increases with stronger IMF.

Volatility

Tendency of a substance to vaporize; higher volatility indicates weaker IMF.

Gas Solubility and Temperature

Gas solubility decreases as temperature increases.

Gas Solubility and Pressure

Gas solubility increases as pressure increases.

Intermolecular Force Trends (Strong IMF)

Stronger intermolecular forces lead to higher boiling point, higher viscosity, higher surface tension, lower vapor pressure, and lower volatility.

Intermolecular Force Trends (Weak IMF)

Weaker intermolecular forces lead to lower boiling point, lower viscosity, lower surface tension, higher vapor pressure, and higher volatility.

London Dispersion Forces

Intermolecular forces present in all atoms and molecules. Their strength increases with greater polarizability (more electrons) and larger molecular contact area.

Dipole-Dipole Forces

Intermolecular forces present between polar molecules due to permanent dipole moments.

Hydrogen Bonding

A strong type of dipole-dipole attraction occurring when $H$ is directly bonded to $N$, $O$, or $F$ (e.g., $H_2O$, $NH_3$, $HF$).

Polarizability

The ease with which the electron cloud of an atom or molecule can be distorted; increases with a greater number of electrons.

Molecular Polarity Requirements

Determined by both bond polarity and molecular geometry. Symmetrical molecules with identical surrounding atoms (such as $CO_2$ and $BF_3$) are nonpolar because dipoles cancel.

Boiling Point vs. Vapor Pressure Relationship

They move in opposite directions. Stronger IMFs result in higher boiling points but lower vapor pressures.

Temperature vs. Vapor Pressure Relationship

They move in the same direction. As temperature increases, vapor pressure increases.

Definition of Boiling Point

The temperature at which the vapor pressure of a liquid equals the external atmospheric pressure.

Altitude and Boiling Point Relationship

As altitude increases, atmospheric pressure decreases, causing the boiling point of a liquid to decrease.

Volatility

The measure of how readily a substance evaporates. High volatility corresponds to weak IMFs, high vapor pressure, and low boiling point.

Viscosity

The resistance of a liquid to flow. Stronger IMFs lead to higher viscosity.

Surface Tension

The energy required to increase the surface area of a liquid. Stronger IMFs result in higher surface tension.

Endothermic Phase Changes

Phase changes requiring energy absorption: Melting (Fusion: solid $\rightarrow$ liquid), Vaporization (liquid $\rightarrow$ gas), and Sublimation (solid $\rightarrow$ gas).

Exothermic Phase Changes

Phase changes releasing energy: Freezing (liquid $\rightarrow$ solid), Condensation (gas $\rightarrow$ liquid), and Deposition (gas $\rightarrow$ solid).

Phase Changes and Entropy ($\Delta S$)

Moving toward the gas phase increases randomness ($\Delta S > 0$); moving toward the solid phase decreases randomness ($\Delta S < 0$).

Relationship Between Enthalpies of Phase Changes

According to Hess's Law, \Delta H_{sub} = \text{\Delta H}_{fus} + \text{\Delta H}_{vap}.

Crystalline vs. Amorphous Solids

Crystalline solids have a rigid, repeating long-range order. Amorphous solids lack a regular 3D arrangement (e.g., glass, rubber).

Four Types of Crystalline Solids

Ionic (held by electrostatic attractions), Molecular (held by IMFs), Covalent Network (held by covalent bonds throughout), and Metallic (held by delocalized electron sea).

Simple Cubic (SC) Unit Cell Properties

Contains $1$ atom per cell, coordination number = $6$, and packing efficiency = $52\%$.

Body-Centered Cubic (BCC) Unit Cell Properties

Contains $2$ atoms per cell, coordination number = $8$, and packing efficiency = $68\%$.

Face-Centered Cubic (FCC) Unit Cell Properties

Contains $4$ atoms per cell, coordination number = $12$, and packing efficiency = $74\%$ (cubic closest-packed).

Unit Cell Atom Contributions

Corner atom = $\frac{1}{8}$, Face atom = $\frac{1}{2}$, Body-center atom = $1$.

Solute vs. Solvent

Solvent is the component present in the greatest amount; solute is the component present in lesser amounts in a homogeneous solution.

Miscibility

The property of two liquids being completely soluble in each other in all proportions (e.g., nonpolar $CCl_4$ and nonpolar benzene).

Saturated vs. Supersaturated Solutions

Saturated solutions contain the maximum equilibrium amount of dissolved solute. Supersaturated solutions contain more than the equilibrium amount.

Gas Solubility and Temperature

Gas solubility decreases as temperature increases. Cold liquids hold more dissolved gas.

Gas Solubility and Pressure (Henry's Law)

Gas solubility increases as pressure increases.

Colligative Properties

Properties of solutions that depend solely on the number of dissolved solute particles, not their chemical identity (VP lowering, BP elevation, FP depression, osmotic pressure).

Master Colligative Property Pattern (Nonvolatile Solute)

Adding a nonvolatile solute causes vapor pressure to decrease, boiling point to increase, and freezing point to decrease.

van't Hoff Factor ($i$)

The ratio of moles of particles in solution to moles of solute dissolved. For nonelectrolytes $i = 1$; for $NaCl$, $i \approx 2$.

Ion Pairing Effect

Real electrolyte solutions can have an effective van't Hoff factor lower than ideal predictions due to associated ion pairs reducing independent particle counts.

Molarity vs. Molality Temperature Dependence

Molarity ($mol\,dm^{-3}$) depends on solution volume, making it temperature-dependent. Molality ($mol\,kg^{-1}$) depends on solvent mass, making it temperature-independent.

Osmosis Direction

Solvent moves across a semipermeable membrane from a less concentrated solution to a more concentrated solution (water moves toward more solute).

Osmotic Pressure Formula

$\Pi = MRT$, where $\Pi$ is osmotic pressure, $M$ is molarity, $R$ is the ideal gas constant, and $T$ is absolute temperature.

What happens to boiling point as intermolecular forces (IMF) become stronger?

Boiling point increases ($\text{BP} \uparrow$) because more heat energy is required to separate the molecules.

What happens to vapor pressure as intermolecular forces (IMF) become stronger?

Vapor pressure decreases ($\text{VP} \downarrow$) because fewer molecules can easily escape into the gas phase.

What happens to viscosity as intermolecular forces (IMF) become stronger?

Viscosity increases because stronger attractions make it harder for molecules to flow past one another.

What happens to surface tension as intermolecular forces (IMF) become stronger?

Surface tension increases because interior molecules pull surface molecules inward more strongly.

What happens to volatility as intermolecular forces (IMF) become stronger?

Volatility decreases because molecules do not evaporate easily.

Which intermolecular force is present in ALL atoms and molecules?

London dispersion forces.

What condition must be met for hydrogen bonding to occur in a molecule?

Hydrogen must be directly bonded to Nitrogen, Oxygen, or Fluorine ($\text{H--N}$, $\text{H--O}$, or $\text{H--F}$).

How does the number of electrons in a molecule affect London dispersion forces?

More electrons increase polarizability, leading to stronger London dispersion forces.

How does molecular shape affect dispersion forces when comparing isomers?

Less compact shapes with greater contact area between molecules produce stronger dispersion forces.

Does a molecule with polar bonds always have a dipole moment (is it always polar)?

No, if the molecular geometry is symmetrical, the bond dipoles may cancel out, making the molecule nonpolar.

Is $\text{CO}_2$ polar or nonpolar, and why?

Nonpolar, because its linear geometry causes the two equal polar $\text{C=O}$ bond dipoles to cancel out.

Is $\text{H}_2\text{O}$ polar or nonpolar, and why?

Polar, because its bent geometry prevents the two $\text{O--H}$ bond dipoles from canceling out.

What is the relationship between boiling point and vapor pressure for different liquids at the same temperature?

They move in opposite directions; a liquid with a lower vapor pressure has a higher boiling point.

What is the effect of increasing temperature on the vapor pressure of a liquid?

Vapor pressure increases ($\text{T} \uparrow \implies \text{VP} \uparrow$) because more molecules gain enough kinetic energy to evaporate.

What is the exact condition required for a liquid to boil?

The liquid's vapor pressure must equal the external atmospheric pressure ($\text{VP} = P_{\text{ext}}$).

How does lower atmospheric pressure (such as at high altitude) affect the boiling point of water?

Lower atmospheric pressure decreases the boiling point of water because less vapor pressure is required for boiling to occur.

Are phase changes from solid to liquid to gas endothermic or exothermic?

Endothermic, because heat energy must be absorbed to break or weaken molecular attractions.

Are phase changes from gas to liquid to solid endothermic or exothermic?

Exothermic, because energy is released as molecules move closer together and establish stronger attractions.

How does entropy change during the phase transition from liquid to gas?

Entropy increases ($\Delta S > 0$) because molecular disorder and randomness increase significantly in the gas phase.

What equation connects the enthalpy of sublimation, fusion, and vaporization using Hess's Law?

$$\Delta H_{\text{sub}} = \Delta H_{\text{fus}} + \Delta H_{\text{vap}}$$

What is the value of Gibbs free energy change ($\Delta G$) at phase equilibrium?

$\Delta G = 0$, which leads to the relationship $\Delta H = T \Delta S$.

What is the key difference between crystalline and amorphous solids?

Crystalline solids have a well-defined, repeating three-dimensional long-range order, while amorphous solids lack a regular repeating structure.

What are the four main types of crystalline solids?

Ionic, Molecular, Covalent Network, and Metallic solids.

What particles and forces hold a molecular crystalline solid together?

Molecules held together by intermolecular forces (dispersion, dipole-dipole, or hydrogen bonding).

What structural feature gives metallic solids high electrical conductivity and malleability?

Delocalized electrons moving freely through a "sea of electrons" around metal cations.

What is the coordination number and packing efficiency of a simple cubic (SC) unit cell?

Coordination number = $6$; packing efficiency = $52\%$.

What is the coordination number and packing efficiency of a body-centered cubic (BCC) unit cell?

Coordination number = $8$; packing efficiency = $68\%$.

What is the coordination number and packing efficiency of a face-centered cubic (FCC) unit cell?

Coordination number = $12$; packing efficiency = $74\%$.

How many net atoms are contained within a single Simple Cubic (SC) unit cell?

$1$ net atom ($8 \times \frac{1}{8}$ corner atoms).

How many net atoms are contained within a single Body-Centered Cubic (BCC) unit cell?

$2$ net atoms ($8 \times \frac{1}{8}$ corner atoms + $1$ body-center atom).

How many net atoms are contained within a single Face-Centered Cubic (FCC) unit cell?

$4$ net atoms ($8 \times \frac{1}{8}$ corner atoms + $6 \times \frac{1}{2}$ face atoms).

How much does an atom on a unit cell corner, face, and body-center contribute to that cell?

Corner = $\frac{1}{8}$, Face = $\frac{1}{2}$, Body-center = $1$.

Which unit cell structures represent closest-packed arrangements with $74\%$ packing efficiency?

Face-Centered Cubic (FCC / Cubic Closest-Packed) and Hexagonal Closest-Packed (HCP).

What defines a saturated solution?

A solution containing the maximum equilibrium amount of dissolved solute at a given temperature, where rate of dissolution equals rate of crystallization.

What is a supersaturated solution?

A solution containing more dissolved solute than its equilibrium saturation amount at a given temperature.