States of Matter and Intermolecular Forces

States of Matter Overview

  • Comparison of Solids, Liquids, and Gases

    • Solid
      • Volume/Shape: Condensed, virtually incompressible, maintains shape and volume.
      • Density: Very high, not affected by pressure.
      • Diffusion: Very low.
      • Flow/Motion of Molecules: Does not flow; molecules vibrate about fixed positions.
    • Liquid
      • Volume/Shape: Condensed, maintains volume, assumes shape of container.
      • Density: Very high, not affected by pressure.
      • Diffusion: Occurs slowly.
      • Flow/Motion of Molecules: Flows readily; molecules slide past one another freely.
    • Gas
      • Volume/Shape: Large spaces between particles; assumes volume and shape of container.
      • Density: Very low, depends on pressure.
      • Diffusion: Very rapid; evenly distributed in space, producing a homogeneous mixture quickly.
      • Flow/Motion of Molecules: Flows readily; very free, random, and dynamic motion.

  • Kinetic Molecular Theory

    • The physical state of a substance depends on the balance between the kinetic energies (KE) of the particles and the strength of attraction between them.
    • External forces affecting state of matter:
      • Temperature: As temperature (T) increases, the average KE of molecules increases. When the KE is large enough to overcome the attractive forces between molecules, a change in state occurs.
      • Pressure: As pressure (P) increases, the molecules of a substance move closer together. This causes an increase in attraction between particles and a decrease in KE, so a change in state occurs.
    • Solid forms of a substance are very condensed and arranged in tightly packed, 3-D configurations. As substances gain energy, they overcome the forces holding them in place and change states of matter. Liquids are held together with little space between molecules, but they are able to flow more freely past each other.
    • The state a substance is at room temperature is determined by the strength of the attraction between its particles: intermolecular forces.

Intermolecular Forces

  • Interatomic vs. Intermolecular Forces

    • Interatomic: between atoms, within molecules (intramolecular) = BONDS
      • Covalent (Molecular) Bonds
        • Shared pairs of electrons where the electron clouds overlap; very strong bond.
        • Low boiling point (bp), low melting point (mp), non-malleable, do not conduct electricity, significant vapor pressure (for molecules under 500 g/mol).
      • Ionic Bonds
        • Transfer of electrons creating a bond that is held together by the electrostatic attraction between cations and anions.
        • All ionic compounds are "salt crystals" and solid at room temperature.
        • High bp, mp, non-malleable, do not conduct electricity in solid form, but do conduct electricity in water, no vapor pressure.
      • Metallic Bonds
        • Atoms of the same metallic element readily bond by sharing their valence electrons (e^-).
        • Metallic ions are fixed and surrounded by a moving cloud of e^-.
        • This gives metals their physical property of conductivity, strength, and luster.
        • Very strong attraction between atoms.
        • Most metals are solid at room temperature.
        • Examples: Na, Al, Fe, etc.
        • Variable melting point (mp), boiling point (bp), do conduct electricity, most have no vapor pressure.
    • Intermolecular: Between Molecules
      • There are 4 types of intermolecular forces (IMFs): all intermolecular forces are due to attractions between positive and negative species (similar to ionic bonds).
        • Ion-dipole forces: between an ion and a polar molecule.
        • Van der Waals Forces: intermolecular attractions between neutral molecules or monatomic gases.
          • Dipole-dipole forces
          • London dispersion forces (LDF)
          • Hydrogen bonds (H-bonds)

  • Types of Intermolecular Forces

    • Ion-Dipole
      • Between an ion and the partial charge on the end of a polar molecule (\delta^+ / \delta^-).
      • The magnitude of the attraction increases as either the charge of the ion or the magnitude of the dipole moment increases. A dipole moment is a difference in charge, or a measure of polarity. A larger dipole moment indicates greater polarity.
      • Example: solution of NaCl in water. \text{NaCl}(s) \rightarrow \text{Na}^+ (aq) + \text{Cl}^- (aq)
    • Dipole-Dipole Forces
      • Exist between polar molecules.
      • Occur when the \delta^+ end of one molecule attracts the \delta^- end of an adjacent molecule (polar molecules attract each other).
      • Generally weaker than ion-dipole forces.
      • Overall effect is net attraction (even with attraction and repulsion).
      • In liquids, for molecules of approximately equal mass and size, the strengths of intermolecular attractions increase with increasing polarity (which can be determined by examining the difference in electronegativity).
      • The boiling points increase with increasing magnitude of dipole moment.
      • Example: Which of the following substances would have dipole-dipole forces? Ar, HCl, HI, CH3Cl, CH4. Would HCl or HI have the strongest dipole-dipole force? Explain.
    • [London] Dispersion Forces (LDF)
      • All atoms, compounds, and molecules have London dispersion forces.
      • This temporary force of attraction is caused by the random, constant motion of electrons in an atom or molecule.
      • Creates an instantaneous dipole moment that is only significant when molecules are very close together.
      • Polarizability
        • A measure of the "squashiness" of the electron (e^-) cloud of an atom or molecule.
        • The ease with which the charge distribution in an atom or molecule can be distorted by an external electric field.
        • Dispersion forces tend to increase in strength with increasing molecular weight (molar mass) because they have more electrons, so the electron cloud is larger, providing more area for distortion and a larger instantaneous dipole moment.
      • Shapes that have access for more contact can have stronger IMFs.
      • Which of the following atoms would you expect to have the strongest London dispersion forces: nitrogen, phosphorus, arsenic? Explain.
    • Hydrogen Bonding (H-Bond)
      • A special type of intermolecular attraction that exists between a hydrogen atom in a polar bond (H-F, H-O, H-N) and an unshared electron pair on a nearby electronegative ion or atom (usually an F, O, or N atom in an adjacent molecule).
      • This occurs because hydrogen has no inner core of electrons.
        • The bare proton of the H nucleus creates the positive side of the bond dipole in the molecule.
        • The hydrogen atom is so small (just a p^+ and n^0), it is attracted to and can approach a highly electronegative atom in a nearby molecule very closely and interact strongly with it.
      • H-bonds are the strongest of the Van der Waals forces.
      • Which of the following substances can have hydrogen bonding? CH2F2, H2O, SiH4, CH3OH, CH2O, NH2.
      • They play an important role in biochemical systems.
        • The folding of proteins, carbohydrates, and nucleic acids involve the formation of H-bonds.
        • This folding is important for creating the 3-D shapes necessary for these molecules to fit together.
        • H-bonds create the “ladder” of the DNA molecule, holding the antiparallel strands together in the double helix structure.
        • Proteins exert their effects by bonding to “binding sites” on molecules in the lock-and-key method.
        • The presence of HFON on biomolecules allows them to be dissolved in the water present in organisms (blood, cytoplasm, etc.).
      • H2O: H-bonds are what cause the density of ice to be less than that of liquid water.
        • It is unusual for the liquid phase of a substance to be denser than its solid phase.
        • When water freezes, the molecules assume an ordered and open arrangement, which leads to a decreased density.
        • Accounts for the basic hexagonal shape of snowflakes.
        • Lake turnover:
          • The seasonal movement of water in a lake has a profound effect on aquatic life in deep lakes (like Lake Champlain).
          • In the fall, as the air cools, the lake water is cooled.
          • When the surface water reaches a temperature of 4 °C, it sinks to the bottom of the lake (because water is its densest at 4 °C).
          • This occurs continually until there is a column of 4 °C water in the lake.
          • At that point, the surface water can drop below 4 °C and freeze.
          • The opposite process occurs in the spring.
          • Why is it important to lake ecosystems?
            • Mixing causes aeration of the entire water column.
            • Nutrients that have settled to the bottom of the lake are distributed throughout the water.

  • Strength of Intermolecular Forces

    • Overview of types and strength of IMF.
    • Intermolecular Forces are used to determine and explain physical properties of a substance.
      • We need to know which IMFs are present to compare strengths of attraction between molecules.
      • Stronger forces of attraction typically indicate that substances are more likely to be solids at room temperature, and weaker forces of IMF typically indicate that substances are more likely to be gases at room temperature.
      • Bonds within molecules are much stronger than attractions between molecules: Metallic bond / Ionic bond (Electrostatic Ion-Ion forces) / Covalent bond >>>> H-bonds> Dipole> LDF
    • IMF effects on Physical Properties
      • (Stronger IMF) Greater energy is needed to break the IMFs (attractions between molecules), thus the substance has higher boiling point, higher melting point, greater surface tension, greater viscosity.
      • (Weaker IMF) Less energy is needed to break the IMFs (attractions between molecules), thus the substance has higher vapor pressure, greater volatility.
    • How do we use this information?
      • Identify the IMFs in each of the following substances.
      • Ranks the substances in order of increasing boiling point: Cl2, NaCl, I2, CH2Cl2, HF

Liquids

  • Viscosity
    • The resistance of a liquid to flow.
    • The stronger the IMFs of a substance, the more slowly it flows, so the higher the viscosity.
      • Examples of slow-flowing liquids:
      • Examples of fast-flowing liquids:
    • Measurement
      • The time it takes for a certain amount of liquid to flow through a small tube.
      • The rate at which a steel sphere falls through the liquid.
    • The greater the IMFs, the greater the viscosity.
      • ↑ IMFs means stronger attractions between molecules
      • It is more difficult to move molecules over each other for flow to occur
    • As T ↑, viscosity ↓
      • As KE ↑, IMFs are more easily overcome.
      • Flow ↑ as IMFs ↓ so viscosity ↓
    • As P ↑, viscosity ↑
      • As molecules move closer together, IMFs ↑
      • With stronger IMFs, flow ↓ and viscosity ↑
  • Surface Tension
    • Definition
      • The energy required to increase the surface area of a liquid by a given amount.
      • A measure of the inward force that must be overcome to expand the surface of a liquid.
    • Water has a high surface tension due to its strong IMFs.
      • Interior molecules are attracted to other molecules in 3D.
      • Surface molecules: there are no molecules above them.
        • Attractions are horizontally and inward.
        • Creates a “skin” at the surface.
        • Accounts for the “beading” of water on a surface, creating the shape of a sphere.
    • Cohesive forces
      • IMFs that attract similar molecules to one another.
      • Water has high cohesive forces.
    • Adhesive forces
      • IMFs that attract a substance to a surface.
      • Meniscus of water vs. mercury
        • Water: molecules adhere to the glass molecules; cohesive forces between water molecules bring other H2O molecules with it as they climb up the glass; adhesive forces > cohesive forces.
        • Mercury: metallic bonds between molecules = strong; cohesive forces > adhesive forces.
      • Capillary action
        • Important in moving water/dissolved nutrients in plants.
        • Liquid will “climb” up a tube until the adhesive and cohesive forces are balanced by gravity.
        • The stronger the IMFs within a liquid, the higher that liquid will rise in a tube.
        • For the same liquid, the more narrow the tube, the higher the substance rises.
  • Vapor Pressure (VP)
    • The pressure exerted by a vapor when the liquid and vapor states are in dynamic equilibrium.
    • Dynamic equilibrium
      • When 2 opposing processes are occurring simultaneously at equal rates.
      • Occurs in a closed system.
      • There is no net change in the system.
        • As the liquid evaporates, the pressure in the space above the liquid ↑
        • When the rate at which liquid molecules enter the gas phase equals the rate at which gas molecules enter the liquid phase, a stable VP results
        • If there are many particles in the gas phase, a high VP results
        • If there are few particles in the gas phase, a low VP results
      • The flask on the right shows the liquid and gas phases in dynamic equilibrium. For every liquid molecule that becomes a gas, a molecule of gas enters the liquid phase. Equilibrium: the number of molecules in the gas phase remains constant. Dynamic: which molecules are in the gas phase is constantly changing. The vapor pressure at dynamic equilibrium is stable.
    • Evaporation - vaporization with no change in T or P
    • Boiling - vaporization with a change in T or P
    • All liquids produce a measurable vapor pressure
      • High VP = weak IMFs = low BP
        • liquid molecules move easily into the gas phase
        • lots of gas molecules in the space above the liquid
      • Low VP = strong IMFs = high BP
        • liquid molecules are more difficult to move into the gas phase
        • not many gas molecules in the space above the liquid
    • What happens to the vapor pressure as the temperature increases? As the temperature increases, the vapor pressure increases. This is because as the temperature increases, more particles have enough kinetic energy and can escape the liquid phase of matter. The more particles of vapor above the liquid there are, the more collisions occur, which causes an increase in pressure.
    • Volatility (volatile = “able to fly”)
      • the tendency of a liquid to evaporate quickly
      • more volatile liquids: have ↑ VP, have ↓ IMFs, have ↓ BP
  • Boiling Point (BP)
    • A liquid has reached its boiling point when its vapor pressure equals the external pressure acting on the surface of the liquid. (Bubbles of vapor form within the interior of the liquid.)
    • Normal BP = BP of the liquid at 1 atmosphere of pressure (H2O= 100 °C)
    • High elevations: ↓ P (the column of air above is decreased), VP is below normal; the BP of water ↓, so T of water is ↓, and the food takes longer to cook
    • Pressure cooker: A sealed container with water, heated so VP is above normal; The BP of water ↑, so T inside is ↑, and the food inside takes less time to cook

Solids

  • Solids are less chaotic (↓KE) and have organized structures in a pattern. They are dense and incompressible.

  • Classification of Solids

    • Metallic- held together by collectively shared valence electrons
    • Ionic- held together by electrostatic attraction between cations and anions
    • Molecular- held together by intermolecular forces
    • Covalent-network- held together by an extensive network of covalent bonds (diamonds, quartz)

  • Types of solids

    • Crystalline solids
      • Most solids are crystalline
      • Made of repeating unit cells in a 3D regular, repeating pattern
      • Repeating pattern is called a crystal lattice
      • There are 7 crystal systems, named for their 3D pattern
    • Amorphous solids
      • “without form”
      • Particles in this solid have no orderly structure and so lack a well-defined shape
      • These solids tend to soften before melting
      • examples: rubber, glass, wax
    • Allotropes
      • Different forms of the same element in the same physical state
      • examples: diamond and graphite

  • Melting Point (MP)

    • The temperature at which a solid becomes a liquid.
    • Process: As T ↑, the KE of the particles ↑ and the particles vibrate more quickly. This ↑ vibration disrupts the orderly pattern of the solid, and they can flow past each other.
    • Point at which the disruptive vibrations of the particles > IMFs holding the particles in place
    • Application:
      • A substance with strong IMFs will have a high MP.
      • Ionic and metallic solids have very high MPs.

Phase Changes

  • Phase Changes

    • The effect of heat on the phase of matter
      • Vaporization
      • Sublimation
      • Melting
    • In order to change phases, you have to change the energy of the system. In order to transition from a more ordered state of matter to a less ordered state of matter, energy must be added to overcome intermolecular forces.

  • Liquid/ Gas Equilibrium and Transitions

    • Evaporation/ Vaporization- requires energy, liquid transformed to gas
      • Heat of Vaporization= heat needed for the vaporization of a liquid: \Delta H_{vap} in kJ/mol. The heat of vaporization for water is 40.7 kJ/mol.
    • Condensation- gas transformed to liquid
      • How much energy to go in the reverse direction? What is the amount of energy associated with the condensation of water?
    • Boiling Point (Condensation Point)
      • Boiling is vaporization with a change in Pressure or Temperature
      • Boiling Point is the temperature where the vapor pressure equals the external pressure (atmospheric pressure) acting on the surface of the liquid.
      • The boiling point occurs at the same temperature as the condensation point.
      • The normal boiling point is the boiling point at 1 atm pressure (water= 100 °C)

  • Solid/ Liquid Equilibrium and Transitions

    • Melting (Fusion)= requires energy, solid transformed into a liquid
      • Heat of Fusion= heat needed for the melting of a solid: \Delta H_{fus} = kJ/mol. The heat of fusion for water is 6.01 kJ/mol.
      • How much heat is needed to melt 95.0 grams of H2O if the heat of fusion is 6.01kJ/mol?
    • Freezing= the transformation of liquid to solid
    • Melting Point (Freezing point)
      • Melting point is the temperature where a solids turns into a liquid.
      • The melting point occurs at the same temperature as the freezing point.
      • The normal melting point is melting point at 1 atm pressure.

  • Solid/ Gas Equilibrium and Transitions

    • Sublimation= the process during which molecules go directly from the solid into the vapor phase
      • Heat of Sublimation= heat needed for the sublimation of a solid: \Delta H_{sub} = kJ/mol
    • Deposition= the reverse process, where molecules make the transition from vapor to solid directly

  • Heating Curves

    • Heating Curve
      • Shows the transitions of a substance (s→l→g) over time with ↑T
      • During a phase change:
        • the graph is flat
        • all added energy goes to overcoming IMFs, not ↑ T
        • once the substance has completed the phase change, any added energy works to ↑ T
      • While a pure substance is changing phases of matter under constant pressure, it remains the same temperature until all of the matter has changed phases.
    • Heating Curve Example:
    • Example: How much energy is required when 55.0 grams of ice at -15.0 °C is converted to water at 58.0°C?

  • Phase Diagrams

    • A phase diagram is a graphical way to summarize the conditions under which equilibria exist between different states of matter. It also allows for the prediction of a phase of a substance that is stable at any given temperature and pressure.
    • General form:
      • Exhibits 3 phases with three curves at various temperatures and pressures
      • A-B is the vapor –pressure curve of the liquid. It represents the equilibrium between the liquid and gas phases. The point on this curve where the vapor pressure is 1 atm is the normal boiling point. The vapor-pressure curve ends at the critical point (B), which is the point where the gas or liquid phase becomes indistinguishable and forms a supercritical fluid (which means “a fluid above the critical point”).
      • The line AC represents the variation in the vapor pressure (the pressure of gas above the solid) of the solid as it sublimes at different temperatures. The sublimation point of a solid is identical to its deposition point.
      • The line from A- D represents the change in melting point of the solid with increasing pressure. For most substances, solid is denser than the liquid, therefore, an increase in pressure usually favors a more compact solid phase. Thus, higher temperatures are required to melt solid at higher pressures. The melting point of a solid is identical to its freezing point.
      • Point A is where the curves intersect is known as the triple point- where all three phases are at dynamic equilibrium at this temp and pressure.
    • Examples:
      • H2O
      • CO2