Chapter 12- Liquids, Solids, and Intermolecular Forces

Chapter 12- Liquids, Solids, and Intermolecular Forces

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Previous Class Chapter 11

11.2 Kinetic Molecular Theory: A Model for Gases

11.3 Pressure: The Result of Constant Molecular Collisions

11.4 Boyle’s Law: Pressure and Volume

11.5 Charles’s Law: Volume and Temperature

11.6 The Combined Gas Law: Pressure, Volume, and Temperature

11.7 Avogadro’s Law: Volume and Moles

11.8 The Ideal Gas Law: Pressure, Volume, Temperature, and Moles

11.9 Mixtures of Gases

11.10 Gases in Chemical Reactions

Multiple Choice

What about gases can be measured?

• Pressure and Volume
• Temperature, Volume, and Pressure
• Volume and Temperature
• Pressure, Temperature, Volume, and Moles

Multiple Choice

In order to convert to Kelvin, you add ______ to the Celsius measurement.

• 372
• 273
• 237
• 732

Multiple Choice

You have a gas that has a pressure of 2 ATM and a volume of 10L.  What would be the new volume if the pressure was changed to 1 ATM?

• 5 L
• 20 L
• It would stay at 10L
• 1 L

Multiple Choice

Boyle's law :  The pressure and volume of a gas show a _____ relationship. 

• No relationship 
• Direct 
• Inverse 
• Equal

Multiple Choice

Which of the following would increase the (gas) pressure of a system?

• Increase the Temperature
• Pump in more gas
• Decrease the volume
• All of these

Multiple Choice

If a balloon is cooled what will happen to the volume?

• Volume will increase
• Volume will decrease
• Volume will not change
• Balloon freezes to death

Multiple Choice

Calculate the volume that a 0.323-mol sample of a gas will occupy at 265 K and a pressure of 0.900 atm.

• 7.18 L 
• 7.81 L
• 4.63 L
• 4.36 L

Multiple Choice

Charles' Law States...

• As Pressure goes up volume goes down
• As Pressure goes up temperature goes up
• As Volume goes up temperature goes up 
• As Pressure goes down volume goes down 

Multiple Choice

What is the pressure of a car tire that had an initial pressure of 1.8 atm but was heated from 38°C to 123°C?

• 0.9 atm

• 2.1 atm

• 1.6 atm

• 3.4 atm

Multiple Choice

A student measures the pressure and volume of an empty water bottle to be 1.4 atm and 2.3 L. She then decreases the pressure to 0.65 atm. What is the new volume?

• 2.1 L

• 5.0 L

• 8.2 L

• 3.9 L

Liquids, Solids, and Intermolecular Forces

12.2 Properties of Liquids and Solids

12.3 Intermolecular Forces in Action: Surface Tension and Viscosity

12.4 Evaporation and Condensation

12.5 Melting, Freezing, and Sublimation

12.6 Types of Intermolecular Forces: Dispersion, Dipole–Dipole, Hydrogen Bonding, and Ion–Dipole

12.7 Types of Crystalline Solids: Molecular, Ionic, and Atomic

Properties of Liquids

Definite Volume but Indefinite Shape:

• Liquids have a fixed volume but take the shape of their container. This is because the particles in a liquid are close together but still have enough freedom to move past each other.

Fluidity:

• Liquids can flow because their molecules slide past each other. This property allows them to be poured and take the shape of their container.

Viscosity:

• Viscosity is a liquid's resistance to flow. It depends on the strength of intermolecular forces; stronger forces (like hydrogen bonding) lead to higher viscosity. Temperature also affects viscosity—raising the temperature usually decreases it.

Surface Tension:

• Surface tension is the "skin-like" effect on the surface of a liquid due to cohesive forces pulling molecules at the surface inward. Water has high surface tension due to hydrogen bonding, allowing certain objects, like light insects, to "walk" on water.

Capillary Action:

• This is the ability of a liquid to flow in narrow spaces without external forces (like gravity). It occurs due to the adhesive forces between the liquid and the container and is visible in phenomena like water rising in plant stems or thin tubes.

Evaporation and Vapor Pressure:

• Evaporation occurs when molecules at the liquid’s surface gain enough energy to overcome intermolecular forces and escape as a gas. Vapor pressure is the pressure exerted by the vapor in equilibrium with its liquid. Substances with weaker intermolecular forces have higher vapor pressures.

Boiling Point:

• The boiling point is the temperature at which the vapor pressure of a liquid equals atmospheric pressure, allowing it to transition to a gas. Stronger intermolecular forces raise the boiling point.

Properties of Solids

Definite Shape and Volume:

• Solids maintain both a fixed shape and volume due to strong intermolecular forces that hold the particles in a rigid structure.

High Density:

• Solids have tightly packed particles, which makes them generally denser than liquids and gases (with some exceptions, like ice compared to water).

Low Compressibility:

• Because the particles are closely packed, solids cannot be easily compressed. They resist changes in shape and volume when pressure is applied.

Types of Solids:

• Crystalline Solids: Have an ordered, repeating structure and include categories like ionic, molecular, covalent network, and metallic solids. These solids have sharp melting points.

• Amorphous Solids: Lack a regular arrangement and have no defined melting point (e.g., glass, plastics). They soften gradually over a range of temperatures.

Properties of Solids

Melting Point:

• The melting point is the temperature at which a solid becomes a liquid. Like the boiling point in liquids, stronger intermolecular forces lead to higher melting points in solids.

Hardness and Brittleness:

• Hardness is the ability of a solid to resist scratching or indentation, while brittleness is the tendency to shatter. These properties depend on the type of bonds and structure—ionic and covalent network solids are often hard and brittle, while metallic solids are usually malleable and ductile.

Conductivity:

• Solids vary in conductivity:

  • Ionic Solids: Conduct electricity when melted or dissolved in water but not in their solid form.

  • Metallic Solids: Conduct electricity due to the presence of delocalized electrons.

  • Covalent Network and Molecular Solids: Generally poor conductors because they lack free-moving charged particles

Key Differences Between Liquids and Solids

Particle Arrangement: In solids, particles are in a fixed, orderly arrangement, while in liquids, they are close together but disordered.

Shape and Fluidity: Solids have a fixed shape, while liquids flow to assume the shape of their container.

Energy and Temperature: Solids require more energy to overcome their intermolecular forces and transition to liquid or gas phases compared to liquids transitioning to gas.

Multiple Choice

In which state of matter will there be the LARGEST space between molecules?

• Solid

• Liquid

• Gas

Multiple Choice

The diagram shows the arrangement of particles in a

• Solid

• Liquid

• Gas

The critical point is the highest temperature and pressure at which a pure material can exist in vapor/liquid equilibrium.

At temperatures higher than the critical temperature, the substance can not exist as a liquid, no matter what the pressure.

Multiple Choice

On a hot, humid summer day drops of water form on a can of cold soda. What process caused this to happen?

• condensation

• freezing

• vaporization

• deposition

Multiple Choice

When a gas forms directly from a solid, skipping the liquid phase. it is called

• deposition

• sublimation

• evaporation

• condensation

Multiple Choice

When a gas forms directly from a solid, skipping the liquid phase. it is called

• deposition

• sublimation

• evaporation

• condensation

Multiple Choice

What would the state of this substance be at 0.5 ATM of pressure at a 100 degrees Celsius?

• Solid

• Liquid

• Gas

• Michigan

Multiple Choice

What is the melting point of this substance at 1 ATM of pressure?

• 40

• 60

• 100

• 110

Multiple Choice

What letter on the diagram represents the triple point?

• A

• B

• C

• D

Multiple Choice

What letter on the diagram represents a solid?

• A

• B

• C

• D

Multiple Choice

What letter on the diagram represents a gas?

• A

• B

• C

• D

Intermolecular forces

​In chemistry, intermolecular forces are the forces of attraction or repulsion that occur between molecules.

These forces are generally weaker than the intramolecular forces (such as covalent or ionic bonds) that hold atoms together within a molecule. Intermolecular forces play a critical role in determining the physical properties of substances, such as boiling and melting points, solubility, and viscosity. Here’s an overview of the primary types of intermolecular forces:

​1. London Dispersion Forces (Van der Waals Forces)

• These are the weakest intermolecular forces and are present in all molecules, whether polar or nonpolar.

• They arise from temporary fluctuations in electron distribution around atoms or molecules, which create temporary dipoles. These dipoles induce a similar shift in neighboring molecules, leading to a weak attraction.

• London dispersion forces become stronger with larger atoms or molecules, as there are more electrons and a larger area for these temporary dipoles to form.

• Example: The forces between nonpolar molecules, like in noble gases or between molecules in substances like methane (CH₄).

Dipole-Dipole Interactions

• These forces occur in polar molecules that have permanent dipoles due to differences in electronegativity between atoms in the molecule.

• The positive end of one polar molecule is attracted to the negative end of another, leading to an overall attraction between molecules.

• Dipole-dipole interactions are stronger than London dispersion forces but are weaker than hydrogen bonds.

• Example: The interaction between HCl molecules, where the positive hydrogen end of one molecule is attracted to the negative chlorine end of another.

Hydrogen Bonding

• Hydrogen bonding is a specific type of dipole-dipole interaction but is stronger than typical dipole-dipole forces.

• It occurs when hydrogen is covalently bonded to a highly electronegative atom (such as nitrogen, oxygen, or fluorine), resulting in a strong partial positive charge on the hydrogen.

• This hydrogen can interact strongly with the lone pairs of electrons on nearby electronegative atoms, creating a strong intermolecular attraction.

• Hydrogen bonds are crucial in biological systems and contribute to the properties of water, which has a high boiling point relative to its molecular weight.

• Example: The strong attraction between water molecules (H₂O) due to hydrogen bonding.

Ion-Dipole Forces

• These forces occur between an ion and a polar molecule.

• They are especially significant in solutions where ionic compounds dissolve in polar solvents, such as salt (NaCl) in water.

• The positive or negative ends of the polar molecules (like water) are attracted to the opposite charge on the ions, stabilizing the ions in the solution.

• Ion-dipole forces are typically stronger than hydrogen bonds, making them one of the strongest types of intermolecular forces.

• Example: The attraction between Na⁺ ions and the partially negative oxygen in water molecules during the dissolution of salt in water.

Solids

Molecular Crystalline Solids

Composition: Molecular solids are composed of molecules held together by intermolecular forces like London dispersion forces, dipole-dipole interactions, and hydrogen bonds.

• Properties:

Low Melting and Boiling Points: Because the forces between molecules are relatively weak, these solids usually melt and boil at lower temperatures compared to ionic or covalent solids.

Poor Conductors: Molecular solids don’t conduct electricity well in any state because they lack free-moving charged particles.

Solubility: Their solubility varies depending on polarity. Polar molecular solids, like sugar, dissolve in polar solvents, while nonpolar ones, like iodine, dissolve in nonpolar solvents.

• Examples:

Ice (H₂O): Held together by hydrogen bonds.

Dry Ice (CO₂): Solid CO₂ held together by London dispersion forces.

Sucrose (Table Sugar): Made of polar molecules bound by dipole-dipole and hydrogen bonds.

Ionic Crystalline Solids

Composition: Ionic solids are composed of positively and negatively charged ions arranged in a repeating lattice structure, held together by strong electrostatic forces (ionic bonds).

• Properties:

High Melting and Boiling Points: Ionic bonds are very strong, requiring considerable energy to break, resulting in high melting and boiling points.

Brittleness: Ionic solids are brittle because shifting the lattice can bring like charges next to each other, causing repulsion and fracture.

Electrical Conductivity: They do not conduct electricity in the solid state but do when melted or dissolved in water, as ions are free to move.

Solubility: Many ionic solids are soluble in polar solvents like water because the solvent can stabilize the individual ions.

• Examples:

Sodium Chloride (NaCl): Common table salt.

Calcium Fluoride (CaF₂): Used in optics and metallurgy.

Magnesium Oxide (MgO): Used as a refractory material due to its high melting point.

Atomic Crystalline Solids

Atomic solids consist of individual atoms held together in a lattice, but they can be subdivided into three categories based on the type of bonding: network covalent, metallic, and noble gas solids.

a) Network Covalent Solids

Composition: Atoms are connected by a continuous network of covalent bonds, creating a very strong and rigid structure.

Properties:

• High Melting Points: Covalent bonds are very strong, so these solids have extremely high melting points.

• Hardness: They are typically hard and durable.

• Poor Conductors: Most lack free electrons or ions and thus don’t conduct electricity.

Examples:

• Diamond (C): Each carbon atom is covalently bonded to four others in a tetrahedral structure, making it extremely hard.

• Quartz (SiO₂): Composed of silicon and oxygen atoms in a rigid lattice.

• Graphite (C): Graphite is a unique network covalent solid where layers of carbon atoms slide over each other, making it soft and a good conductor.

Metallic Solids

• Composition: Metallic solids consist of metal atoms held together by a "sea" of delocalized electrons that move freely, allowing for conductivity.

• Properties:

  • Conductivity: The free electrons enable metals to conduct electricity and heat.

  • Malleability and Ductility: The flexibility of the electron "sea" allows metals to be shaped without breaking.

  • Variable Melting Points: Melting points vary widely depending on the metal and its bond strength.

• Examples:

  • Iron (Fe): Used in construction and manufacturing.

  • Copper (Cu): Known for its excellent electrical conductivity.

  • Gold (Au): Highly malleable and resistant to corrosion.

​Water is indeed a remarkable molecule with unique properties that arise from its molecular structure and hydrogen bonding. These properties make water essential to life and influential in many natural processes.

• Molecular Structure and Polarity: Water (H₂O) consists of two hydrogen atoms covalently bonded to one oxygen atom. The oxygen atom is more electronegative than hydrogen, creating a polar molecule with a bent shape. This polarity means the oxygen side of the molecule has a partial negative charge, while the hydrogen side has a partial positive charge. This arrangement enables water molecules to interact strongly with each other and with other polar molecules.

• Hydrogen Bonding: The partial charges on water molecules lead to hydrogen bonding, where the positive hydrogen of one water molecule is attracted to the negative oxygen of a neighboring molecule. Hydrogen bonding is responsible for many of water’s unique properties, including its high boiling point, surface tension, and ability to dissolve many substances.

• High Boiling and Melting Points Compared to other small molecules, water has an unusually high boiling and melting point due to hydrogen bonding. It takes a significant amount of energy to break these bonds, keeping water in a liquid state over a wide temperature range. This property is crucial for life, as it enables water to exist as a liquid in most environments on Earth.

• High Specific Heat Capacity Water has a high specific heat capacity, meaning it can absorb a lot of heat energy without a significant temperature change. This stabilizes climates and allows organisms to maintain a constant internal temperature. In natural environments, large bodies of water moderate temperatures, preventing extreme fluctuations that would make many ecosystems unsustainable.

High Surface Tension Due to hydrogen bonding, water molecules are highly cohesive, giving water an unusually high surface tension. This property enables phenomena like capillary action (important in plant water transport) and allows certain insects to "walk" on water.

Density Anomaly of Ice Most substances become denser as they cool and solidify, but water is an exception. Ice is less dense than liquid water due to the open, hexagonal structure of water molecules in the solid state, caused by hydrogen bonds.This lower density allows ice to float on water, insulating aquatic life during cold seasons and preventing bodies of water from freezing solid from the bottom up.

Excellent Solvent: Water is often called the "universal solvent" because it can dissolve a wide variety of substances, particularly ionic and polar molecules. This ability stems from water’s polarity, which allows it to surround and separate ions and molecules, breaking their intermolecular bonds. In biological systems, this solvent property is crucial for transporting nutrients, gases, and wastes in cells and organisms.

Transparency Water’s transparency to visible light is essential for aquatic ecosystems. Sunlight can penetrate water, allowing photosynthesis to occur in aquatic plants and supporting life in oceans and freshwater systems.

Role in Chemical Reactions Water is an essential reactant in many biochemical reactions, including hydrolysis and dehydration synthesis. It also plays a role as a medium for reactions and as a source of hydrogen and oxygen in photosynthesis and cellular respiration.

High Heat of Vaporization Water’s high heat of vaporization means that it requires a lot of energy to transition from liquid to gas. This property is essential for evaporative cooling, a process organisms use to regulate temperature through sweating and transpiration in plants.

Multiple Choice

Which of the following will take the longest to evaporate?

• CH₃CH₂OH
• CH₃OH
• CH₃CH₃

Multiple Choice

Intermolecular forces are the forces

• within molecules
• between molecules

Multiple Choice

H₂S has what kind of intermolecular force?

• dipole dipole
• dispersion
• H-bond
• ionic

Multiple Choice

Type of intermolecular force present in HF.

• dipole dipole
• dispersion
• H-bond
• ionic

Multiple Choice

Changing from a liquid to a gas

• condensation

• vaporization

• sublimation

• melting

Multiple Choice

Which grouping lists the states of matter in order of increasing distance between particles?

• gas, liquid, solid

• liquid, solid, gas

• solid, gas, liquid

• solid, liquid, gas

Multiple Choice

Which grouping lists the states of matter in order of increasing distance between particles?

• gas, liquid, solid

• liquid, solid, gas

• solid, gas, liquid

• solid, liquid, gas