X Unit 13 States of Matter

Reviewing States of Matter 2) 1) 3)

States of Matter  Solid Has definite shape & volume  Liquid Has constant volume but takes the shape of its container  Gas Expands to fill its container & takes the shape of its container

Forces of Attraction  There are two kinds of attractive forces at the molecular level –  The forces inside a molecule holding the individual atoms together  The forces between molecules holding different molecules together in a sample

Intramolecular Forces  “Intra-” prefix = within  The forces inside a molecule holding the individual atoms together Ex.) Covalent bonds in H 2 O

Intermolecular Forces  “Inter-” prefix = between  Short range forces between molecules in a sample  There are 3 main types of intermolecular forces Hydrogen bonding Dipole-dipole forces London Dispersion forces

Dipole-Dipole Forces

 Dipole – a molecule or part of a molecule that contains both positively and negatively charged regions A.k.a. a polar molecule Dipoles have δ + (partial positive) & δ - (partial negative)  Dipole-Dipole Forces – forces of attraction between POLAR molecules Dipoles must get close together in correct orientation (positive end must be near negative end)

Dipole-Dipole Forces  Dipole-dipole forces will raise melting and boiling points.  A dipole can temporarily attract electrons from another molecule causing an induced dipole.

London Dispersion Forces

 Intermolecular attractions resulting from the uneven distribution of electrons and the creation of temporary dipoles Present in all substances (polar molecules, non- polar molecules, and noble gases) The weakest intermolecular force  Does the number of electrons play a role regarding strength of the attraction??

London Dispersion Forces  Electrons are constantly moving around the nucleus therefore electron density can fluctuate This effect becomes stronger with increasing number of electrons  Example: F 2 Br 2 I 2

Hydrogen Bonding

 Attractive forces in which a hydrogen covalently bonded to a very electronegative atom (F, O, N) is also weakly bonded to an unshared electron pair on another electronegative atom (another F, O, or N atom) Hydrogen Bonding

 Elements that undergo H-bonding Hydrogen bonding is FON! (Fluorine, Oxygen, and Nitrogen)  Hydrogen bonding raises melting and boiling points because more energy is required to break the forces between molecules. NO Hydrogen Bonding!!!!

Hydrogen Bonding – Effects on Physical Properties  H 2 O is one of the most notable examples of hydrogen bonds Ice forms rigid, open structures  Increases volume upon freezing (floats) Molecules w/ higher molar mass have lower BP than H 2 O

Hydrogen Bonding in DNA  Hydrogen bonding plays a key role in maintaining the double helix structure of DNA

Example: Identify the types of intermolecular forces present in compounds of: Hydrogen Fluoride (HF) Pentane (C 5 H 12 ) Hydrochloric Acid (HCl) Ethanol (Ethyl Alcohol)

Relative Strength of Intermolecular Forces (Weakest)(Strongest)

Liquids & Solids Even though they are more restricted than gas molecules, the molecules of solids and liquids are constantly in motion as well! (This idea comes from the Kinetic Molecular Theory – we’ll come back to this idea)

Liquids  Viscosity – a measure of the resistance of a liquid to flow The particles in a liquid are close enough together that their attractive forces slow their movement as they flow past one another The stronger the attractive forces (intermolecular forces), the more viscous the liquid is.  As temperature increases, viscosity decreases.

 Surface tension – an inward force that tends to minimize the surface area of a liquid A measure of the inward pull by particles in the interior The stronger the intermolecular forces, the higher the surface tension Liquids In water, this is due mainly to hydrogen bonding!

Liquids  Surfactant – any substance that interferes with the hydrogen bonding between water molecules & reduces surface tension

Surfactants used to clean up oil spills Exxon Valdez oil spill in 1989 spilled over 700,000 barrels of oil into the water near Alaska

Classes of Solids  Crystalline Network Covalent Metallic Molecular Ionic  Amorphous

Solids  Crystalline solid – a solid in which the atoms, ions, or molecules are arranged in an orderly, geometric, three-dimensional structure  Unit cell – the smallest arrangement of connected points that can be repeated to form the lattice A.k.a. The representative part of the whole crystal

Crystalline Solids Pyrite (cubic) Rutile (tetragonal)

Crystalline Solids Borax (monoclinic) Copper sulfate (triclinic)

Network Covalent Solids Atoms that can form multiple covalent bonds (such as C, Si, and other Group 14 elements) are able to form network covalent solids. All atoms in the entire structure are bonded together with covalent chemical bonds.

Metallic Solids  Metallic solids – positive metal ions surrounded by a sea of mobile electrons Mobile electrons make metals malleable and ductile because electrons can shift while still keeping the metal ions bonded in their new places  Metallic solids are good conductors of heat and electricity Metallic Bonds Video Clip – Metals & Metallic Bonds

Amorphous Solids  Amorphous solid – a solid in which the particles are not arranged in a regular, repeating pattern “Amorphous” = “without shape” Often form when a molten material cools too quickly to allow enough time for crystals to form  Common examples: glass, rubber, many plastics

Molecular Solids  Formed when atoms are bonded together with covalent bonds  Forms a molecule Properties of Molecular Solids:  Formed from 2 nonmetals  Do not conduct electricity (can be used as insulators)  Low melting & boiling points

Ionic Solids  Formed from attraction between cations & anions  Properties of Ionic Solids:  Formed from metal (cation) & nonmetal (anion)  Strong electrolytes – will conduct electricity when dissolved in water (or when molten)  High melting & boiling points  Forms crystal lattice

Ionic Solids Forms crystal lattice made of formula units Ionic formula units show the ratio of cation to anion in the crystal lattice.

Gases & the Kinetic Molecular Theory  KMT describes the behavior of gases (and solids/liquid to some extent) in terms of particles in motion Kinetic = movement  Assumptions: 1. Gas particles have negligible volume compared to the volume of their container 2. Particles move in constant, random, straight line motion 3. Particles collide with themselves and walls without losing energy (elastic collisions) 4. There are no intermolecular forces between gas molecules

Kinetic Molecular Theory  If gas particles are always in this constant, random motion, what keeps them going? ENERGY!! (Kinetic Energy to be exact)  Temperature is a measure of the average kinetic energy in a substance. Higher temp. = more kinetic energy = particles move faster!

Temperature  3 Main Temperature Scales Fahrenheit Celsius (°C)  °C = K Kelvin (K)  K = °C + 273

Temperature Conversions  Convert the following temperatures: 28 °C = ________ K 200 K = ________ °C -15 °C = _________ K

Remember…  Gases expand to fill their container & take the shape of their container In other words, gases will spread out until they can’t spread out anymore Gases will also move according to the principles of diffusion and effusion.

Racing Gases: If concentrated HCl is at one end of the tube and concentrated NH 3 is at the other end, which gas do you think will move farthest and fastest down the tube? HCl (g)NH 3 (g)

RACING GASES  The gases will diffuse down the tube Diffusion – tendency of molecules to move from areas of higher concentration towards areas of lower concentration  Example: spraying perfume and smelling it across the room Originally Over Time

RACING GASES  The gases diffused at different rates Why does NH 3 move further and faster? BECAUSE IT’S LIGHTER!! (Has a lower molar mass)

Graham’s Law of Effusion  Graham’s Law of Effusion – gases of lower molar masses effuse (and diffuse) faster than gases with higher molar masses Effusion – when a gas escapes (diffuses) through a tiny hole in its container  Example: Helium balloons shrinking compared to normal balloons

Graham’s Law  Graham’s Law be applied to both the effusion and the diffusion of a gas Gases with lower molar masses (lighter gases) diffuse faster than gases with higher molar masses (heavier gases) The lighter the gas, the faster it moves!!!

Graham’s Law Which gas would diffuse and effuse faster… 1. Methane (CH 4 ) or carbon dioxide (CO 2 )? 2. Chlorine (Cl 2 ) or oxygen (O 2 )? 3. Hydrogen sulfide (H 2 S) or carbon monoxide (CO)?

Graham’s Law Formula The rates are simply the speed or velocity at which the gas is traveling. So, this formula will compare the speed of one gas to the speed of the other gas.

Gas Pressure  Pressure of a gas = the force of a gas exerted on the surface of a container  As gases bounce around in a container, each collision with a container wall exerts a force  More collisions = higher pressure  Less collisions = lower pressure  Vacuum = Empty space with no particles (no pressure)

Pressure of a Gas  Atmospheric pressure – collision of air molecules with objects As elevation increases, air density and therefore pressure decrease  Barometers measure atmospheric pressure

Pressure of a Gas  Vapor pressure – pressure due to force of gas particles above a liquid colliding with walls of container  Higher temp. = higher vapor pressure

Units of Pressure  SI unit of pressure: pascal (Pa)  Other common pressure units: Millimeters of mercury (mm Hg) Atmospheres (atm) 1 atm = 760 mmHg = kPa = 760 torr

 Partial pressure – the contribution of each gas in a mixture to the total pressure  Dalton’s Law of Partial Pressures – for a mixture of gases, the total pressure is the sum of the partial pressure of each gas in the mixture P total = P 1 + P 2 + P 3 + … (at constant volume and temperature) Dalton’s Law of Partial Pressures

Practice with Dalton’s Law Determine the total pressure of a gas mixture that contains oxygen, nitrogen, and helium. The partial pressures are: P O 2 = 20.0 kPa, P N 2 = 46.7 kPa, and P He = 26.7 kPa. P total = P 1 + P 2 + P 3 + …

Practice with Dalton’s Law Air contains O 2, N 2, CO 2, and trace amount of other gases. What is the partial pressure of oxygen (P O 2 ) if the total pressure of the system is kPa and the partial pressures of N 2, CO 2, and the other gases are kPa, kPa, and 0.94 kPa, respectively? P total = P 1 + P 2 + P 3 + …

Phase Changes What are phase changes? A change in a substance’s state of matter What are some examples of phase changes?

Phase Changes that Require Energy  If you have to put energy into a reaction to make it happen, it is called an endothermic reaction.  Endothermic Phase Changes: Melting a.k.a “fusion”(solid  liquid) Vaporization (liquid  gas) Sublimation (solid  gas)

Phase Changes that Release Energy  If energy is released or given off by a reaction, it is called an exothermic reaction.  Exothermic Phase Changes: Condensation (gas  liquid) Deposition (gas  solid) Freezing (liquid  solid)

Heating Curve