X Unit 13 States of Matter

Reviewing States of Matter

Reviewing States of Matter 1) 2) 3)

Solid States of Matter Has definite shape & volume Particles are tightly packed Can expand when heated

Liquid States of Matter Has constant volume but takes the shape of its container Fluid Less closely packed particles than solid particles Can expand when heated

States of Matter Gas Expands to fill its container & takes the shape of its container Fluid Much less closely packed than solid particles Expands when heated

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 H2O

Demo What happened to the paper clip when placed in the beaker of water vs. the beaker of acetone? Explain your observations.

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-Dipole Forces Dipole – a molecule or part of a molecule that contains both positively and negatively charged regions δ+ (partial positive) or 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 H-F H-F H-F H-F 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

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: F2 Br2 I2

Hydrogen Bonding

Hydrogen Bonding 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 bonds are strong intermolecular forces

Hydrogen Bonding Hydrogen bonding raises melting and boiling points because more energy is required to break the forces between molecules.

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

Think Back to the Demo Question… Why did the paperclip float on water but not acetone?

Before You Go: Identify the types of intermolecular forces present in compounds of: Hydrogen Fluoride Pentane (C5H12) Hydrochloric Acid Ethanol (Ethyl Alcohol)

Relative Strength of Intermolecular Forces (Strongest) (Weakest)

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.

In water, this is due mainly to hydrogen bonding! Liquids 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 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 as well Exxon Valdez oil spill in 1989 spilled over 700,000 barrels of oil into the water near Alaska

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 Copper sulfate (triclinic) Borax (monoclinic)

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

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

Kinetic Molecular Theory Describes the behavior of gases (and solids/liquid to some extent) in terms of particles in motion Assumptions: Gas particles have negligible volume compared to the volume of their container Particles move in constant, random, straight line motion Particles collide with themselves and walls without losing energy (elastic collisions) 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 Celsius (°C) Kelvin (K) 3 Main Temperature Scales Fahrenheit Celsius (°C) °C = K - 273 Kelvin (K) K = °C + 273

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

Gases 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 diffusion and effusion.

Racing Gases Demo Racing Gases Demo: If concentrated HCl is at one end of the tube and concentrated NH3 is at the other end, which gas do you think will move farthest and fastest down the tube? Racing Gases Demo HCl (g) NH3 (g)

RACING GASES DEMO 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

DIFFUSION Originally Over Time

(Has a lower molar mass) RACING GASES DEMO The gases diffused at different rates If the white ring forms closer to the HCl end of the tube, which gas moved farthest and fastest? Why did this gas move faster? BECAUSE IT’S LIGHTER!! (Has a lower molar mass)

Graham’s Law of Effusion The racing gases demo is related to Graham’s Law 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

The lighter the gas, the faster it moves!!! 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… Methane (CH4) or carbon dioxide (CO2)? Chlorine (Cl2) or oxygen (O2)? Hydrogen sulfide (H2S) 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.

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 the a container wall exerts a force More collisions = higher pressure Less collisions = lower pressure Empty space with no particles and no pressure is called a vacuum

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

Pressure of a Gas When baking, there are different instructions for baking at high altitudes…why? Boiling point of water decreases since lower pressure Liquid leaves faster so food must bake longer Gravity

Pressure of a Gas SI unit of pressure: pascal (Pa) Other common pressure units: Millimeters of mercury (mm Hg) Atmospheres (atm) 1 atm = 760 mmHg = 101.3 kPa = 760 torr

Practice Converting Units 1 atm = 760 mmHg = 101.3 kPa A tire pressure gauge records a pressure of 450 kPa. What is the pressure in atmospheres? In mm Hg?

DALTON’S LAW

(at constant volume and temperature) Dalton’s Law 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 Ptotal = P1 + P2 + P3 + … (at constant volume and temperature)

Practice with Dalton’s Law Determine the total pressure of a gas mixture that contains oxygen, nitrogen, and helium. The partial pressures are: PO2= 20.0 kPa, PN2= 46.7 kPa, and PHe= 26.7 kPa. Ptotal = P1 + P2 + P3 + …

Practice with Dalton’s Law Air contains O2, N2, CO2, and trace amount of other gases. What is the partial pressure of oxygen (PO2) if the total pressure of the system is 101.3 kPa and the partial pressures of N2, CO2, and the other gases are 79.10 kPa, 0.040 kPa, and 0.94 kPa, respectively? Ptotal = P1 + P2 + P3 + …

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 (solid  liquid) a.k.a “fusion” 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