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Chapter 11 Reading Notes

by: Hannah Huffman

Chapter 11 Reading Notes CHE107

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Hannah Huffman


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These notes cover both class material and material in the book, and I have made it easier to understand by using normal person language!
Gen Chem 2
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This 16 page Class Notes was uploaded by Hannah Huffman on Monday February 1, 2016. The Class Notes belongs to CHE107 at a university taught by Blue in Spring 2016. Since its upload, it has received 33 views.


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Date Created: 02/01/16
Chapter 11 Notes ~ Book Notes ~ Class Notes States of Matter - By changing Temperature and Pressure, most substances can exist as solid, liquid and Gas - Forces of attraction determine physical state o Strong attractions cause solids and liquids. o Weak attractions cause gas. - Phase: homogenous (same chemical composition) part of a system is in contact with other parts of a system, but is separated by well-defined boundary. o Ex: 2 phases of water in a cup (ice and liquid water) 3 Common Phases of matter - Solids o Fixed: keeps shape when placed in a container o Compressible? No o Strength of IMF  Very strong o Density: high o Closely packed, fixed in position o Able to vibrate o Close packing  Allows substance to remain dense  Incompressible (retain volume)  Inability to move:  Why it retains shape  Does not flow o Crystalline Solids:  Molecules are geometrically ordered in 3D array.  Means that you can connect the dots diagonally and horizontally.  Ex: salt, sugar, diamond o Amorphous Solids:  No long range order.  IMFs are very strong so it doesn’t have to have a specific order.  Ex: glass, plastic, wax. - Liquids o Indefinite: takes shape of container o Volume: fixed. o IMF: intermediate to strong o Density: low o Compressible: No o Particles in a liquid are:  Closely packed  Less close than solids  Mobile in position  Able to vibrate easily Chapter 11 Notes ~ Book Notes ~ Class Notes o Close packing:  Means they are dense and hard to compress.  Retains volume. o Ability to move:  Adopt shape of container.  Flow. - Gas o Ideal Gas Law  PV=nRT  P= pressure in atm  V= volume in L  n= moles  T- temperature in K  R= ideal gas constant= .08206 (L*atm)/(mol*K)  Corrected to include IMF’s= (V-nb)= (nRT/P)  B=constant that depends on the gas. o Indefinite shape and volume o Density: low o Volume: indefinite o Compressible? Yes o IMF’s: weak o Particles in a gas are:  Loosely packed (large spaces between particles)  Mobile  Collide with one another and container  Rapidly vibrating o Loose packing:  Low density  Molar volume:  Compressible o Ability to move:  Adopt shape of container  Expand to fill container  Flow Degrees of Freedom (of motion of particles) - Translational Freedom: ability to move from one position to another. o Gas has large freedom - Rotational Freedom: ability to change orientation of particle without changing center of gravity. - Vibrational Freedom: ability to oscillate about a point in space. Kinetic-Molecular Theory (Chapter 5) State of a material depends largely on: 1. Kinetic energy (KE) of particles 2. Strength of attraction between particles. - These two factors oppose one another Chapter 11 Notes ~ Book Notes ~ Class Notes o Molecules in gases have complete freedom of motion  KE >> attractive forces o Molecules in a solid are locked in place  KE < attractive forces  Still vibrate but cannot escape o As temperature goes down, KE slows down enough for IMF’s to kick in making it go to a liquid (in H2O).  Therefore, IMF’s play a part in determining the substance’s state. Conversion between States - Condense o Gas to solid  Ex: frost - Freeze o Liquid to solid - Condense o Gas into liquid - Melt (fuse) o Solid to liquid - Boil (evaporation) o Liquid to gas - Sublimation o Solid to gas  Ex: dry ice Intermolecular Forces: attractive forces between molecules (atoms) - KE increases with Temperature - Attractive forces constant with T - Attractive forces << ionic attractions or covalent bonds o Attractive forces holding water molecules together are normally weaker than bonds holding individual water molecules together. - Electrostatic attractions based on charges, partial charges, and temporary charges. o Partial charges of some whole molecules attract to negative partial charges of other molecules. - Attractions listed in decreasing strength 1. Ion-ion/Bond  q: full charge r:short distance (100-200pm)  Attraction between two ions.  Strong enough to be called a bond.  Between two ions.  Ionic bond: bond where there is a complete transfer of electrons, resulting in oppositely charged ions.  Covalent bond: Bond where electron pairs are shared between atoms.  Polar covalent bond: electrons are shared unevenly causing a partial negative and positive charge. Chapter 11 Notes ~ Book Notes ~ Class Notes  Nonpolar covalent bond: electrons are shared equally between two atoms. 2. Ion-dipole a. q: partial charge r:longer distance (200-400pm) b. Positively charged end of a polar molecule is attracted to negative charged ions. i. Hydration c. Ion and polar molecule. d. Occurs when an ionic compound is mixed with a polar compound. i. Positive and negative ions react with partial positive and negatives of other molecules in the compound. e. Responsible for substances ability to make solutions with water. 3. Hydrogen Bonding a. Happens when Hydrogen is directly bonded through IMF’s to Fluorine, Oxygen, or Nitrogen. b. IMF’s are strong because of the fairly large partial positivity of hydrogen, and fairly large partial negativity of F, O, or N. Also, these atoms are small, so they can get very close to each other. c. Chemical bonds: bonds between atoms. d. IMFs: bonds between separate MOLECULES. i. Hydrogen bonding is an IMF. e. Ex: on right, partial positivity of H causes partial negativity on the opposite side, causing partial positivity of next H to be attracted to the other molecule. i. H is partially positive because it is donating an electron, F is partially negative because it is receiving an electron. 4. Dipole-dipole a. Exists in all polar molecules. i. Polar molecules: have an uneven distribution of electrons creating a permanent dipole. ii. Polar bond: uneven sharing of electrons between all molecules in a compound, creates a dipole moment. iii. Permanent dipole: positive or negative end of an atom that interacts with surrounding atoms creating a dipole-dipole attraction. iv. Greater dipole moment (separation between positive and negative charges) causes stronger dipole-dipole attraction resulting in a higher boiling point. b. Miscibility: the ability to mix without separating into two states. Chapter 11 Notes ~ Book Notes ~ Class Notes i. Polar liquids are miscible with other polar liquids, but not with nonpolar liquids. 5. Dipole-induced dipole 6. Dispersion forces (with model) a. Present in all molecules. b. Caused by changes in electron distribution/position in molecules. c. More electrons means stronger dispersion forces. i. Results from increasing molar mass that makes it more polarizable. ii. More electrons make valence electrons farther from control of nucleus. d. Results from asymmetrical distribution of electrons or TEMPORARY dipole. i. Induced dipole: temporary dipole cause other “induced” dipoles that are attracted to the positive and negative side of the other neighboring molecules. e. Instantaneous/temporary dipole: change in charge separation in a molecule i. Ex: due to different positions of electrons, the right side or the left side of the molecule may have a slightly negative or positive charge. f. When one side of the atom is positive, it attracts the negative side of the neighboring atom, vice versa. g. Magnitude of dispersion force depends on how easily atoms can move or polarize. i. Larger electron clouds are more polarizable because more electrons are not held as tightly to the nucleus. ii. Therefore, dispersion forces increase with more molar mass. h. Longer molecules have stronger dispersion forces because there is more area for attraction and IMF’s. Review: Bonding - Ionic: o Metal and nonmetal o Electrons transferred - Covalent: o Nonmetal and nonmetal Chapter 11 Notes ~ Book Notes ~ Class Notes o Electrons shared  Since nonmetals have high ionization energies, electrons are not easy to remove so they are shared. - Metallic: o Metal and metal o Electrons pooled - Polar Covalent Bonds o Electronegativity: the ability for an atom to attract electrons to itself.  Increases going across and up the periodic table (Fluorine is the most electronegative).  With the exception of hydrogen, which is relatively electronegative.  Polarity in bonds depends on the electronegativity difference between bonding atoms.  Greater the difference, the more polar (uneven) the bond is. o Ex: C2 is the same element, so the bond is covalent, NaCl is far apart in electronegativity, so the bond is ionic, but HCl is polar covalent because there is an intermediate difference in their electronegativity’s. o Ex: Fluorine would be more negative than hydrogen when they bond because Fluorine has a higher electronegativity which means it will attract more electrons to itself. IMF’s in Action: Surface tension, Viscosity, and Capillary Action Surface Tension: the tendency of liquids to minimize their surface area. The energy required to increase the surface area of a liquid by a unit amount. - Surface molecules are less stable (more Potential Energy) because it has less neighboring molecules to interact with. o Molecules want to be stable, so they tend to try to minimize surface area (or minimize the amount of unstable molecules). o For things to sink, more water molecules must come to the top, which requires energy, so water and other liquids resist sinking things. - Surface tension decreases as IMF’s decrease. o Spheres minimize SA by volume, so many liquids form spheres due to IMF’s. Viscosity: the resistance of a liquid to flow. - Measured in Poise (P) which is defined as 1g/cm*s - Viscosity increases with IMF’s. o Also increases sometimes with molar mass b/c of greater dispersion forces. Chapter 11 Notes ~ Book Notes ~ Class Notes - Viscosity also increases in longer molecules because they have a greater area that can become entangled. - Viscosity decreases as temperature increases. o Because temperature can overcome IMF’s. Capillary Action: the ability of a liquid to flow against gravity, up a narrow tube. - Results from a combination of two forces 1. Cohesive forces: attraction between molecules in a liquid. a. Causes liquid to stay together and go up the tube when some molecules are not touching the wall. 2. Adhesive forces: attraction between molecules and the surface of a tube. a. Causes a liquid to spread out over the surface of a tube. b. If adhesive forces are less than cohesive forces, liquid does not rise up at all. Vaporization and Vapor Pressure Vaporization: the process by which thermal energy can overcome IMF’s to change a liquid to a gas. - Process of Vaporization: o The higher the temperature, the greater the average energy of the collection of molecules. o Weaker IMF’s increase the rate of vaporization.  Surface molecules and molecules with enough energy break free and become a gas.  When water is spilled, more vaporization takes place because more surface area, more heat, more vaporization takes place b/c there is more KE. o Condensation: transition from gas to a liquid.  Evaporation usually takes place more quickly because the gas escapes into the atmosphere. o Volatile: evaporates easily. o Nonvolatile: evaporated not as easily. - The Energetics of Vaporization o Vaporization is an endothermic process which means it TAKES energy to vaporize molecules. o Condensation is exothermic: energy in the form of heat is released when a gas condenses to a liquid.  Ex: steam burn: as steam condenses to liquid on your hand, it releases heat. o Heat (or enthalpy) of Vaporization (∆H vap: amount of heat required to vaporize one mole of a liquid to a gas.  Heat of vaporization of water: +40.7 kJ/mol  Heat of vaporization is always positive because it is an endothermic process.  At lower temperatures, heat of vaporization is slightly greater because the substance contains less original heat energy. Chapter 11 Notes ~ Book Notes ~ Class Notes o Condensation: when a substance condenses, the same amount of heat is RELEASED instead of used.  Condensation of water: -40.7 kJ/mol  Sign is negative, because process is exothermic and releases heat. o Heat capacity: the amount of heat it takes to raise the temperature of an object by one degree.  q=C s(∆T)  q= energy  Cs= specific heat  m= mass  ∆T= (T FTI o Final temperature-initial temperature. o The longer the carbon chain, the higher the boiling point because there is a bigger surface you have to heat up. Vapor Pressure and Dynamic Equilibrium - Dynamic Equilibrium: the rate of vaporization and condensation are equal. - Vapor pressure: the pressure of a gas in dynamic equilibrium with its liquid. o Depends on IMF’s.  Weak IMF’s have high vapor pressure b/c it is easier for the substance to become a vapor (volatile).  Also IMF’s are easy to overcome with thermal energy. o Substances tend to get back to dynamic equilibrium even if the substance is disturbed (for example, if the volume of the container is changed).  If volume becomes greater, pressure initially drops, then returns to dynamic equilibrium, vice versa. - Temperature Dependence of Vapor Pressure and Boiling Point o When temperature rises, vapor pressure rises b/c more molecules have enough energy to vaporize. o Boiling point: temperature at which the liquids vapor pressure equals the external pressure. o Normal boiling point: the temperature at which vapor pressure equals 1 atm (760 torr)  Additional heating simply makes it boil more rapidly, it does not change the temperature of the liquid above the boiling point.  As long as liquid water is present, the temperature cannot exceed 100 degrees C. o After all water is converted to steam, the temperature can raise. - Clausius-Clapeyron Equation: shows the exponential relationship between vapor pressure and temperature. Chapter 11 Notes ~ Book Notes ~ Class Notes ln P = −∆H vap 1 +lnβ o vap R (T o Mimicked y=mx+b form  P = vapor pressure vap  β = a constant depending on the gas.  ∆ H vap heat of vaporization  R = gas constant(8.314J/molK)  T = Temperature in K o Even though the relationship between temperature and vapor pressure isn’t linear, the inverse of temperature and the natural log of vapor pressure are. o Measuring heat of vaporization in lab: measure vapor pressure as a function of temperature and plot the natural log of vapor pressure versus inverse of temperature, slope of the line is the heat of vaporization.  Graph on next page gives an example.  Remember to check units! - Two-point form of Clausius-Clapeyron Equation P2 −∆ H vap1 1 o lnP = T (T − T ) 1 2 1 - Can be used if we want to predict vapor pressure at any temperature if we know the enthalpy of vaporization and normal boiling point (or the vapor pressure at some other temperature). - Converting from C to K o K=Degrees C+ 273.15 The Critical Point: the Transition to an Unusual State of Matter - Supercritical Fluid: neither a liquid nor a gas. o Works as a good solvent (dissolver). Chapter 11 Notes ~ Book Notes ~ Class Notes o Occurs at the Critical Temperature (T ) c  Liquid form cannot occur over this temperature, regardless of pressure. o Occurs at critical pressure (P c Sublimation and Fusion Sublimation: molecules have enough energy to break free from a solid and go directly to a gas. - Happens at a quicker rate than deposition because gas particles escape into the atmosphere. Deposition: opposite of sublimation, goes directly from gas to solid. Fusion: - Melting point (fusion/melting): Increased temperature make molecules vibrate faster and faster until they have enough energy to overcome molecular forces. o The substance does not raise above 0 degrees Celsius until all the solid is melted. o Melting is endothermic which means it absorbs thermal energy from the surroundings. Energetics of Melting and Freezing - Heat of Fusion: the amount of heat it takes to melt 1 mol of a solid. o For water it is: 6.02kJ/mol - To freeze, the numbers are the same, but the sign is opposite since freezing is exothermic. o For water: -6.02kJ/mol Heating Curve for Water - Amount of heat required to achieve the state change: q=n∆H Chapter 11 Notes ~ Book Notes ~ Class Notes A-B: solid ice is warmed from -25°−0° C - Since there is not a state change, amount of heat required is found by: o q=mC ∆s o Cs, ice.09J/g ℃ B-C: added heat does not change temperature because all heat is absorbed by state change. - Amount of heat to change ice to water is given by: ∆ H o q=n fus  ∆ Hfuseatof fusion  Heat of fusion for water= 6.02kJ/mol C-D: liquid water is warmed from 0 ° to 100 ℃ . - Amount of heat given by: o q=mC ∆s o However, now we use heat capacity of liquid water which is 4.18J/g ℃ . D-E: water transitions from liquid to gas. - Amount of heat given by: o q=n ∆ H vap Chapter 11 Notes ~ Book Notes ~ Class Notes o ∆ Hvap for water is 40.7 kJ/mol E-F: steam is warmed. - Heat required is given by: o q= mC ∆s o Heat capacity of steam: 2.01 J/g ℃ When doing problems, remember to keep units consistent and add all the steps together at the end! Phase Diagrams: a map of the state or phase of a substance as a function of pressure and temperature. Major Features of a Phase Diagram - Regions: solid, liquid, and gas regions. o Represents conditions where that particular state is stable.  Low temperature and high pressure favor solids.  High temperature and low pressure favor the gas state.  Intermediate conditions favor the liquid state. - Lines: represent the temperature and pressures where substances are in equilibrium. o Sublimation curve (dividing solid and gas). o Fusion curve: separating solid and liquid. - Triple point: three states are equally stable and at equilibrium. - Critical Point: represents the temperature and pressure at which there is a supercritical fluid (liquid and gas form it). Crystalline Solids: composed of atoms or molecules arranged in structures with long range order. - X-ray diffraction: a way to look at and measure distance between atoms and molecules in crystalline solids. o Interference: electromagnetic (light) waves interact with each other by cancelling each other out or reinforcing each other.  Constructive Interference: two waves interact with their crests and troughs in alignment.  Destructive Interference: two waves interact with the crests of one aligning with troughs of another.  Interference pattern: alternate bright and dark lines made by interference when travelling through 2 slits comparable to wavelength.  Bragg’s Law: nγ=2dsinθ  Act like γ is upside down. Crystalline Solids: Unit Cells and Basic Structures - Crystalline Lattice: arrangement of atoms within a crystalline solid. Chapter 11 Notes ~ Book Notes ~ Class Notes - Unit cell: represents the crystalline lattice. - Lattice point: the point in space occupied by an atom, ion, or molecule. - Cubic unit cell: equal edge lengths and 90 degree corner angles. o Simple cubic:  Cube with one atom at each corner.  Edge length=2r  r= radius of atom  Coordination Number: 6 o Body-centered cubic:  A cube with one atom at each corner and one at the center of the cube.  Atoms touch along diagonal line running from one corner, across the cube, to the other, NOT across the edge.  Coordination number: 8 4r  Edge Length: √3 o Face-centered cubic:  One atom in each corner and one atom in the center of each cube face.  Atoms touch along the diagonal FACE NOT the edge.  Edge length= 2√2r  Coordination Number: 12 - Coordination Number: the number of atoms with which each atom is in contact with. - Packing Efficiency: percentage of the volume of the unit cell occupied by the spheres. o Higher coordination number, higher packing efficiency. Chapter 11 Notes ~ Book Notes ~ Class Notes Molar Mass/ Avogadro’s number (6.022E23 gives you g/atom. 1 pm= 1E-12 m Crystalline Solids: The Fundamental Types - Molecular Solids: o Composite units are molecules.  Ex: H2O and CO 2 o Held together by IMF’s o Low to moderately low melting points depending on IMF’s. - Ionic Solids o Composite units are ions.  Ex: NaCl o Coordination number represents cation-anion interactions. o All unit cells must be neutrally charged b/c the structure tries to decrease PE making it more stable. o If cations and anions are similar radii, or size, the higher the coordination number. o Held together by Coulombic forces (ionic bonds) so they have much higher boiling points. - Atomic Solids o Composite units are individual atoms. o Ex: solid Xenon (Xe), iron (Fe) and SiO2  Nonbonding atomic solids:  Held together by dispersion forces.  Form close-packed solids to maximize coordination number and minimize space between each other.  Low melting points that increase with molar mass. o Ex: noble gases in their solid form.  Metallic atomic solids:  Metallic Bonds o Held together by metal cations and sea of electrons that surround them.  Form close-packed structures.  Varying strengths, so varying melting and boiling points.  Network Covalent atomic solids:  Held together by covalent bonds  Do not form close-packed structures, because covalent bonds tend to be more directional.  Does not conduct electricity.  High melting point. Crystalline Solids: Band Theory - Band theory: combining of atomic orbitals of atoms within a solid crystal to form orbitals that are delocalized over the entire crystal. Chapter 11 Notes ~ Book Notes ~ Class Notes o Atomic orbitals combine and form delocalized bands of electrons over the solid. o Model for bonding in solids. o Applies to metallic and covalent solids. o Conduction band: unoccupied orbitals that hold mobile electrons. o Valence band: occupied molecular orbitals.  When electrons move from valence band to conductive band, they become mobile. o Band gap: present in semiconductors and insulators between the conduction and valence bands, causes less conductivity.  Metals do not have a band gap. - Doping: Controlling the Conductivity of Semiconductors o Cause holes in valence band which leads to more electrons in conduction band, or simply can add electrons to conductive band, allowing for more conductivity.  N-type semiconductor: an element with additional electrons is added to another which leads to more electrons in the conductive band since valence band is already full.  P-type semiconductor: causes empty molecular orbitals (holes) in the valence band so electrons can fit through and go to conductive band. Each hole acts as a positive charge.  P-n junctions: tiny spots that are p-type on one side and n- type on the other.  Act as diodes(allow flow of electrons in only one direction) or amplifiers (making a small electrical current into larger one). Bonding Metals: the Electron Sea Model - Metallic bonding: bonding between two metals. o Metals lose electrons easily, so they each donate at least 1 electron to an “electron sea” when bonding.  Sodium becomes a cation which is then attracted to the “sea” of electrons, which holds the solid together.  Mobile electrons allow for conductivity of electricity and heat because electrons are free to move around and disperse thermal energy.  Malleability: the ability to be pounded into sheets.  Ductility: ability to be drawn into wires.  Since there are no localized or “specific” bonds, the electrons accommodate deformation by moving relatively easily. The General Properties and Natural Distribution of Metals - Opaque - Good conductors - High malleability and ductility - Noble metals: metals with low reactivity. - Minerals: homogenous, natural, inorganic crystalline solids. Chapter 11 Notes ~ Book Notes ~ Class Notes o Metals occur in natural oxidation states with minerals. o Ore: rock that contains a high concentration of a specific mineral.


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