Gen Chem 2 Week 4 Notes
Gen Chem 2 Week 4 Notes chem 10061-001
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This 5 page Class Notes was uploaded by Jessica Brown on Monday February 15, 2016. The Class Notes belongs to chem 10061-001 at Kent State University taught by David bowers in Summer 2015. Since its upload, it has received 28 views. For similar materials see general chemistry 2 in Chemistry at Kent State University.
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Date Created: 02/15/16
General Chemistry 2 Week 4 Notes Polarizability Trends Smaller the atom is the harder it is to polarize o This is due to the fact that the electrons have less space to move away Larger atoms are easier to polarize o This is due to the fact that electrons have more space to get away from each other Polarizability increases down a group because atomic size also increases down a group o Decreases across a period Polarizability effects all intermolecular forces! London (dispersion) Forces Temporary and instantaneous interaction o They are constantly interacting and then breaking Most universal intermolecular force o They’re present in all molecules and ions Polarizability is important for London forces Cylindrical chain molecules have more London forces then compact molecules o More points of contact Very weak Large molecule=large amount of London forces Deciding Stronger Dispersion Forces 1. Polarizability 2. Surface area a. Larger=more points of contact for forces to occur i. More points of contact/forces = higher boiling points Phases and Uniqueness of H2O Liquids: combine the ability to flow with strong intermolecular forces Least understood of all the phases Gases: random arrangement is same at any place in container 2 Surface Tension: The energy requires to increase the surface area (J/m ) Two Types of Liquid Molecules 1. Interior Molecules a. Are in the middle of the water (not toward the top) i. Therefore, they are attracted by other molecules intermolecular forces on all sides of them 2. Exterior Molecules a. Located at the surface of water i. Only attracted from molecules below and on the sides 1. Due to this it experiences a downward net force b. In order to increase its attractions and become more stable (like interior molecules) it must break its intermolecular forces ( which requires energy)—surface tension i. The stronger the intermolecular forces the more energy that is required to increase surface area (greater surface tension) ii. Surfactants: decrease the strength of H2O by collecting at the surface and disrupting the hydrogen bonds 1. Ex.) Soap, petroleum recovering agents c. Stronger intermolecular forces = Higher surface tension Macroscopic properties of Liquid Capillarity: Rising of a liquid against the pull of gravity through a narrow space (capillary action) o Results from the competition of intermolecular forces within a liquid (cohesive forces) and those between the liquid and the tube wall (adhesive forces) H2O in a glass vs. Hg in a glass H2O is a concave meniscus o H bonding to walls (SiO2) Hg convex meniscus o Cohesive forces > adhesive forces Viscosity: The resistance of a fluid to flow o Result of intermolecular forces that impede the movement of molecules around and past each other Cause friction and stop/slow the flow o Liquids and gases flow but viscosity of liquid is much greater o Factors that affect viscosity Temperature Increased temperature increased movement breaking of intermolecular forces Molecular Shape o Small spherical molecules have a lower viscosity o Long chain molecules have a higher viscosity Types of Solids Chapter 12.6 Crystalline Solids: Well defined shapes because of their particles (atoms, molecules, ions) occur in an orderly shape Is composed of particles packed in an orderly 3D array—Crystal lattice o Lattice consists of all points with identical surroundings Crystals are a result of time—they have time to get to where they want to be before solidifying Amorphous Solids: non crystalline structures; they have a lack of defined shape due to the particles not having an orderly arrangement Think of boiling a dissolved substance they get disorganized and you’re left with a blob of a solid Unit Cell: Smallest portion of a crystal which gives you the total crystal when repeated in all 3 dimensions Coordination Number: number of nearest neighbors of a particle The higher the Coordination number, the greater number of particles in a given volume Counting atoms in a unit cell A view of one atom of a simple cell In simple cubic: only one total atom (in entire cube) Remember orange example Different Unit Cells Simple Cubic The centers of 8 identical particles define the corners of a cube o Particles touch along the cube edges o Coordination number: 6 Body Centered Identical particles lie at each corner and in the center of each face but not in the center of the cube o Particles at the corner do not touch each other, they only touch the center atom 1/8 atom at 8 corners 1 atom at center o Coordination number: 8 Face Centered Identical particles lie at each corner and in the center of each face but not in the center of the cube o Particles at the corner touch the center of each face but they do not touch the other corner atoms 1/8 atoms at 8 corners ½ atoms at 6 faces o Coordination number: 12 Packing Efficiency: measure of the total volume occupied by spheres Think of how you would stack oranges in 2D to occupy the most space available o Simple Cubic Packing: 36 oranges o Close Packing: 42 oranges Most prevalent in nature Simple Cubic: 52% efficiency Body Centered: 68% efficiency Face Centered: 74% efficiency o Best we can do most common in nature Types and Properties of Crystalline Solids 1. Atomic: solids consist of individual atoms a. These are held together mainly by dispersion forces b. Soft c. Low melting point d. Poor electric and thermal conductor e. Cubic closest packing for Ar (s) f. Atoms separated by VDW distance i. All noble gases 2. Molecular: solids consist of individual molecules a. These are commonly held together by various intermolecular forces i. H bonds, dipole-dipole or dispersion is most common b. Fairly soft c. Low to moderate melting point d. Poor electrical and thermal conductor e. Molecules separated by VDW distance 3. Ionic: solids consist of a regular array of cations(smaller) and anions(larger) a. No intermolecular forces because ionic is bonding b. Smaller of two ions lies in space (holes) formed by packing of the larger ions c. Unit Cell: has the same cation:anion ratio as empirical formula d. Very strong e. Brittle f. High melting and boiling points g. Electrical conductivity i. Solids: electrons are held tightly by anions which causes it to have a low conductivity ii. Molten Aqueous: separated bonds cause them to have a high conductivity due to mobile charges 4. Metallic: solids that exhibit organized crystal structure a. No intermolecular forces because metallic is bonding 5. Network Covalent: solid consists of atoms covalently bonded a. No intermolecular forces because covalent is bonding b. Diamonds, Graphite etc. fall into this category Some Examples Ionic Solid Sodium Chloride—face centered cube o Recall periodicity: Group 1 halides + - o A Na face centered cube and a Cl face centered cube combined Na in the holes of Cl o Coordination number of (Na)= Coordination number of (Cl)= 6 Each surrounded by 6 of the other o Unit cell contains 4 Na+ ions and 4 Cl- ions 1:1 ratio NaCl Zinc Blende Structure (ZnS) Two faced centered cubic arrays Combine so that each atom is touching 4 ions of the other Fluorite Structure Common 1:2 cation anion ratios Face centered array of Ca+ Packing of polyatomic ions causes them to be softer than others because the packing isn’t as efficient Metals and Alloys Softhard Low very high melting points Excellent conductivity Malleable and ductile o Think of copper wire Cubic or close packed structures Network Covalent Atoms are linked via covalent bonds Very high melting and boiling points Silicates (SiO2) o Graphites, Diamond These are the strongest
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