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General Chemistry 2 Rules for Organic Compounds 1. Carbon always will have 4 bonds 2. Nitrogen will typically only have 3 a. If 4 bonds are present it is cationic such as ammonium 3. Oxygen will typically have 2 bonds 4. Halogens only contain 1 bond Organic Chemistry: study of carbon compounds Mainly deals with hydrocarbons o Others can be present such as H,N,O,F,P,S,Cl,Br and I Organic Molecules are unique 1. Organic molecules will have structural complexity a. Carbon is not energetica4+y fav4-able to form ions i. Does not form C or C ii. Forms 4 covalent bonds (2 sigma and 2 pi bonds) iii. Electronegativity value is 2.5 which is midway on the scale iv. Has the ability to catenate (bond to itself) 1. Only group 14 is able to catenate v. Carbons small size and good orbital overlap means hybridized orbitals 1. Also small size means shorter bonds. This means it has good orbital overlap and strong bonds vi. Carbon to Carbon bonds are short enough that they can Pi bond. This allows for a larger variety of compounds that can be formed using Carbon. 1. Multiple bonds restrict rotation around a bond vii. Carbon has molecular stability (due to its ability to catenate) 2. Organic molecules will have chemical diversity a. There are millions of carbon based compounds that exist. They can exist for 3 reasons i. They can form stable bonds with heteroatoms 1. They will form linear and ring structures 2. They can contain single, double and triple bonds ii. Due to their electron density and reactivity 1. Chemical reactions will occur when an area of high electron density on one molecule meets an area of low electron density on another 2. Pi bonds or C-X bonds will normally have higher reactivity (X=heteroatom) a. When looking at molecular polarity you must consider electronegativity and how the bonds are arranged because some arrangements will cause dipoles to cancel iii. They have functional group importance 1. Functional Groups: Specific combination of atoms, typically C-C multiple bonds or C-X bond, that reacts in a characteristic way no matter what molecule it occurs in a. Most reactions of organic compounds tend to happen in the functional group Recap of why Carbon is unique Ability to catenate Creates stable bonds with heteroatoms Multiple bond formation Can form in linear forms or ring forms Hydrocarbon: an organic compound consisting of only C and H Examples: Methane, ethane, benzene etc. Remember single bonds rotate freely and the molecule can shift in space. Double and triple bonds can not rotate! Alkene: hydrocarbons that contain only single bonds and are referred to as saturated hydrocarbonds Formula for alkene is C n 2n+2n must be a positive integer) Shape: Tetrahedral Hybridization: sp3 Naming Organic Compounds Prefix+ Root + Suffix Root: Is determined by how many carbons are in the longest continuous chain Suffix: indicates the type of organic compound and is found after the root Prefix: Identifies any group that is attached to the main chain. Also tells the position of this group Steps to Naming 1. Name the longest chain of Carbons (root) a. Longest chain of continuous carbon atoms 2. Name the type of compound based on what type of bonds you have a. Single bonds: -ane b. Double bonds: -ene c. Triple Bonds: -yne 3. Name Branches (groups attached to the main chain) a. End the branches in –yl b. Branch names will come after the chain name. When you have 2 or more branches their names must go in alphabetical order. c. Also specify where the branch is located along the chain. Number the main chain Carbon atoms consecutively and try to make your branches have the lowest numbers. General Chemistry 2 Week 2 Geometrical Isomers: cis-trans versions of a molecule Have different physical properties Alkynes: hydrocarbons that contain at least one C-C triple bond. They are known as unsaturated carbons CnH2n-2 Rotation is restricted at the triple bond Linear Shape Sp hybridization e- rich and behaves as a functional group Practice Drawing structures 3,5—Dimethyl—2—octene 2,3,3—trimethyloctane 3—ethyl—3,4 dimethyldecane Trans—2,2—dimethyl—3—hexene When deciding if something is cis or trans you are looking directly before and after the double bond! Aromatic Hydrocarbons: cyclic molecules with delocalized pi electrons They must have alternating double and single bonds to be an aromatic hydrocarbon Skip naming these molecules Resonance forms have alternating single and double bond o Shows delocalized electrons as unbroken or dashed Functional Group There are a lot of functional groups in organic chemistry Functional Groups determine o Physical properties o Chemical properties o Reactivity Functional groups will determine the polarity of a compound o Polarity determines what intermolecular forces are present Intermolecular forces and polarity depend on each other Regions of high and low density will be determines by functional groups Common Abbreviations Used In Organic Chemistry I. R—an organic group of atoms bound by carbon a. Exception—when carbon is bound to oxygen II. R/H—used for substituents a. Can be either R or H III. X—halide a. F,Cl,Br,I **Know the prefixes and suffixes in table 15.5 functional groups. (halogen is the only one with a prefix)* Alcohol Consists of carbon bond to –OH Named by replacing –e at end of hydrocarbon name with –o They have high melting point o This is due to the fact that they form H bond Less acidic and basic compared to H O2 Haloalkanes Contain a carbon bound to a halogen Amines Contains a nitrogen atom It is viewed as a derivative of N3 Weak bases o Due to lone pair on the nitrogen Three types o Primary=NRH 2 o Secondary=NR H 2 o Tertiary= NR 3 o Primary and secondary can form hydrogen bonds but tertiary can not. Carboxyl (not functional group) They are parts of functional groups o Aldehydes, ketones, carboxylic acids, ester and amides Have a C-O double bond o Partial positive charge is located on the carbon and there is a partial negative charge on the oxygen There will always be a dipole on a carboxyl Aldehydes and Ketones Aldehydes have a hydrogen at the terminal end Ketones have two terminal carbons These are very easy to confuse. Be careful! Carboxylic Acid Had the functional group –COOH in it Weak acid Amides Has a nitrogen in it In most drugs, pain meds and LSD Also a peptide bond in biology Skip polymers General Chemistry 2 Week 3 Notes Highlight=Vocabulary Word Highlight=Important information Organic Chemistry (chapter 15) Highlights 1. The uniqueness of Carbon 2. Alkanes, Alkenes, Alkynes and aromatic rings 3. Drawing and naming structures 4. Functional Groups General Chemistry 1 concepts we can apply to organic chemistry 1. Pi and sigma bonds 2. Hybridization 3. Molecular polarity a. This is extremely important!! Chapter 12 Matter occurs as a solid liquid or gas based on the conditions it is in H20(s) H 2 (l) Phase: a physically distinct homologous part of a system o Solids and liquids are considered to be a “condensed” state of matter Intermolecular forces (IF’s) are a form of potential energy. It holds particles together o Intermolecular Force: weak forces that exist between two separate molecules. Higher charge—the stronger the force will be o Intramolecular Force: Bonding forces (ionic, covalent, metallic) Understand that these are different! Recall Coulombs law o Electrostatic potential energy depends on… Charge of molecules Distance of molecules Kinetic energy is responsible for the phase changes o When kinetic energy outweighs potential energy the substance melts or evaporates o When potential energy outweighs kinetic energy substance will condense or freeze Imagine yourself in a flask amongst H2O molecules If you were shrunk down to molecular size and could physically see the H2O molecules you would notice two different types of electrostatic interactions occurring o Intramolecular forces (bonding) Bent and polar H and O atoms being held together through covalent bonds o Intermolecular forces (nonbonding) Separate H2O molecules interacting with each other in a characteristic way H bonds Potential and Kinetic Energy in Phases Potential Energy: energy of attractions (aka intermolecular forces) Kinetic Energy: Random motion of individual particles Gas molecules o High kinetic energy o Low potential energy Liquid molecules o Similar amounts of kinetic and potential energy Solid Molecules o Low kinetic energy o High potential energy Phase changes and enthalpy (delta H) As temperature increases, kinetic energy increases o Endothermic (delta H= +) For endothermic reactions heat is absorbed by the system causing intermolecular forces that hold the molecules together, to be disrupted Particles are moving faster and it is easier to overcome potential energy As temperature decreases, kinetic energy decreases o Exothermic (delta H= -) Particles slow down making attractions (potential energy) become stronger Melting is Delta H fusion Recall Hess’s law Key Equations 1. Heat transferred as a part of a phase change a. Or also known to be within a phase i. q=(moles)(molar heat capacity)(∆Temp) 1. q=mc∆t ii. Recall Change in temp: ∆T=T finainitial 2. During phase change a. q=(moles)(∆H phase change) i. Calculations used to quantify heat change b. Heat removed=q Boiling Point The temperature where the phases of gas and liquid are at equilibrium ±∆H vap Melting Point The temperature where the phases of solid and liquid are at equilibrium ±∆H fusL Liquid-Gas Equilibrium Equilibrium: condition existing when a process and its reverse process proceed at equal rates Represented by a double headed arrow A liquid has a vapor present above it—this causes it to have a vapor pressure Vapor pressure : pressure exerted by the vapor at equilibrium o Rate that molecules turn to vapor and cool/condense happen at equal rates o Recall—the vapor pressure demonstration in class with ether and ester Liquids do not burn, but the vapors above them do This is why ether had a higher flame—it’s vapor pressure was higher Open System: nonequilibrium process o Ex.) H2O in a glass eventually disappears Closed System: Equilibrium process o Constant temperature and a vacuum applied Apply vacuum and remove gas molecules Close vacuum You will notice the gas molecules return o This is because of vapor pressure o Does not have to be boiling to do this Effect of temperature on molecular speed Increase temperature molecules faster (higher kinetic energy) more molecules vaporize o Higher temperatures help overcome the intermolecular forces Increase temperature = higher vapor pressure Vapor Pressure Increase vapor pressure (@ constant temperature) = lower intermolecular forces Boiling Point: Temperature at which vapor pressure equals external pressure o Boiling pressure can change depending on the elevation Higher boiling points=higher intermolecular forces Quantifying the effect of temperature Nonlinear relationship between pressure and temperature can be converted to a linear one through the Clausius equation ln P=−∆ Hvap 1 +C R (T P2 −∆Hvap 1 1 ln = ( − ) P1 R T 1 T 2 For logs and not log (ln) o Log (a/b)= log a- log b o Ln (a/b)= ln a-ln b Inverse log o Log 10^( ) o Ln e^( ) Liquid Gas Equilibrium Solids fixed into a 3D lattice Higher temperatures=higher kinetic energy of molecules meaning that particles can break free from the lattice o Melting Pressure changes have minimal effect o Solids are not easily condensed Because they are already in the most condensed form they can be in No clausius equation Solid Gas Equilibrium Sublimation requires a high vapor pressure Most solids will not sublime Combination of high vapor pressure and low intermolecular forces o It’s not enough to keep particles close enough to be a liquid Non polar molecules with weak intermolecular forces only have van der whals forces Intermolecular Forces 12.3 Intramolecular Force: attractive forces within the same molecule Attraction between cation and anion in ionic bonding Attraction between nuclei and e- pair in covalent Attraction between metal cations and delocalized e- in metallic bonding Intermolecular Force: attractive forces between molecules, atoms and ions Can arise due to partial forces Can also arise between ions and molecules Weak compared to bond forces o This is explained Coulombs Law Types of Forces (in order of strongest to weakest) 1. Bonding (ionic, covalent, metallic) a. This is NOT a intermolecular force!!! All the ones below this are 2. Ion-Dipole 3. Hydrogen 4. Dipole-Dipole 5. Ion-Induced Dipole 6. Dipole- Induced Dipole 7. London Dispersion Van Der Whal (VDW) How close can molecules get to one another? o Minimum distance is the sum of VDW radii o Distance: Bond length < VDW distance o VDW distance: Closest two molecules can get to each other Increases down a group o VDW Radius: ½ of the shortest distance between the nuclei of identical nonbonded atoms Decreases across a period Ion Dipole Forces Strongest intermolecular force Interaction between an ion and a polar molecule Electrostatic motion o Ionic compound dissolved into H2O Dipole-Dipole Attraction between two dipoles of separate molecules Forces more orderly in a solid compared to a liquid Ion Dipole Ions + + o N 3 , Na etc Ion and a molecule with a dipole Dipole-Dipole 2 polar covalent molecules Hydrogen Bond Force that involves an H bound to an F, O or N directly and another very electronegative element o X—H-------Y X and Y are most electronegative X or Y can also be bound to other atoms H bonded compounds will have a higher boiling point and melting point then others o Many H bonds in DNA Ion Induced Dipole and Dipole Induced Dipole A nearby electric field can “induce” or create a distortion in an e- cloud. o This pulls e- density toward a positive pole o Or it pushes it away from a negative pole It’s only a temporary dipole 2 scenarios o Nonpolar Molecules Induces temporary dipole moment o Polar Molecules Enhances (or strengthens) an already existing dipole moment Typically will occur in a solution Strength of the force will depend on the amount of charge present and the polarizability of it Polarizability: The ease at which an e- cloud is distorted (ability to form pole) 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 General Chemistry 2 Exam 1 Study Guide Chapter 15 Organic Chemistry: the study of carbon compounds Mostly deals with hydrocarbons Rule of Thumb for Organic Compounds 1. Carbon will always have 4 bonds 2. Nitrogen typically has 3 bonds a. 4 bonds makes it cationic i. Ex.) Ammonium 3. Oxygen typically contains 2 bonds 4. Halogens only contain 1 bond Uniqueness of Organic Molecules 1. Organic Molecules have structural complexity a. Recall: Carbon 2. Organic Molecules have chemical diversity Carbon It is not energetically favorable for Carbon to form ions Has the ability to catenate Short enough bonds to be able to pi bond o This allows larger variety of compounds that can be formed Form stable bonds with heteroatoms o Form linear and ring structures o Can contain single, double, triple bonds Functional Groups: Specific combination of atoms, typically C-C multiple bonds or C-X bond, that reacts in a characteristic way no matter what molecule it occurs in o Most reactions of organic compounds tend to happen where the functional group occurs Hydrocarbon: an organic compound consisting of only C and H o Ex.) methane, ethane, benzene etc. Alkene: Hydrocarbons that contain only single bonds o Referred to as saturated o Formula is CnH2n+2 o Tetrahedral o Sp3 hybridizes Naming Organic Compounds Prefix + Root + Suffix Root: determined by the amount of carbons are in the longest continuous carbon chain Suffix: indicates the type of organic compound. Is found after the root. Prefix: Identifies any group that is attached to the main chain and the position of that group Steps to Naming 1. Name the longest chain of carbons (root) 2. Name the type of compound based on what type of bonds you have a. Single bonds: -ane b. Double bonds: -ene c. Triple bonds: -yne 3. Name Branches (groups attached to the main chain a. End of branches are named with –yl b. Branch names come after the chain name. i. When there are two or more branches the names must go in alphabetical order c. Specify where the branch is located by numbering your carbon chain to determine where the branch is i. Make sure the carbons are numbered so that the branch has the lowest possible number Geometrical Isomers: cis-trans versions of a molecule Cis and trans will have different properties from each other o When deciding if something is trans or cis you must look at the bond directly before and after the double bond Alkynes: hydrocarbons that contain at least one C-C triple bond Known as unsaturated carbons Formula: CnH2n-2 Rotation is restricted at the triple bond Linear shape Sp hybridized Electron rich and therefore act as a functional group Aromatic Hydrocarbons: cyclic molecules with delocalized pi electrons Single and double must alternate for them to be considered an aromatic hc Resonance forms exist—remember because of alternating single/double bonds Do NOT need to know how to name these Functional Group Many functional groups exist Functional groups determine o Physical properties o Chemical properties o Reacitivity Determine the polarity of a compound o Polarity determines what intermolecular forces are present Intermolecular forces depend on polarity and vice versa Regions of high and low electron density is determined by functional groups Common abbreviations in organic Chemistry You’ll use these when looking at functional groups I. R—an organic group of atoms bound by carbon a. Exception: when carbon is bound to oxygen II. R/H—used for substituents a. Can be either R or H III. X—halide a. F,Cl,Br,I *Know the functional groups and their suffixes!!!!!* Alcohol Carbon bound to –OH Named by replacing –e at the end of the hydrocarbon name with an –o High melting point o Due to the fact that they H bond Less acidic and basic compared to H2O Haloalkanes Carbon bound to a halogen Amines Contains a nitrogen atom Viewed as a derivative of NH3 Weak bases o Due to lone pair on N Three Types o Primary: NRH 2 o Secondary: NR H2 Primary and secondary can hydrogen bond. Tertiary can not. o Tertiary: N3 Carboxyl (not a functional group) They are parts of functional groups o Aldehyde, ketones, carboxylic acids, ester and amides Contain a C=O double bond o Partial positive charge is located on the carbon. o Partial negative charge is on oxygen Always has a dipole Aldehydes and Ketones Aldehydes have a hydrogen at the terminal end Ketones have two terminal carbons Easy to confuse so be careful! Carboxylic Acid Contains the functional group –COOH Weak acid Amides Contains a nitrogen In most drugs Also a peptide in biology General Chemistry 1 concepts to apply to Organic Chemistry 1. Pi and sigma bonds 2. Hybridization 3. Molecular polarity a. Very important Chapter 12 Matter occurs as a solid liquid or gas depending on the conditions it’s in Phase: a physically distinct homologous part of a system o Solids and liquids are considered to be a “condensed” state Intermolecular forces: form of potential energy that holds particles together o Intermolecular Force: weak forces that exist between two separate molecules. High charge—stronger the force will be o Intramolecular Force: bonding forces (ionic, covalent, metallic) Understand that these are different Kinetic Energy is responsible for phase changes o If kinetic outweighs potential substance melts or evaporates o If potential outweighs kinetic substance condenses or freezes Potential and Kinetic Energy in Phases Potential Energy: energy of attraction (intermolecular forces) Kinetic Energy: random motion of individual particles Gas molecules o High Kinetic Energy o Low Potential Energy Liquid Molecules o Similar of kinetic and potential Solid o Low kinetic energy o High potential energy Phase changes and enthalpy (delta H) Temperature increases kinetic energy increases o Endothermic (delta H= +) Heat is absorbed Causes intermolecular forces between molecules to be interrupted Particles moving faster makes it easier to overcome potential energy Temperature decreases potential energy increases o Exothermic (delta H= -) Particles slow down making the intermolecular forces stronger Key equations 1. Within a phase change a. q=(moles)(molar heat capacity)(∆temp) i. q=mc∆t 2. During phase change a. q=(moles)(∆H phase change) b. Used to quantify heat change i. q= heat removed Liquid-Gas Equilibrium Equilibrium: condition existing when a process and its reverse process proceed at equal rates Liquids have a vapor pressure above them—this causes it to have a vapor pressure Vapor pressure : pressure exerted by the vapor at equilibrium o Remember flame demonstration Open System Non equilibrium process o H2O in a glass eventually disappears Closed System: Equilibrium process Increase temperature means an increase in the amount of molecules turning to vapor o Higher temps are helping overcome intermolecular forces o Increased vapor pressure=lower intermolecular forces Quantifying the effect of temperature Nonlinear relationship between pressure and temperature can be converted to a linear relationship −∆Hvap 1 L nP= R (T +C lnP2 =∆Hvap 1 − 1 P1 R (T 1 T 2) *Know how to do problems using these equations! Practice problems can be found on learn and in the book!* Intermolecular Forces 12.3 Intramolecular Forces: attractive forces within the same molecule Attraction between cation and anion in ionic bonding Attraction between nuclei and electron pair in covalent Attraction between metal cations and delocalized electron in metallic bonding Intermolecular Force: attractive forces between molecules atoms and ions Can arise due to partial forces Can also arise between ions and molecules Weak compared to bonding forces Van Der Whal distance: Closest two molecules can get to each other Intermolecular forces Ion Dipole Forces Strongest intermolecular force Interaction between an ion and a polar molecule Electrostatic motion o Ionic compound dissolved in H2O Dipole Dipole Attraction between two dipoles of separate molecules Forces are more orderly in a solid compared to a liquid Ion Dipole Ion and a molecule with a dipole Hydrogen Bond Force involving H bound to F, N, or O directly and another very electronegative element o X—H-----Y X and Y are most electronegative X and Y can also be bound to other atoms H bonded compounds have higher boiling points then others Ion Induced Dipole and Dipole induced Dipole A nearby electric field can “induce” or create a distortion in an electron cloud o Pulls electron density toward positive pole and pushes it from negative pole Only temporary dipole Nonpolar molecules o It will induce them to have a temporary dipole moment Polar Molecules o Enhances (or strengthens) an already existing dipole moment Polarizability: The ease at which an electron cloud is distorted (ability to form pole) Polarizability Trends Smaller atom = harder to polarize o Electrons have less space to move away in a smaller atom Larger atoms = easier to polarize Polarizability increases down a group o Follows same trend as atomic size Polarizability effects ALL intermolecular forces London Forces Temporary and instantaneous reaction o Constantly interacting and then breaking Most universal intermolecular force o Present in all molecules and ions 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 Uniqueness of H2O Liquids: combine the ability to flow with strong intermolecular forces Gases: Random arrangement is same at any place in a container 2 Surface Tension: The energy required to increase the surface area (J/m ) Types of Liquid Molecules 1. Interior Molecules a. Are in the middle of water (not toward the top) i. Therefore, they have attractions on all sides 2. Exterior a. Located at the surface of water i. Only attracted from molecules below and on the sides 1. Causes it to have a downward net force ii. In order to increase its attractions and become more stable like the interior molecules, it must break it’s intermolecular forces (this requires energy)—surface tension iii. Surfactants: decrease the strength of H2O by collecting at the surface and disrupting the H bonds Water Properties Capillarity: Riding of a liquid against the pull of gravity through a narrow space Viscosity: The resistance of a fluid to flow Surface Tension: described above Chapter 12.6 Types of Solids Crystalline Solids: well defined shapes because of their particles (atoms, molecules, ions) occur in an orderly shape Amorphous Solids: non crystalline structures; lacking defined shape due to particles not having an orderly arrangement Unit Cell: Smallest portion of a crystal which gives you the total crystal when repeated in all 3 dimensions Coordination number: number nearest neighbors of a particle o Higher the CN, greater the particles in a given volume Body Centered Cube o 1/8 atoms at 8 corners o 1 atom at center o CN: 8 Face Centered o 1/8 atoms at 8 corners o ½ atoms at 6 faces o CN:12 Packing Efficiency: measure of the total volume occupied by spheres o Simple Cubic: 52% o Body Centered: 68% o Face Centered: 74% Best we can do, most common in nature With unit cells, try to practice visualization. Do not stress over names.
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