AP Chem Units 11 and 12 - steps to identify intermolecular forces like dipole dipole, london dispersion PDF

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steps to identify intermolecular forces like dipole dipole, london dispersion and heat of fusion, heat of vaporization and vapor pressure notes ...

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AP Chem Units 11 & 12 INTRAmolecular forces (forces within a molecule) are STRONGER than INTERmolecular forces (forces that connect molecules. Intermolecular Molecules: 

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The attraction between molecules are not nearly as strong as the intramolecular bond attraction that holds compounds together. They are however strong enough to control physical properties like such as boiling and melting points, vapor pressures and viscosity (ability to flow through substances).

Van der Waal’s Forces (intermolecular) 

Dipole-Dipole interactions (not as strong): molecules that have permanent dipoles are attracted to each other.  The positive end of one is attracted to the negative end of the other and vice versa  These forces are only important when molecules are close to each other (only in solid and liquid phases).  The more POLAR (has lone pair) the molecule is, the higher the boiling point is, because there are stronger intermolecular forces which take more energy to break.

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London Dispersion Forces: (weakest): **everything on Earth have them.  These forces are present in all molecules whether they are polar or nonpolar. Usually nonpolar have only London Dispersion or Elements.  Polarizability: the tendency of an electron cloud to distort in this way  The shape of the molecule affects the strength of dispersion forces: long, skinny molecules like n-pentane tend to have stronger dispersion forces (higher boiling points) than short fat ones like neo-pentane.  Linear molecules like hexane have stronger dispersion, higher boiling point, due to increased surface area. The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds which are easier to polarize. Hydrogen Bonding: (Strongest): the dipole- dipole interactions experienced when H is bonded to N, O, and F.

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**With any intermolecular force, when putting it in order and stuck in between 2 of the same forces look at which has the greater mass to see which has the stronger force/higher boiling point.  Ionic and covalent, the strongest. Ion-Dipole interactions: are important in solutions of ions. Only happens when salt dissolved in polar solvents (water). The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

The stronger the intermolecular forces, the higher the boiling point (must mention this when explaining with order of highest to lowest boiling points). Viscosity: Resistance of a liquid to flow. The greater the resistance, the greater the viscosity. Ex: honey, syrup.  It’s related to the ease with which molecules can move past each other.  Viscosity increases with stronger intermolecular forces and decrease with higher temperatures.  Increases with increasing molecular weight. Surface Tension: results from the net inward force experienced by the molecules on the surface of the liquid. a) As temperature increases, viscosity and surface tension decrease because molecules move with more kinetic energy, thus intermolecular forces will decrease. b) As intermolecular forces of attraction become stronger, viscosity and surface tension increase. Gas Liquid = condensation, LiquidGas= vaporization, GasSolid= deposition, SolidGas= sublimation. Heat of Fusion (Hf): is the energy required to change a solid at its melting point to a liquid. 

q=Hf x m. Hf should be given, m=mass. Use this equation when melting or freezing, if freezing Hf would be opposite sign (so negative)

Heat of Vaporization: the energy required to change a liquid at its boiling point to a gas. 

q=Hv x m. use this equation when vaporizing or condensing, if condensing, then Hv opposite sign.

q=mc∆T- use this equation when NO PHASE CHANGE occurs. ** Remember to convert from KJ to J.    

Slanted lines are Phases (Solid, liquid, gas) On Slanted lines Kinetic energy increases, and potential energy decreases. On straight lines, potential energy increases, kinetic decreases Straight lines, phase changes occur (melting, vaporization) *** The temperature of a substance DOES NOT rise during a phase change!!

Vapor Pressure: at any temperature, some molecules in a liquid have enough energy to escape. More molecules escape at higher temperatures.   

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As temperature rises, the fraction of molecules that have enough energy to escape increases. As more molecules escape the liquid, the pressure they exert increases. As temp go up, vapor pressure go up. The boiling point of a liquid is the temperature at which its vapor pressure equals its atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure 760 torr or 1 atm. The higher the boiling point, the stronger the intermolecular forces. Ethyl glycol will have the highest boiling point and strongest intermolecular force.

Solids: can fall into 2 groups  

Crystalline- particles are in highly ordered arrangement (hard to break apart). Amorphous- no particular order in the arrangement of particles (easy to break).

Metallic Bonding: metals are not covalently bonded but the attractions between atoms are too strong to be Van der Waals forces. In metals valance electrons are delocalized (able to move) throughout the solid.    

Because the electrons in metals are delocalized, none of them belong to any metal cations. High electrical conductivity in metals results from loosely held, mobile electrons. Electrical current easily moves through the electron sea. High thermal conductivity results because the mobile electrons can readily transfer kinetic energy (heat) through the solid. Alloys have properties of pure metal, but contain 2 or more elements!

Substitutional Alloy: a solid solution where atoms of another element replace some of the metal atoms. 

Ex: Gold alloys with silver or copper. Silver and copper are slightly smaller than gold, therefore can readily replace them. They improve the strength and shine of gold when replacing the gold atoms by altering the crystal structures.

Interstitial Alloy: another atom occupies holes between the metal atoms. Typically an interstitial element is non-metal with a much smaller bonding radius than the metal.  

The interstitial atoms for covalent bonds with the metal, making the metal lattice harder, stronger, and less malleable/ductile. Ex: steel, an interstitial alloy of iron and carbon.

Ionic Solids: Compounds of metal cations and nonmetal anions   

NaCl2, Al2O3, K2SO4, FeS Crystal lattice of ions locked in place by relatively strong ionic bonds (electrostatic force). Brittle, high melting nonconductors (insulators) in pure form, conductors in solution (aq).

Metallic Solids: metal ATOMS ONLY! (Al, Fe, Cr, Ni).  

Metal ions surrounded by uniform sea of delocalized (mobile) valance electrons Good conductors of heat and electricity, malleable, ductile, mixtures of metals from alloys.

Covalent-Network Solids (STRONGEST): carbon (diamond), metalloids (Si, Ge), and compounds of metalloids (SiO2, SiC, BN).  

Held together by chains or networks of covalent bonds. Every atom is covalently bonded to other atoms. No intermolecular forces, no molecules. Hard, brittle (not always), high melting point, poor conductor (nonmetals) or semi-conductor (metalloid).

Molecular Solids (weakest): Compounds of only Nonmetals: H2O (s), P2O5, C6H12O6.   

Individually and covalently bonded molecules held together by intermolecular forces. Soft, low melting point, non-conductors Ex: Graphite is a molecular solid in which atoms are held together by Van der Waals forces.

Semi-conductors- usually metalloids!! •

Doping: The process of adding small amounts of impurities to the crystal lattice of a material to influence its electrical conductivity

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N-type semiconductor (negative charge): can carry a negative charge, by replacing a few atoms of pure silicon, each which has 4 valence electrons, with atoms of phosphorus, which contain five valence electrons. The ‘extra’ valence electrons greatly increase the conductivity of silicon.

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P-type semiconductor: (positive charge) silicon doped with a Group 13 (3A) element having three valence electrons. The resulting material has a deficiency of valence electrons called holes. A hole can be thought of as having a positive charge. Conductivity is increased because electrons can jump from hole to hole as they move through the material.

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N and P type semiconductors are the basis for many electronic devices such as diodes, transistors, and solar cells....