Title | Cheat sheet print outs exam 3 |
---|---|
Author | Courtney Soliz |
Course | General Inorganic Chemistry II |
Institution | University of Houston-Victoria |
Pages | 2 |
File Size | 331.3 KB |
File Type | |
Total Downloads | 75 |
Total Views | 128 |
Download Cheat sheet print outs exam 3 PDF
Name Heat capacity Heat capacity: hot & cold Molar heat Energy change of system Conservation of energy Enthalpy: At Constant V Enthalpy: At Constant P Energy in state changes Overall Energy
Formula
Formal Charge Exs:
q = m c ΔT (heat required=mass*specific heat* change in temperature)
-q ( ) = q ( ) q = n cm ΔT (Heat required = moles × molar heat capacity × ΔT)
ΔE = Efinal – Einitial ΔE = q + w (change in energy = heat +work) ΔE = qv Work requires motion against an opposing force, no motion, w = 0 ΔE = qp + watm = ΔH + watm
q = n * ΔHfusion OR q = n * ΔHvaporization qoverall = qstep 1 + qstep 2 + qstep 3 + qstep 4 + qstep 5
Standard Molar Enthalpy of Formation
ΔH = {(nproducts)+(ΔHfproducts)} – {(nreactants)+(ΔHfreactants)}
Formal charge of each atom
(# of valence e-) – (e- on the atom)
Heat capacity: H2O(l) = 4.184 J/gC (Water) H2O(s) = 2.06 J/gC (Ice) H2O(g) = 1.996 J/gC (Steam) ΔHfusion = 6.02 kJ/mol ΔHvaporization = 40.7 kJ/mol Pay attention to states (s, l, g)!
H: 1eF: 6e-(long pair) + 1e- = 7 e-
Formal Charge Cont.: If there is choice between Lewis structures: • Smaller formal charges are favored • Negative formal charges should be on the most EN atoms • Like charges should not be on adjacent atoms • Note: sum of formal charges = molecular charge • How to calculate formal charge for given formula • 1st find out # electrons of each atom • 2nd find valence electron of each atom
T= Tf - Ti solving for Tf with formula -q=q: remember Tf must be between hot and cold. (+) sign = q, heat added, endothermic, absorbs E, internal E (-) sign = -q, heat lost, exothermic, releases E, internal E 1J=103 kJ so 3 decimal places Enthalpy: ΔHfusion = heat to melt a solid ΔHfusion = − ΔHfreezing
Conversion of energy and changes of state: Endothermic: heat moves from surroundings into system. (melting/ boiling) s l / l g Exothermic: heat moves out of a system into surroundings. (condensing/ freezing) g l / l s Examples: Water boils: H2O(l) H2O(g) endothermic Steam condenses: H2O(g) H2O(l) exothermic
Enthalpy Changes for Chemical Reactions grams
moles
kJ (heat)
During a chemical rxn: -old bonds break: requires E (endothermic) -new bond forms: releases E (exothermic) -sign for E is always (+) for ΔH when we break bonds. -Always label melting/ boiling points on heating curves! -As specific heat (c) the slope of the line (m) Example: Convert ice at - 50°C to water at + 50°C Notice: T is constant when states changing at 0°C & 100°C
Hess’s Law: Multiply a reaction, multiply ΔH ° Reverse a reaction, change the sign of ΔH ° (multiply by -1) Standard Molar Enthalpy of Formation -When writing a formation, you need to have one mole of that substance on the right side of the equation. -If the physical state changes, the formation state changes too.
Quantum Numbers: Heating curve problems: -When changing state (slope), use heat capacity formula. -When constant, use energy in state changes formula. 5 steps total: Step 1: Tice increase Step 2: solid liquid (fusion) Step 3: Twater increase Step 4: liquid gas (vaporization) Step 5 Tsteam increase Qtotal= step 1 + step 2 + step 3 + step 4 + step 5
-Remember that there can be overlaps in quantum numbers. -when determining ml, count how many values m has. When l=0, m=0 -S has 1 orientation, spherical shape -P has 3 orientations, dumbbell shape -D has 5 orientations. Four leaf clover -orientations can be called orbitals! • The number of orbitals in subshells = 2 l + 1 • The total number of orbitals in a shell = n2 -For ms: electron spin remember that one electron is positive and the other is negative.
Atom Electron Configurations:
Isoelectronic: L to R are equal
Bond Polarity and Electronegativity: Nonpolar: equally shared (identical atoms) Polar: unequal sharing (different attraction for e-) Polar means closer to one side and further away from the other side In order to figure out chemical bond need to know polarity.
Electron Configurations of Transition Metals:
Para-magnetism and Unpaired Electrons: Diamagnetic: all magnets cancel, paired eParamagnetic: unpaired e-, weak, sum is very small. Have magnets, but very weak Ferromagnet: permanent, all pointing one direction. Magnet is strong when they add up. Example: Fe (If individual atom-magnets line up in a bulk sample)
Note: ½ filled and filled shells have extra stability Periodic Trends: Atomic Radii: 2 important conclusions: -down column = grow Atoms grow down a group. -across column = shrink Atoms shrink across a period A cation is smaller than its neutral atom. (likelier to be metals) An anion is larger than its neutral atom.
Valence electrons: Remember, # of valence e- = A group #
Covalent bonding: Two atoms make 1 covalent bond Octet rule: used to maintain stability…. Exception: Hydrogen Ions, ionic compounds or covalent compounds used in order to achieve octet rule (8 valence electrons) Bonding pair = actual chemical bond and shared electron pair---- aka shared electron pairs between oxygen and hydrogen example Lone pair = no chemical bond, unshared, both electrons belong to the same atom. Electronegativity: Remember: Fluorine has a stronger electronegativity than Oxygen and Nitrogen which are also strong. Smaller than 1.7- nonpolar covalent bond Between 0.7-1.9= polar covalent bond Greater 1.9- ionic covalent bond Different ΔEN, determine bond polarity
Bond Length and Bond Energy: bigger the atom, longer the distance Triple bond will hold atoms the shortest distance together Single bond is longer than double and longer than triple Bond ABSORBS energy so bond enthalpy is positive More bond formed between atoms= stronger bond...