Exam 3 cheat sheet PDF

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Spot on cheat sheet to study for practical. ...

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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 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

H: 1eF: 6e-(long pair) + 1e- = 7 e-

q = n * ΔHfusion OR q = n * ΔHvaporization qoverall = qstep 1 + qstep 2 + qstep 3 + qstep 4 + qstep 5

Standard Molar Enthalpy of Formation

ΔH = S{(nproducts)+(ΔHfproducts)} – S{(nreactants)+(ΔHfreactants)}

Formal charge of each atom

(# of valence e-) – (e- on the atom)

Heat capacity: H2O(l) = 4.184 J/g°C (Water) H2O(s) = 2.06 J/g°C (Ice) H2O(g) = 1.996 J/g°C (Steam) ΔHfusion = 6.02 kJ/mol ΔHvaporization = 40.7 kJ/mol Pay attention to states (s, l, g)!

Formal Charge Exs:

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 • 2 nd find valence electron of each atom

DT= 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 DE # (-) sign = -q, heat lost, exothermic, releases E, internal DE ¯ 1J=10 3 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...