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Chemistry 2
| Question | Answer |
|---|---|
| Energy wavelength and frequency flashcards | c = λ v E = h v REMINDER: put units in your calulations |
| what does c mean in the equations mentioned in an earlier slide | c = speed of light constant: 3.00 x 10^8 m/s meters per second |
| what does v mean | frecuency unit: hertz (Hz) or cycles per sec (/s) both units are equal in value |
| what does λ mean | wavelength the unit for the equation is meters however, you might see wavelength measured in nm 1 nm = 10^-9 meters |
| what does E mean | energy joules |
| what does h mean | planck's constant 6.626 × 10⁻³⁴ joule-seconds (J⋅s). |
| define wavelength | how long your wave is how much distance between repitions (he regular, complete patterns of the waveform, with one full repetition defining a wave cycle) the distance between successive crests of a wave |
| frecuency | how many cycles per second |
| converstion between cal and J | 1 cal = 4.184 J |
| conversion between kcal | 1 kcal = 1,000 cal |
| conversion between J and kj | 1 J = 0.001 KJ |
| BOHR atom | -electrons are found in levels (n = 1, 2, 3,...) -Bohr diagram -can jump to higher levels (gain E) to an exctoed state (deltaE = +_ -can jump to lower levels (loses E as light/photons) (deltaE = -) -deltaE = E final - E intial |
| BOHR atom VS BOHR diagram | Bohr atom: concept that e- orbit a nucleus in fixed energy levels or shells, not nessisarily a drawing Bohr diagram: a visual, simplified representation of Bohr atom , typically showing the nucleus and the orbiting electrons in concentric circles |
| How to calculate the energy level of 1 mole of e- in a specific shell of the Bohr atom | divide the negitive Ionization E of 1 mol of the element's atoms in Kcal by the principle quantum number squared (n^2) or megitive ionization E in kj divided by n^2 if a problem wanted it for a single e-, divide by 6.022 x 10^23 |
| why does the equation Eionization/N^2 x -1 only work for atoms like H with only 1 electron | other atoms have electron-electron repulsion and sheilding |
| how to know the number of electrons in an atom | 1. find atomic # (the number of protons = the number of electrons) 2. the charge tells how many e- are lost or gained 3. add or subtract from atomic # based on the charge |
| after finding the energy level of 1 mol of e- in the bohr atom, what does the value of E mean? | bound electrons have negitive energy E = 0 means that the electron is free/has escaped the atom an e-'s E is measured relative to a free e-, the closer to being free, the more energy |
| why dont all of the electrons crowd the same shell | although the e- is alawys a -1 charge and every e- in an atom is atraccted to the same nucleus with the same nuclear charge, they still pread out due to electron-electron repulsion and shielding -nigher nuclear charge = e- closer to nucleus |
| electron-electron repulsion | the e-'s repel each other because they are all negitively charged. This pushes some of the e- further back |
| elctron shielding | inner e-s block some of the nucleus's positive pull from outer electrons |
| how to excite differnt electrons in different orbitals | d e- absorb/are excited by visual light s & p e-s (in lower E shells) need mega E to turn color becuase E gaps between s → p or p → higher orbitals is larger than visible light E -UV or higher to excites it -s & p block are colorless in visible light |
| what happens when an electron becomes excited | |
| When a d electron absorbs visible light: | It jumps to a higher d orbital (excited state) We see the complementary color of the absorbed light → that’s why many transition metal complexes are colored. |
| electrons in which orbitals release certain types of light when they jump to lower shells | d electrons five off visible light s and p electrons give off light energy much higher than visible light |
| electromagnetic spectrum | the spectrum of light |
| wavelength and E have a _____ realtionship | inverse |
| frecuency and E have a __ relationship | direct |
| wavelngth and frecuency have a __ relationship | inverse |
| how differences in absorption or emition of photons in different spectrial regions relates to different types of molecular motion | -molecules move in differnt ways depending on the light being pushed on it Different types of light (photons) have different amounts of energy. Molecules respond differently depending on the energy of the photon: |
| how microwave radiation affects molecular motion | associated w/ transitions in molecular rotaional levels (rotation/spinning) |
| how infrared radiation affects molecular motion | associated w/ transitions in molecular vibrational levels (shaking) -leads to an increase in temperature |
| how ultraviolet/visable radiation affects molecular motion | UV/Visible radiation is associated w/ transitions in electronic E levels (electrons start jumping between levels) |
| coulomb's law | explains force between 2 charged particles (p+ and e-) 1. higher charge = more attraction (electrons would have a higher attraction to 3 protons than 1) 2. closer particles = stronger attraction stronger attraction = higher Ionization E |
| def photon | a particle of light |
| what is a ground state electron configuration | the “normal” arrangement of electrons in an atom when it is at its lowest possible energy alos called resting |
| def Excited electron configuration | One or more electrons have absorbed energy and moved to higher orbitals than they normally occupy. mostly electrons jump from s --> d but in rare occasions e- can jump from p --> d or s --> f, etc. |
| exceptions from the ground state electron configuration (these have excited electron configurations) | Cr - 1s2 2s2 2p6 3s2 3p6 !(4s1)! 3d5 Cu - 1s2 2s2 2p6 3s2 3p6 !(4s1)! 3d10 Ag - 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 !(5s1)! 4d10 Au - 1s1 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 !(6s1)! 4f14 5d10 e- has jumped from an s orbital to a d |
| why do some elctrons jump to higher orbitals (excited electron configurations) ? | -it provides lower overall energy e- might move from a full or half-full S orbital to a d orbital to get a more stable configuration (the s --> d transition is more in transition metals) a full s orbital and a d4 is less stable than a 1/2-filled d5 |
| why dont all electrons jump to higher orbitals and avoid being paired | lower elergy e-s are stuck in lower energy orbitals and it would cost them more E to move them |
| excited state def | same number of electrons, but one or more are at a higher E the amount of energy it absorbs is equal to the gap between its original orbital and the one it jumps to EX: 1s2 2p1 (3 resting electrons) 4s2 4d3 (5 excited electrons) |
| Hund's Rule | electrons prefer their own orbital fill an orbital with +1/2 (upward arrow) before filling it with -1/2 (downward arrow) |
| impossible electron configurations 1 -- 1s2 2s2 2p1 4s3 2 --- 2d10 | 1 -- s can only have 1-2 electrons 2 -- there isn't a 2d subshell |
| low energy happiness chart | 1/2 subshell is decent full subshell is still good full shell is ideal |
| shapes of the orbitals | s - sphereical p - dumbell d - cloverleaf shapes (4 lobes) f - multi-lobed patterns (6–8 lobes). |
| shorthand electron configurations | backtrack to the last Nobel gas write what is left over (start at the nobel gas and write the configuration until you get the element EX: [Ne} 3s2 Mg |
| how colors work | -white things reflect every color of light (reflects IR & absorbs UV) -black things absorb every color (they are warm becuase they absorb visable, infradred , and UV ight) -something only relfects the color of light that it is and absorbs the rest |
| outer electron configuration | -AKA valence (bonding) electrons 1. write shorthand 2. get rid of the Nobel gas and d or f electrons (only keep the outer shell electrons) |
| how to write the electron configuration for Ions -the element's configuration may match another element, but IT DOES NOT BECOME ANOTHER ELEMENT | 1. + e- for negitive charges 2. subtract e- for + charges 3. usually a nobel gas configuration (atoms accept or give away e- to gain an octet) trantitons: take from s e- 1st, then the d e-s (take from outer shell 1st and + to the outer shell 1st) |
| photoelectron spectroscopy def | a technique that is used to gather information about the elctrons in an atom technique used to study the E of e-s in atoms & MCs by shining high-E photons on a sample & measuring kinetic E of the emitted e-'s (binding E, how strongly they were held). |
| how to use photoelectron spectroscopy | the further to the left the peak is, the more E it takes to remove the e-s from the subshell -higher the peak, the more electrons in the subshell -imagine the nucleus is on the far left of the chart -use coulumbs law to explain why some take more E |
| photoelectron spectroscopy Jones defintion | you take an atom, zap it with light of different energies, and you see what energies the e-'s fly off at |
| binding energy VS ionization energy | binding: energy required to remove that electron from its orbital ionization - a specific case of binding E: minimum energy required to remove an outermost (valence) electron from a neutral atom in the gas phase. |
| quantum numbers | Principle quantum number sublevel orbital spin quantum number |
| pauli exclusion principle | no two electrons can have the same four quantum numbers |
| principle quantum numbers | n = 1,2 ,3 4, etc. the shell or energy level |
| sublevel quantum number | -subshells (l) s = 0 p = 1 d = 2 f - 3 |
| orbital quantum number | m or ml draw out orbitals (orbital notation) put zero under the center orbital and fill in the rest like a number line start filling out the orbitals as normal and choose the one you end up on, the number beneath it is the answer |
| spin quantum number | s or ms use hund's rules (all arrows up then down) an up arrow is +1/2 a down arrow is -1/2 |
| how many orbitals does each subshell have | s = 1 p = 3 d = 5 f = 7 |
| paramagnetism | contains unpaired electrons |
| contains unpaired electrons | paramagnetism |
| diamagnetism | contains only paired electrons all of group 2, 12, and 18 are diamagnetic |
| contains only paired electrons | diamagnetism |
| same electron configurations (ions!) | isoelectronic |
| isoelectronic | same electron configurations (ions!) same number of electrons (a nutral atom changes its configuration when it gains or loses electrons) |
| effective nuclear charge (Zeff) | z = atomic number net positive charge felt by an e- in a multielectron atom how much pull from teh nucleus an outer electron acctually feels |
| why is the term "effective" used when describing efffective nuclear charge? | the shielding effect of negatively charged e- prevents higher orbital e- from expieriencing the full nuclear charge by repelling affect of inner-layer electrons |
| Zeff equation | atomic number - non-valence electrons (the inner shells - number of e- between the nuclus and the e- in question) Zeff is unitless |
| atomic radius | not definite - probability of size Fr is the largest atomic radii is expressed in nm (10-9 m) |
| why does size increase as you go down a group | shells get further from the nucleus and more electron-electron repulsion |
| why does size decrease as you go right | as the Zeff increases, the stronger sttraction pulls the outer electrons in closer, shrinking the size. Shell stays the same |
| ion radius (how losing electrons affects the size) | shrink in size (losing e- reverts it back to lower level and an excess of p+ pulls outer electrons closer) -same protons pulling less electrons -Zeff |
| ion radius (how gaining electrons affects the size) | increase in size, extra electrons in outer level repel one another, spreading them out |
| ionization energy | E needed to remove an e' relative magnitude can be estimated through coulomb's law also, it is hard to break into filled levels (full shell, full subshell, half-full subshell) |
| why is the I.E. of He so high | it only has 2 e-, no electron, electron repulsion or shielding the e- are extreamely close to the nucleus because it only has one shell it has a full shell (1s2) & breaking a full shell requires a lot of energy |
| deltaE1 meaning deltaE2 meaning | 1. energy needed to romove the 1st e- 2. energy needed to remove the 2nd electron |
| electronegativity | ability to attract e- pair forming a covalent bond F is the most electronegative F = 4.0 (exclude nobel gases, they are mostly satisified with an octet) |
| why does F have a great affinity for e- | needs only one more elctron to get an octet small atomic radius and high efective nuclear charge (9 protons) The lightest and smallest mean that val. e-s are very close to the nucleus -little electron-electron repulsion of shielding |
| electron affinity | the energy change associated with addition of e- to a gaseous atom |
| how to tell if a single atom is paramagnetic or diamagnetistic | 1. write outer-orbital diagrams for the element 2. observe whether there are paired or unpaired electrons |
| how to form lewis dot structures from orbital diagrams | 1. draw the orbital disgrams for central atom (outer shell only) 2. draw a box to show sharing (cov. bond) (box orbitals with only 1 e-) 3. bonds (the boxes) are represented by ----- 4. unshared e- are represented by dots (orbitals with 2 e- already) |
| adjoining atoms | atoms that are chemically bonded to each other (sharing or tansfering of electrons) |
| how to identify the central atom in a chemical compound? | -it is NVER H and rarely F -choose the one with the lowest electronegativity (central atom must be able to share not hog all the electrons) If there’s a tie, choose the one that can form the most bonds. |
| what to do when you have an odd number of e- | the most electronegative atom will get the extra e-/ will get the octet |
| when you are drawing a lewis structure for an ion.. | put the structure in brackets and write the charge as a superscript |
| deficient octet | an atom in a MC that has fewer then 8 val e- in it's valence shell |
| when there are multiple structures for one molecule (ressonance) | formal charge decides best structure (has closest correlation between # of valence e- used by each atom) -charge an atom would have if all bonding e-s were shared equally , ignoring differences in electronegativity. Lower formal charges → more stable |
| formal charge equation | contributed - aquired contributed = val e- aquired = nonbonding e- + 1/2 bonding e- -do this for each atom -when finding aquired, each bond (----) counts as 1, so double bonds are 2 etc |
| coulombs law formula | F = k * (q1q2) / r²* (the r is squared) F = force k = constant q = charge of one particle r = distance |
| how to determine the hybridization of the central atom in a covalent compound | # of domains (count each electron pair around the central atom both --- and dots) a line is one and each single dot is 1 # of electron pairs determines this 2 = sp 3 = sp2 4 = sp3 5 = sp3d 6 = sp3d2 triple & double bonds count as 1 domain |
| how a molecule is nonpolar | dipoles cancel when they are equal in magnitude and opposite in direction. how to know the megnitude of an --> : The bigger the difference between the bonded atoms, the more polar the bond is, and the stronger the dipole arrow. |
| how to determine if a tetrahedral is polar or nonpolar | Tetrahedral + identical substituents → nonpolar. (CH4) Tetrahedral + mixed substituents → polar. (CH2Br2) |
| can polyatomic ions have a dipole | yes, as long as it's lewis structure is asymetrical (still mention that it's an ion) |
| can ionic compounds have a dipole | not really |
| why does an unshared pair of electrons has such a strong repulsion with adjoining atoms in a molecule | unshared pair of e- exert greater repulsion than e- pairs in bonds |
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Blessed Rectangle 9