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Chemistry 2
https://www.studystack.com/flashcard-4505540
| Question | Answer |
|---|---|
| look at the link in the description | |
| Energy, wavelength, and frecuency | 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). |
| 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 |
| 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 |
| 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 |
| wavelength and E have a _____ realtionship | inverse |
| frecuency and E have a __ relationship | direct |
| wavelngth and frecuency have a __ relationship | inverse |
| 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 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 |
| Hund's Rule | electrons prefer their own orbital fill an orbital with +1/2 (upward arrow) before filling it with -1/2 (downward arrow) |
| 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). |
| 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). |
| 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 |
| 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 |
| 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 |
| Zeff equation | atomic number - non-valence electrons (the inner shells - number of e- between the nuclus and the e- in question) Zeff is unitless |
| 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 |
| 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 |
| 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 |
| 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 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 pair of dots is 1 # of domains 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 |
| when we get an odd number of electrons ... | the most electronegative atom will get the octet |
| delocalized | if a resonating pair is said to be delocalized it doesn't have a permanent location |
| resonance reminder!! if adjoining atoms are all the same element, & there is only one double bond needed, then it is resonance if the adjoining atoms are different elements, the most electronegative atom gets the double (also factor in formal charge) | |
| internuclear distance | distance between atoms |
| In a COVALENT bond, the bond length is influenced by both the size of the atom's core and the bond order (single double, triple) | the size of atom's core & bond order (single double, triple) bonds w/ higher order are shorter (stronger bonds pull atoms closer) In ressonance, it's an average length (if one version has a single and the other a double, reality is in the middle) |
| how to find bond length of a bond in a resonance structure when given the length of the bonds | add up the bond lengths (should be the same amount of measuremnts as the number of structures) then divide by # of structures BOOM |
| why is every covalent bond hybridized | hybridization maximizes bonding cites for atoms to share eletrons (creates more bonding cites) without hybridization, most atoms would not be able to bond. |
| hybridization octet exceptions | Be - 4 e- B - 6 e- C and Si - 8 e- (sp3 is most common, WHEN IN DOUBT) P - 8 e- S - 12 e- |
| isomers | same structure, different formula |
| sigma and pi bonds | single bonds = 1 sigma double bonds = 1 sigma + 1 pi triple bonds = 1 sigma + 2 pi |
| VSEPR | uses Coulombic repulsion between e-s to predict the arrangement of e- pairs around a central atom atoms orientate as far as possible & lone pairs even further than bonding pairs atoms will repell each other & e- pairs will especially repel each other |
| VSEPR effect | lone pairs tend to compress the angles between bonding pairs |
| what determines the polarity | electronegativity and geometry for example, a linear molecule of 2 H is non polar becuase the shape is BOTH symetrical and both Hs have the same electronegativity BUT if one is H and the other F, the it is polar and F is more negitive |
| how to determine polarity based on the difference of electronegativities of the bonding atoms | 2.5 = ionic 1.7 = 50% ionic bond (strong polar) 1.0 = slightly polar 0.5 = nonpolar |
| guarenteee polar molecules | bent distorted tetrahedron triangular pyramids T shaped Square pyramid |
| How to find the second ionization energy of atoms E2 = ?? | |
| make flashcard for how to do number 71b on pg 159 as well as the steps for other calculations too ALSO watch hybrid videos | |
| how to write free responce questions | 1. explain similarities between 2 mcs, atoms, etc. (# of P+, # of e-, # of shells, size, electroneg., ionization E, etc.) 2. explain differences 3. explain why the differences cause whever they are asking about -give specific #s and configurations |
| memorize molecular geometry from the notes | |
| what type of orbitals form sigma and pi bonds | In a multiple bond, sigma bond can be made using hybrid orbitals. π bonds, cannot be made using hybrid orbitals —they only come from unhybridized p orbitals. |
| why do hybrid orbitals only form sigma bonds and why can only p orbitals form pi bonds | -hybrid orbitals overlap end-to-end (or vertically) (perfect for sigma bonds), while unhybridized p orbitals can overlap side-to-side (think paraell lines), which is necessary to form pi bonds. |
| when we talk about pi and sigma bonds, we talk about overlapping orbitals. Don't electrons form bonds not orbitals | e- only bond inside orbitals Two electrons from different atoms can form a bond only if their orbitals overlap. |
| can only p orbitals form pi bonds | Yes — only unhybridized p orbitals can form pi bonds. Pi bonds require side-by-side (parallel) overlap and only p orbitals keep the correct shape for that. |
| Can p orbitals and other original orbitals form sigma bonds? | Yes, absolutely. ANY orbital that can overlap end-to-end can form a sigma bond. |
| do all isotopes have the same properties | Chemical properties: ✔️ Almost identical between isotopes Physical and nuclear properties: These can be very different: heavier atoms behave differently than lighter ones |
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Blessed Rectangle 10