Term
| the one distinguishing characteristic of all organic compounds |
|
Definition
| they all contain the element carbon |
|
|
Term
|
Definition
| the study of carbon compounds |
|
|
Term
| Why is C special? Why do most presently known chemical compounds contain C? |
|
Definition
| because of its structure and its consequent position on the periodic table |
|
|
Term
| why C is able to form so many compounds |
|
Definition
because of its ability to bond
-can share 4 ve's and form 4 strong covalent bonds
-C atoms can bond to each other, forming long chains and rings |
|
|
Term
| some elements commonly found in organic compounds |
|
Definition
|
|
Term
| the color used to represent H |
|
Definition
|
|
Term
| the color used to represent C |
|
Definition
|
|
Term
| the color used to represent N |
|
Definition
|
|
Term
| the color used to represent O |
|
Definition
|
|
Term
| the color used to represent F |
|
Definition
| some kind of pale turquoise color |
|
|
Term
| the color used to represent P |
|
Definition
|
|
Term
| the color used to represent S |
|
Definition
|
|
Term
| the color used to represent Cl |
|
Definition
|
|
Term
| the color used to represent Br |
|
Definition
|
|
Term
| the color used to represent I |
|
Definition
|
|
Term
| the general composition of an atom |
|
Definition
-small, dense nucleus with protons and neutrons
-electrons circling it at a relatively large distance |
|
|
Term
| the electron density in an atom |
|
Definition
| denser closer to the nucleus than towards the outer edge |
|
|
Term
|
Definition
| atoms with the same atomic number but different mass numbers |
|
|
Term
| the behavior of a specific electron in an atom can be described by... |
|
Definition
|
|
Term
| the solution to a wave equation |
|
Definition
|
|
Term
|
Definition
| the solution to a wave equation |
|
|
Term
|
Definition
this is the Greek letter psi
this denotes the orbital, which is the solution to a wave function |
|
|
Term
| how an orbital describes the volume of space around a nucleus that an electron is most likely to occupy |
|
Definition
| by plotting the square of the wave function, Ψ2, in 3D space |
|
|
Term
|
Definition
|
|
Term
| the orbitals we're primarily concerned with in OCHEM and biochem |
|
Definition
| s and p because they're the most common |
|
|
Term
| the shape of an s orbital |
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
4 of the 5 are shaped like a 4-leaf clover
the other one is shaped like a dumbbell inside a donut |
|
|
Term
| the atoms in an orbital are organized into... |
|
Definition
| different electron shells |
|
|
Term
|
Definition
| a group of an atom's electrons with the same principal quantumn number |
|
|
Term
| each orbital in a shell can be occupied by how many electrons? |
|
Definition
|
|
Term
| composition of the first energy level in an atom |
|
Definition
| a 1s orbital, thus just 2 electrons |
|
|
Term
| composition of the second energy level in an atom |
|
Definition
| a 2s orbital and 3 2p orbitals, thus 8 electrons |
|
|
Term
| composition of the third energy level in an atom |
|
Definition
| a 3s orbital, 3 3p orbitals, and 5 3d orbitals, thus 18 electrons |
|
|
Term
|
Definition
|
|
Term
| how the p orbitals are arranged |
|
Definition
| along mutually perpendiculat directions,
denoted px, py, and pz
[image] |
|
|
Term
| the lobes of an orbital are separated by... |
|
Definition
| an area of zero electron density called a node |
|
|
Term
|
Definition
| an area of zero electron density between the nodes of an orbital |
|
|
Term
| how the lobes of the orbital are denoted in the wave function |
|
Definition
| with the algebraic signs + and - |
|
|
Term
| ground-state electron configuration |
|
Definition
| the orbitals occupied by an atom's electrons at its lowest energy arrangement (its ground state) |
|
|
Term
| the 3 rules for listing the orbitals occupied by an atom's electrons (WRITE THIS DOWN!) |
|
Definition
1: Aufbau principle; the lowest energy orbitals fill up first, according to the order 1s --> 2s --> 2p --> 3s --> 3p --> 4s --> 4d. Note that the 4s orbital lies between the 3p and 3d orbitals.
2: Pauli exclusion principle; electrons act in some ways as if they were spinning around an axis, somewhat like how the Earth spins. This spin can have different orientations, denoted as up (/|\) and down (\|/). Only 2 electrons can occupy an orbital, and they must be of opposite spin.
3: Hund's rule; if 2 or more orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half full. |
|
|
Term
|
Definition
| The lowest energy orbitals fill up first, according to the order 1s --> 2s --> 2p --> 3s --> 3p --> 4s --> 4d |
|
|
Term
| Pauli exclusion principle |
|
Definition
| Only 2 electrons can occupy an orbital, and they must be of opposite spin |
|
|
Term
|
Definition
| If 2 or more orbitals of equal energy are available, one electron occupies each with spins parallel until all orbitals are half full |
|
|
Term
| the orbitals that fill up first |
|
Definition
| the lowest energy orbitals |
|
|
Term
| how many electrons can occupy an orbital? |
|
Definition
|
|
Term
| how must the electrons in an orbital spin? |
|
Definition
| they must be of opposite spin |
|
|
Term
| what electrons do when 2 or more orbitals of equal energy are available |
|
Definition
| one electron occupies each with spins parallel until all orbitals are half full |
|
|
Term
| how to predict ground state electron configuration: |
|
Definition
|
|
Term
| the valency of carbon in organic compounds |
|
Definition
| tetravalent; always forms 4 bonds when it joins other elements to form stable compounds |
|
|
Term
|
Definition
| always forms 4 bonds when it joins other elements to form stable compounds |
|
|
Term
| the spatial orientation of a carbon atom's bonds |
|
Definition
| a regular tetrahedron (basically a trigonal pyramid) |
|
|
Term
| how bonds coming out of the page toward the viewer |
|
Definition
[image]
heavy, wedge-shaped line |
|
|
Term
| how bonds on the same plane as the page are represented |
|
Definition
|
|
Term
| how bonds receding back behind the page away from the viewer are represented |
|
Definition
|
|
Term
| why do atoms bond together? |
|
Definition
| because the the compound that results is more stable and lower in energy than the separate atoms |
|
|
Term
| flow of energy when a bond is formed |
|
Definition
| energy, usually as heat, always flows out of a chemical system when a bond is formed |
|
|
Term
| flow of energy when a bond is broken |
|
Definition
| energy must be put into a chemical system to break a bond |
|
|
Term
| why do atoms want an octet? |
|
Definition
| to obtain noble gas configuration |
|
|
Term
|
Definition
| electrostatic attraction between an anion and a cation |
|
|
Term
|
Definition
| sharing of electrons between atoms |
|
|
Term
|
Definition
| neutral collection of atoms held together by covalent bonds |
|
|
Term
|
Definition
|
|
Term
| a stable molecule results when... |
|
Definition
| a noble gas configuration is achieved for all the atoms |
|
|
Term
| Kekulé structures aka line-dot structures |
|
Definition
| kinda like Lewis structures, but the bonds are represented with lines instead of electron pairs |
|
|
Term
|
Definition
valence electrons that are not used for bonding
aka nonbonding erlectrons |
|
|
Term
| 2 models that have been developed to describe covalent bonding |
|
Definition
-valence bond theory
-molecular orbital theory |
|
|
Term
|
Definition
| According to this theory, a covalent bond forms when 2 atoms approach each other closely and a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom. The electrons are now paired in the overlapping orbitals and are attracted to the nuclei of both atoms, thus bonding the atoms together. |
|
|
Term
|
Definition
covalent bond formed by head on overlap of atomic orbitals
this is basically what happens in a single bond |
|
|
Term
| why molecules are more stable than the free atoms |
|
Definition
| because the product molecule has less energy than the starting free atoms |
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
| the distance between 2 nuclei that leads to maximum stability; this is at the minimum energy point between the 2 nuclei |
|
|
Term
|
Definition
| see if you can find it in the video |
|
|
Term
|
Definition
| hybrid of an s orbital and three p orbitals |
|
|
Term
| depiction of unhybridized s and p orbitals |
|
Definition
|
|
Term
| depiction of sp3 hybrid orbitals |
|
Definition
|
|
Term
| why there's directionality in a tetrahedral molecule |
|
Definition
| because the sp3 hybrid orbitals have 2 lobes and are unsymmetrical about the nucleus, giving them directionality and allowing them to form strong covalent bonds when they overlap with orbitals from other atoms |
|
|
Term
| why sp3 hybrid orbitals form stronger bonds than unhybridized s and p orbitals |
|
Definition
| because one of the 2 lobes is larger and can therefore overlap more effectively with an orbital from another atom to form a bond |
|
|
Term
| why sp3 hybrid orbitals are asymmetrical |
|
Definition
| because the 2 lobes of the p orbital have different algebraic signs (+ and -). Thus, when a p orbital hybridizes with an s orbital, the positive p lobe adds to the s orbital but the negative p lobe subtracts from the s orbital, making the orbital strongly oriented in one direction |
|
|
Term
| the bond length between C and H |
|
Definition
|
|
Term
| the bond angle in a tetrahedral molecule |
|
Definition
|
|
Term
| chemical formula for methane |
|
Definition
|
|
Term
| chemical formula for ethane |
|
Definition
|
|
Term
| the simplest molecule with a C-C bond |
|
Definition
|
|
Term
| the overlap between the C atoms in an ethane molecule |
|
Definition
|
|
Term
| the hybrid orbital that comes from each C atom in an ethane molecule |
|
Definition
|
|
Term
| length of the C-C single bond |
|
Definition
|
|
Term
| the bond angles when an atom has 4 electron domains |
|
Definition
|
|
Term
|
Definition
| this occurs when an s orbital hybridizes with the px and py orbitals, but not the pz orbital; it remains unchanged |
|
|
Term
| the orientation of sp2 hybrid orbitals |
|
Definition
| trigonal planar with the 3 lobes 120° apart from each other and the unchanged p orbital perpendicular to the sp2 plane |
|
|
Term
|
Definition
| covalent bond between 2 p orbitals |
|
|
Term
| the regions occupied by the sigma (σ) bond |
|
Definition
| the region centered between nuclei |
|
|
Term
| the regions occupied by the sigma (π) bond |
|
Definition
| the regions above and below a line drawn between the nuclei |
|
|
Term
| length of the C-H single bond |
|
Definition
|
|
Term
| length of the C=C double bond |
|
Definition
|
|
Term
| why a double bond is less than twice as strong as a single bond |
|
Definition
| because the sideways overlap in the pi (π) part is not as much as the head on overlap in thew (σ) sigma part |
|
|
Term
|
Definition
| this happens when an s orbital hybridizes with the px orbital, but not the py and pz orbitals; they remain unchanged |
|
|
Term
|
Definition
| sp-sp bond between the sp orbitals and py-py pi bond between the py orbitals and pz-pz pi bond between the pz orbitals |
|
|
Term
| length of then C-C triple bond |
|
Definition
|
|
Term
|
Definition
| often 109 pm, but sometimes 106 pm |
|
|
Term
| electron domain alignment associated with sp hybrid orbitals |
|
Definition
|
|
Term
| electron domain alignment associated with sp2 hybrid orbitals |
|
Definition
|
|
Term
| electron domain alignment associated with sp3 hybrid orbitals |
|
Definition
|
|
Term
| bond angle associated with sp hybrid orbitals |
|
Definition
|
|
Term
| bond angle associated with sp2 hybrid orbitals |
|
Definition
|
|
Term
| bond angle associated with sp3 hybrid orbitals |
|
Definition
|
|
Term
| why P can form 5 covalent bonds |
|
Definition
| due to its position on the Periodic Table, it can expand its outer shell octet and form more than the typical number of covalent bonds |
|
|
Term
| why S can form 4 covalent bonds |
|
Definition
| due to its position on the Periodic Table, it can expand its outer shell octet and form more than the typical number of covalent bonds |
|
|
Term
|
Definition
| com[pounds that contain a P atom bonded to 4 O, with one of the O bonded to a C |
|
|
Term
|
Definition
| compound with a S atom bonded to one H and one C |
|
|
Term
|
Definition
| compound that has a S atom bonded to 2 C |
|
|
Term
| molecular orbital (MO) theory |
|
Definition
| describes covalent bond formation as arising from a mathematical combination of atomic orbitals (wave functions) on different atoms to form molecular orbitals, so called because they belong to the entire molecule rather than to an individual atom |
|
|
Term
| the 2 ways for a molecular orbital to occur |
|
Definition
-the additive way
-the subtractive way |
|
|
Term
| the additive way for molecular orbitals to occur |
|
Definition
| this combination leads to a molecular orbital that is lower in energy and is roughly egg-shaped |
|
|
Term
| the subtractive way for molecular orbitals to occur |
|
Definition
| this combo leads to a molecular orbital that's higher in energy and has a node between nuclei |
|
|
Term
| bonding molecular orbital (MO) |
|
Definition
| molecular orbital that's lower in energy than the atomic orbitals from which it is formed |
|
|
Term
| antibonding molecular orbital (MO) |
|
Definition
molecular orbital that's higher in energy than the atomic orbitals from which it is formed
can't contribute to bonding because the electrons it contains can't occupy the central region |
|
|
Term
|
Definition
| shorthand method for writing structures in which C-H and C-C single bonds aren't shown, but instead understood
example: propane being written as CH3CH2CH3 |
|
|
Term
|
Definition
| shorthand way of writing structures in which C atoms are assumed to be at each intersection of 2 lines (bonds) and at the end of each line |
|
|
Term
| rules for drawing skeletal structures (WRITE THIS DOWN!) |
|
Definition
1: C atoms aren't usually shown. Instead, a C atom is assumed to be at each intersection of 2 lines (bonds) and at the end of each line. Occasionally, a C atom might be indicated for emphasis or clarity.
2: H atoms bonded to C aren't shown. Because C always has a valence of 4, we mentally supply the correct number of H atoms for each C.
3: Atoms other than C and H are shown. |
|
|
Term
| what the end of the line in a skeletal structure represents |
|
Definition
|
|
Term
| what a 2-way intersection in a skeletal structure represents |
|
Definition
|
|
Term
| what a 3-way intersection in a skeletal structure represents |
|
Definition
|
|
Term
| what a 4-way intersection in a skeletal structure represents |
|
Definition
|
|
Term
| bonds that have a circular cross-section and are formed by head-on interaction |
|
Definition
|
|
Term
| bonds formed by sideways interaction of p orbitals |
|
Definition
|
|
Term
| shape of a hexagonal benzene molecule |
|
Definition
|
|
Term
| hybridization of C+ ion when it forms 3 single bonds |
|
Definition
|
|
Term
| geometry of C+ ion when it forms 3 single bonds |
|
Definition
|
|
Term
| electronic relationship refers to... |
|
Definition
|
|
Term
| singlet (spin-paired) orbital |
|
Definition
| has pair of electrons, opposite spin |
|
|
Term
| triplet (spin-unpaired) orbital |
|
Definition
|
|
Term
| how the chemical symbols in an organic compound's molecular formula are ordered |
|
Definition
| first C, then H, then alphabetical from there |
|
|
Term
| why the 4 bonds of C can be arranged in a variety of ways |
|
Definition
| because of differences in hybridization |
|
|
Term
| bond angle of tetrahedral bonds |
|
Definition
|
|
Term
| hybridization of tetrahedral |
|
Definition
|
|
Term
| hybridization of trigonal planar |
|
Definition
|
|
Term
| bond angle of trigonal planar |
|
Definition
|
|
Term
| where the p orbital is in trigonal planar |
|
Definition
| perpendicular to the sp2 orbitals on the trigonal planar |
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
| some other elements that are often found in organic compounds |
|
Definition
|
|
Term
|
Definition
| small collection of atoms that gives a molecule specific properties |
|
|
Term
| volume of an atom has to do with... |
|
Definition
|
|
Term
| the most common C isotope |
|
Definition
C-12
99% of the C in mature |
|
|
Term
| distance between energy layers vs. distance from nucleus |
|
Definition
| the further you get from the nucleus, the closer and closer those energy layers get |
|
|
Term
| s orbitals always cone ______ at a time |
|
Definition
|
|
Term
| p orbitals always cone ______ at a time |
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
|
|
Term
|
Definition
| point of zero electron density between the nodes of an orbital |
|
|
Term
| the difference between the p orbitals |
|
Definition
| the orientation; they're the same except for the axis |
|
|
Term
|
Definition
| exactly the same in energy |
|
|
Term
| d orbitals always cone ______ at a time |
|
Definition
|
|
Term
| f orbitals always cone ______ at a time |
|
Definition
|
|
Term
| why the 1st energy level can have only 2 electrons |
|
Definition
| because it has only 1 s orbital |
|
|
Term
| stability of orbital vs. distance from nucleus |
|
Definition
| the further from the nucleus, the less stable the orbital |
|
|
Term
| where the node is in a p orbital |
|
Definition
|
|
Term
| the most common elements we will deal with in OCHEM |
|
Definition
|
|
Term
|
Definition
| an atom's affinity for electrons |
|
|
Term
| general trend for electronegativity |
|
Definition
|
|
Term
| what causes polarity in bonds? |
|
Definition
|
|
Term
| some elements that have the same electronegativities as C |
|
Definition
|
|
Term
| why the C-O bond is more polar than the C-Cl bond |
|
Definition
| because O is 2 elements to the right of C and Cl is just 1 element to the right of S |
|
|
Term
| # of nodes vs. energy of orbital |
|
Definition
| the more nodes, the higher the energy |
|
|
Term
| the difference between a 1st energy orbital (such as 1s) and a 2nd energy orbital (such as 2s) |
|
Definition
|
|
Term
| the type of node in an s orbital |
|
Definition
|
|
Term
| why the valence electrons are so special |
|
Definition
| because the define the chemistry and size of the element |
|
|
Term
| trend in size of atoms as you go down the columns of the PTable |
|
Definition
| size increases, but by increasingly smaller increments |
|
|
Term
| trend in size of atom as you go right along the rows of the PTable |
|
Definition
| size of the atom decreases |
|
|
Term
| electronegativity vs. size of atom in the same row |
|
Definition
|
|
Term
| the 2 main types of bonds |
|
Definition
|
|
Term
|
Definition
| bond created by attraction of electric charge |
|
|
Term
| the elements that typically yield anions |
|
Definition
|
|
Term
| the elements that tend to form anions are the ones that are more... |
|
Definition
|
|
Term
|
Definition
| sharing of electron density between 2 separate nuclei |
|
|
Term
|
Definition
| covalent bonds with unequal sharing of electrons |
|
|
Term
| is bond formation endothermic or exothermic? |
|
Definition
|
|
Term
| which is more stable? molecule or individual atoms? |
|
Definition
|
|
Term
| is bond breaking endothermic or exothermic? |
|
Definition
|
|
Term
|
Definition
| the energy input needed to break bonds apart |
|
|
Term
|
Definition
[image]
the bond length is nominal bond length |
|
|
Term
|
Definition
|
|
Term
| what happens to the energy of the bond as the nuclei get further apart? |
|
Definition
|
|
Term
| how wavelength is related to energy |
|
Definition
ΔE = hv
ΔE = change in energy
h = Planck's constant
v = wavelength (nm) |
|
|
Term
|
Definition
|
|
Term
| bond length vs. bond strength |
|
Definition
|
|
Term
|
Definition
| 1 / 1 billionth
1 x 10-9 meter |
|
|
Term
|
Definition
| 1 / trillionth
1 x 10-12 meter |
|
|
Term
| frequency of vibration of bonds vs. length of bond |
|
Definition
|
|
Term
| why the length and strength of bonds between atoms is important for chemistry |
|
Definition
| because it helps predict reactivity |
|
|
Term
| length of bond between atoms vs. reactivity |
|
Definition
|
|
Term
| how size of atom affects reactivity |
|
Definition
| affects internuclear distance, thus affecting bond length and frequency |
|
|
Term
| is the covalent bond between 2 atoms static? |
|
Definition
|
|
Term
| how a covalent bond between 2 atoms vibrates |
|
Definition
| at a specific frequency characteristic to that bond |
|
|
Term
| depiction of where this equation, ΔE = hv, fits into internuclear bonds |
|
Definition
|
|
Term
| Heisenberg's uncertainty principle |
|
Definition
| you can know either the location or energy of the electron, but not both |
|
|
Term
| the amount of time an electron spends in its specific orbital |
|
Definition
|
|
Term
|
Definition
|
|
Term
| depiction of how to draw a p ordital |
|
Definition
[image]
one lobe is shaded to indicate that one of the numbers is positive
the horizontal line is a plane where the node is |
|
|
Term
| the 2 key types of orbitals you gotta pick up on in molecular orbital theory |
|
Definition
|
|
Term
|
Definition
| as you increase the energy, you increase the # of nodes |
|
|
Term
| depiction of how s orbitals interact in covalent bonding |
|
Definition
[image]
the one on top shows antibonding orbitals as a result of destructive interference at high energy and the one on bottom shows bonding molecular orbitals as a result of constructive interference at low energy
the arrows at the bottom represent the opposite spins of paired electrons |
|
|
Term
| what a theory is used for |
|
Definition
| it's used to make predictions |
|
|
Term
| example of how to characterize a bond (ON TEST!) |
|
Definition
[image]
this shows the orbitals involved in forming the bond |
|
|
Term
| depiction of how electron configuration contributes to covalent bonds when bonding atoms are tetrahedral |
|
Definition
|
|
Term
| how to characterize double bonds when the bonding atoms are trigonal planar |
|
Definition
|
|
Term
| why the bond angles that are assumed are often approximations |
|
Definition
| because of the molecule not being entirely symmetrical |
|
|
Term
| the orbitals that are relevant in VSEPR theory |
|
Definition
|
|
Term
| the type of orbitals used by sigma bonds |
|
Definition
|
|
Term
| how to characterize triple bonds when the bonding molecule is linear |
|
Definition
[image]
remember, a triple bond contains 1 sigma bond and 2 pi bonds |
|
|
Term
| what a triple bond contains |
|
Definition
| 1 sigma bond and 2 pi bonds |
|
|
Term
| depiction of unhybridized s and p orbital overlapping |
|
Definition
[image]
the shaded area of the p orbital is positive
everything unshaded is negative |
|
|
Term
| depiction of sp hybrid orbital |
|
Definition
[image]
the shaded area is positive |
|
|
Term
| the effect of hybridization on the sp orbital |
|
Definition
destructive interference causes the positive lobe to decrease in amplitude and the negative lobe to increase in amplitude
[image] |
|
|
Term
| the orbitals that define shape of an atom |
|
Definition
|
|
Term
|
Definition
| the electrons that are in the hybridized orbitals |
|
|
Term
| what causes π (pi) bonds? |
|
Definition
|
|
Term
| the symmetry a sigma bonding orbital has |
|
Definition
| cylindrically symmetrical (cross section along any point in the bond is circular) |
|
|
Term
| if a bond is cylindrically symmetrical, it's a ______ bond |
|
Definition
|
|
Term
| if a bond is not cylindrically symmetrical, it's a ______ bond |
|
Definition
|
|
Term
|
Definition
| the electrons in the unhybridized π (pi) bonds |
|
|
Term
| which electrons are more energetic? those in the sigma system or those in the pi system? |
|
Definition
|
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Term
| why are the electrons in the pi system more energetic than those in the sigma system? |
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Definition
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Term
| which electrons are most likely to be reactive in a chemical reaction? |
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Definition
| the more energetic ones; those in the pi system |
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Term
| can electrons that influence shape of an atom be part of the pi system? |
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Definition
| no, they're part of the sigma system |
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Term
| principal quantum number (n) |
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Definition
| this defines the energy level of an electron |
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Term
| angular momentum quantum number (l) |
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Definition
| this defines the shape of the orbital |
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Term
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Definition
| set of electrons that have the same n and l values |
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Term
| how a subshell is designated |
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Definition
| by a letter and a number, such as s1, s2, and p2 |
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Term
| magnetic quantum number (ml) |
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Definition
describes the orientation of the orbital in space
can have values of l to -l, including 0 |
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Term
| value of l for the s orbital |
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Definition
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Term
| value of l for the p orbital |
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Definition
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Term
| subshell designation for the s orbital |
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Definition
ns
n = principal quantum number |
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Term
| subshell designation for the p orbital |
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Definition
np
n = principal quantum number |
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Term
| possible values of ml for the s orbital |
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Definition
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Term
| possible values of ml for the p orbital |
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Definition
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Term
| the elements with their valence electrons in the s orbital |
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Definition
| He and the Group 1A and 2 A elements |
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Term
| the elements with their valence electrons in the p orbital |
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Definition
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Term
| the elements with their valence electrons in the d orbital |
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Definition
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Term
| the elements with their valence electrons in the f orbital |
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Definition
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