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CHEM 2
| Question | Answer |
|---|---|
| when we solve the wave equation, what do we get? | the orbital (wave function) |
| what is the final model of the electron? | quantum model |
| what 3 theories describe the nature/behavior of electrons? | De Broglie eq, Heisenberg Uncertainty Principle, Schrodinger's eq |
| De Broglie eq: | λ = h/mv (mass x velocity) |
| Heisenberg Uncertainty Principle: | ΔxΔ(mv) ≥ h/(4π) |
| Shodinger’s Equation: | ĤΨ=EΨ |
| what does De Broglie eq tell us? (2) | Electrons behave both like particles and a wave, their distance from the nucleus is related to the λ of the e- |
| what does the Heisenberg Uncertainty Principle tell us? | We can only know the position or speed of an electron at one time, but not both |
| what does Shodinger’s Equation tell us? | The location of electrons inside the atom |
| 2 ways Shodinger’s Equation gives us electron locations: | By using wavefunction/orbitals, by using Quantum Numbers |
| Are wavefunctions the same as orbitals? | yes |
| what are QNs? | Quantum numbers are the answers we get from solving Shodinger’s Equation |
| what are all the variables equal to in ĤΨ=EΨ? | E = total energy of the system Ψ = wave function/orbitals Ĥ = Hamiltonian Operator (a set of conditions that define the total energy of the system) |
| wave function = | answers and solutions to Schrodinger's wave equation |
| the answers of Schrodinger's wave equation give us what 2 things? | orbitals/wavefunctions and quantum numbers |
| how are orbitals/wavefunctions and quantum numbers related? | QNs describe orbitals/wavefunctions |
| what does x mean in Heisenberg Uncertainty Principle? what does p or mv mean? | uncertainty in position, uncertainty in momentum |
| QNs tell us... | the address of an electron inside an atom |
| what is an orbital? | a region of space where you will find an electron |
| each orbital has what two types of spins? | upspin and downspin |
| every electron in an atom has ___ unique QNs | 4 |
| QN - n (3) | principle QN, describes size and energy levels, values will be integers (1,2,3) |
| QN - l (3) | angular momentum QN, describes the shape of the orbital, can be less than or equal to n |
| which is subshell and which is shell? for l and n | l is subshell and n is shell |
| go through what each l value means for QN | l = 0 (s shape - 1 orbital) l = 1 (p shape - 3 orbitals) l = 2 (d shape - 5 orbitals) |
| QN - ml (3) | magnetic QN, gives us the orbital's orientation, dependent on l value |
| go through what each ml value will be dependent on l | When l=0, m=0 When l=1, m=-1,0,1 When l=2, m=-2,-1,0,1,2 (ml is between -l and l) |
| each orbital can have ___ electrons because of what principle? | 2, Pauli's exclusion principle |
| overview/breakdown of each quantum number and what they're for | n, principle QN (for energy level) l, angular momentum QN (for shape of orbitals) ml, magnetic QN (for orientation of orbitals) |
| STEPS FOR ASSIGNING QNs: | Assign l: With n=1, l can only be 0, bc QN rule says l will start with 0 and go up to n-1. Assign ml, which is based on l QN. Have values from -2 thru 0, +2. l=0, ml = 0 (s value) l=1, m = -1, 0 +1 (p value) l=2, , ml(l) = -2, -1, 0, +1, +2 (d value) |
| Pauli’s exclusion principle: (2) | Each orbital can only have 2 electrons (OPPOSITE SPINS) Every e- has its own unique set of QNs |
| QN - ms (3) | magnet spin, describes electron spin, will be either +1/2 or -1/2 |
| electrons always prioritize filling a(n) ______ orbital first | empty |
| break down periodic table into section: | Rows are principle quantum numbers (n=1, n=2, etc… there’s 7 total). First two column are s block, as well as helium. Right six columns is p block. Middle ten column is d block. Bottom 2 rows are f blocks |
| Aufbau principle: | Fill the lowest energy first, then the higher energy orbitals |
| Hund's rule: (3) | If there's degenerate orbitals, one electron goes into each until all are half-filled, these electrons will all have the same spin One electron goes into each orbital until all are half-filled And these electrons will all have the same spin QN |
| Orbitals with the same energy level are ________ orbitals | degenerate |
| which two elements are exceptions to QN common rules? | Cr and Cu |
| Which of the following transitions in a hydrogen atom will emit a light with the longest wavelength: n=3 to n=1, n=4 to n=2, or n=5 to n=2 | n=4 to n=2 |
| how do you find the number of orbitals from n=4 (or any n value) | square the number!! (this would be 16) |
| columns on the periodic table are called _____ | groups |
| columns on the periodic table are called _____ | periods |
| as the atoms go down the column, they get _____ | bigger |
| looking down the column: as the number of electrons/protons increases, radius _______ | increases |
| looking down the period: as the number of electrons/protons increases, radius _______ | decreases |
| Why is fluorine’s radius smaller than lithium's, when F has more electrons than Li? | can be explained by effective nuclear charge (Zeff), Zeff is higher for fluorine because its smaller, because other electrons are feeling the energy of it moreso (closer together) Lithium’s Zeff is lower because its larger and more spread out |
| MAKE SURE U KNOW ATOM SIZING CHARTS | - |
| equation for finding Z effective: | Zeff = Z (# protons) - Shielding e- (# of core electrons) |
| what does Z mean in the Zeff equation? | # protons |
| what does shielding e- mean in Zeff equation? | # of core electrons |
| what is the Zeff for F? | +7 |
| can we just use the number of valence electrons as the Zeff? why AND why not? | yes, but only for main group elements and neutral atoms. no for some ions, since the number of valence electrons is different than neutral atoms |
| Zeff: | the actual charge from the nucleus felt by outer electrons |
| core e- are... | electrons in the inner energy levels (or shells) of an atom, these electrons are not involved in chemical bonding and are shielded from interactions with other atoms by the outer (valence) electrons |
| can determine # core e- by looking at .... | e- configuration |
| which are core electrons and which are outer electrons for Mg 1s²2s²2p⁶3s² | [1s²2s²2p⁶] are inner energy levels (10 of these for Mg) [3s²] are outer electrons, valence electrons |
| cation sizing (2): | Lost e-, positively charged Smaller than its neutral atom |
| anion sizing (2): | Gained e-, negatively charged Larger than its neutral atom |
| isoelectronic ions (3): | Having the same number of electrons (different atoms, same # of electrons) I.e. F- and Na+ Can use Zeff to decide which is larger/smaller (if table is not given) |
| Electrons closer = what for ion sizes? | ion sizes are smaller |
| ionization energy: | energy required to remove an electron from an atom (cation formation bc we’re losing electrons) |
| IE increases as _____ gets larger | Z (protons) |
| IE decreases as _____ gets larger | n (bc of shielding) |
| Why is the IE lower for O than N? | Can be explained by looking at the orbital filling diagrams and the electron being removed (red arrow indicates electron being removed) |
| what does IE2 mean? (or IE#) | energy to remove 2nd electron from an atom |
| electrons that are closer to the ______ and in a _____ shell are harder to remove | nucleus, full |
| 3 periodic trends: | sizing trends, ionization energy, electron affinity |
| how can bond polarity be calculated? | by using electro negativity difference |
| when ΔEN = 0, we have | a pure covalent bond |
| when ΔEN = 0.4, we have (2) | this is our cutoff, non-polar covalent bond |
| when ΔEN = greater than 0.4, we have | a polar covalent bond |
| what makes ΔEN = 0 pure? | there's a symmetric distribution of electrons |
| the arrow for bond polarity points to the _____ sign | negative |
| steps to draw a lewis structure: | sum the valence e- from all the atoms, write the skeletal structure (arranged around central atom), use a pair of e- to form a bond (line), arrange remaining e- to satisfy octet rule |
| octet rule: | atoms must be surrounded by 8 valence electrons |
| when do we use double/triple bonds: | when you don't have enough electrons for the octet rule |
| how do we write a skeletal structure? (2 notes about it) | the least EN atom is usually the central atom, H is always the terminal atom (the one at the edge) |
| lewis structure: (2) | help us visualize molecule in a 2D space, generally obeys the octet rule |
| some exceptions for the octet rule? what rule do we use instead? | H & HE. we will use the duet rule |
| what type of bond does the dash represent in a lewis structure? | covalent |
| what is a lone pair of electrons? | 2 electrons not bonded to anything |
| what is a bonding pair of electrons? | 2 electrons between atoms |
| are smaller or larger electrons going to be more reactive? | larger |
| lower ionization energy = ____ reactive | more |
| electron affinity (EA): | energy released when an e- is added to an atom (ex: anion formation) |
| for electron affinity values, what does the - mean? | energy was released |
| where does electron affinity peak? | at halogens (17/17A) |
| where does electron affinity dip? | alkali earth metals |
| what does it mean if electron affinity peaks at halogens and dips at alkali earth metals? (3)-what does it mean for each | Halogens like to gain e- to form anions (easier to gain it) Alkali earth metals more likely to lose e- and become cations (easier to lose it) Fully filled ones (Mg): almost never form as an anion/don’t like to lose electrons |
| covalent bonds: | when one or more pairs of valence electrons that are shared between atoms of nonmetals (valence electrons: outermost shell e-) |
| one bond has how many electrons? | 2 |
| 2 types of covalent bonds: | polar and nonpolar |
| electronegativity (EN): | a measure of the ability of an atom to attract bonding electrons to itself |
| which one is a scale? EN or EA? | EN |
| write out the EN scale and check | - |
| ionic bonds: | when valence electrons are transferred from the metal to the nonmetal atoms |
| what is bond energy? | energy required to break bonds |
| review atomic sizing | - |
| atomic radii increase down the _____, decrease across the ______ | group, period |
| IE increases as we move ___ and ___ on the PT | up and right |
| metric prefix: mono= | 1 |
| metric prefix: di= | 2 |
| metric prefix: tri= | 3 |
| metric prefix: tetra= | 4 |
| metric prefix: penta= | 5 |
| metric prefix: hexa= | 6 |
| metric prefix: hepta= | 7 |
| metric prefix: octa= | 8 |
| metric prefix: nona= | 9 |
| metric prefix: deca= | 10 |
| mercury(I) chemical name: | Hg2 2+ |
| ammonium chemical name: | NH4 + |
| nitrite chemical name: | NO2 |
| nitrate chemical name: | NO3- |
| sulfite chemical name: | SO3 2- |
| sulfate chemical name: | SO4 2- |
| hydrogen sulfate (or bisulfate) chemical name: | HSO4 - |
| hydroxide chemical name: | OH - |
| cyanide chemical name: | CN- |
| phosphate chemical name: | PO4 3- |
| hydrogen phosphate chemical name: | HPO4 2- |
| dihydrogen phosphate chemical name: | H2PO4 - |
| thiocyanate chemical name: (2) | NCS - or SCN - |
| carbonate chemical name: | CO3 2- |
| hydrogen carbonate (or bicarbonate) chemical name: | HCO3 |
| hypochlorite chemical name: (2) | ClO - or OCl - |
| chlorite chemical name: | ClO2 - |
| chlorate chemical name: | ClO3 - |
| perchlorate chemical name: | ClO4 - |
| acetate chemical name: (2) | C2H3O2 - or CH3COO - |
| permanganate chemical name: | MnO4 - |
| dichromate chemical name: | Cr2O7 2- |
| chromate chemical name: | CrO4 2- |
| peroxide chemical name: | O2 2- |
| oxalate chemical name: | C2O4 2- |
| thiosulfate chemical name: | S2O3 2- |
| Hg2 2+ | mercury(I) |
| NH4 + | ammonium |
| NO2 | nitrite |
| NO3- | nitrate |
| SO3 2- | sulfite |
| SO4 2- | sulfate |
| HSO4 - | hydrogen sulfate (or bisulfate) |
| OH - | hydroxide |
| CN- | cyanide |
| PO4 3- | phosphate |
| HPO4 2- | hydrogen phosphate |
| H2PO4 - | dihydrogen phosphate |
| NCS - or SCN - | thiocyanate |
| CO3 2- | carbonate |
| HCO3 | hydrogren carbonate |
| ClO - or OCl - | hypochlorite |
| ClO2 - | chlorite |
| ClO3 - | chlorate |
| ClO4 - | perchlorate |
| C2H3O2 - or CH3COO - | acetate |
| MnO4 - | permaganate |
| Cr2O7 2- | dichromate |
| CrO4 2- | chromate |
| O2 2- | peroxide |
| C2O4 2- | oxalate |
| S2O3 2- | thiosulfate |
| lewis structures helped visualize the molecule in the ____ space | 2D |
| lewis structures obey the ___ rule | octet |
| what does the octet rule enforce? | main group atoms can gain/lose/share electrons to keep at 8 valence electrons |
| rules for making lewis structures (5) | 1) Count total # of valence electrons 2) The center atom should be the least EN (except H, terminal atom) 3) Use pairs of e- to form bonds between all atoms 4) Use remaining e- to complete octets. 5) run out of e-? multiple bonds can be used |
| how do we find the number of valence electrons? | look at group number (last digit) and add, remember to multiply by 2 if Cl2 or something like that |
| covalent bonds are between ______ atoms | nonmetal |
| ionic bonds are between _____ and _____ atoms | nonmetal and metal |
| covalent bonds naming rules/formula: | need a prefix (mono = not on first element). formula = 1st element name (with prefix) + 2nd element name (with prefix) and end in "ide" |
| 2 examples for covalent bond naming: name CF4 and N2O4 | carbon tetrafluoride, dinitrogen tetroxide |
| ionic bonds naming rules/formula: | cation first, anion second. formula = element name + (roman numberals for type II cations) + either binary compound name ending in "ide"/or polyatomic ion name |
| 2 examples for ionic bond naming: name SrCl2 and PtMnO4 | strontium chloride, platinum (II) maganate |
| the lewis structure represents 2D models. what type represents 3D models? | VSEPR model |
| VSEPR stands for... | Valence Shell Electron Pair Repulsion |
| what does the VSEPR model give us? (2) | the shape/geometry of a molecule, whether the molecule is polar/nonpolar |
| 3 exceptions to the octet rule | expanded octet, incomplete octet, molecules with an odd number of valence electrons |
| Expanded octet: (2 instances and why) | The central atom n has more than 8 e- around it/ If the atom is in the third period (S, Cl, P) or higher (4th, 5th, 6th period, etc) Why? Based on quantum rules, at n=3, electrons can go to the d orbitals, thus able to accommodate the extra e- |
| Incomplete octet: (3) | Compounds of B, some unusual bonding characteristics, often quite reactive |
| Example of an incomplete octet: (practice drawing it out too) | BF3 |
| Molecules with an odd number of valence electrons: (2) | Have at least one of them unpaired Are generally called free radicals (fairly reactive) |
| if the atoms in the structure have FC>0... (2) | Generally the negative charge is on the more electronegative atom, and the positive charge is on the less electronegative one |
| Formal Charge (FC) = | number of valence electrons in the neutral atom - number of lone pair electrons + ½ bonding e- |
| Be able to recognize resonance structures and calculate formal charge | - |
| What are dimers? What molecule tends to form them? | when the lone e- on N is shared, NO2 |
| 5 elements that have exceptions for naming (know what the exceptions are) | Silver, cadmium, platinum, gold, mercury |
| Difference between bond polarity and molecular polarity | BP → Looks at one bond btwn 2 atoms. If atoms have different ENs, electrons are shared unequally → that bond is polar. MP → Looks at the whole molecule. Even if individual bonds are polar, the molecule may be nonpolar if the bonds cancel out |
| again, 4 things VSEPR model tells us | To predict molecular GEOMETRY and SHAPE of molecules To determine if a molecule is POLAR or NONPOLAR |
| basis of the VSEPR model (3) | Electron pairs in bonded atoms repel each other And keep themselves as far away from each other as possible This results in the geometry and shapes we observe |
| Definition of an electron pair (aka electron group or effective pair) | 1 electron pair= 1 lone pair (of electrons) 1 single bond 1 double bond 1 triple bond |
| KNOW ALL THE SPATIAL GEOMETRIES AND SHAPES | - |
| 180 degrees, line shape | linear |
| 120 degrees, triangle shape | trigonal planar |
| 120 degrees, triangle shape with one lone pair | bent/v shaped |
| 109.5 degrees, 4 bonds | tetrahedral |
| 109.5 degrees, 3 bonds and 1 lone pair | trigonal pyramidal |
| 109.5 degrees, 2 bonds and 2 lone pairs | bent/v shaped |
| 90/120 degrees, 4 bonds and 1 lone pair | see-saw |
| 90/120 degrees, 3 bonds and 2 lone pairs | T-shaped |
| 90/120 degrees, 2 bonds and 3 lone pair | linear |
| 90 degrees, 6 bonds | octahedral |
| 90 degrees, 5 bonds and 1 lone pair | square pydramidal |
| 90 degrees, 4 bonds and 2 lone pairs | square planar |
| 90 degrees, 3 bonds and 3 lone pairs | T-shaped |
| 90 degrees, 2 bonds and 4 lone pairs | linear |
| how to use VSEPR model to predict molecular geometry/shape/polarity (5 steps) | draw Lewis structure, count # of electron pairs around central atom (should be as far away from each other as possible), determine geometry, determine the shape, draw bond polarity arrows (using EN) - if arrows cancel out, the molecule is nonpolar |
| When arrows cancel each other out, the molecule is ________. If they don’t it is a ________ molecule | nonpolar, polar |
| the arrow always points to the _____ EN value | more |
| how do we determine if the arrows cancel out? | when they are equal in magnitude and opposite, they will cancel out |
| _____ molecules are typically nonpolar, but only if the same atoms are attached to the central atom | linear |
| molecules with ______ geometry and shapes will be nonpolar, but only if the same atoms are attached to the central atom | symmetric |
| examples of symmetric geometry/shapes: | linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, sq planar |
| example of a POLAR linear molecule, and what would make this occur | HCN, when different atoms are attached to the central atom |
| if the arrows are not equal in magnitude/direction, it is a _____ molecule | polar |
Created by:
stuisl