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Chem Exam 2
| Term | Definition |
|---|---|
| characteristics of waves | wavelength, frequency, speed, electromagnetic radiation |
| electromagnetic radiation | provides an important means of energy transfer; one of the ways energy travels through space |
| wavelength | the distance between two consecutive peaks or troughs in a wave |
| frequency | number of waves (cycles) per second that pass a given point in space; low wavelength = high frequency, high wavelength = low frequency |
| speed of light | wavelength x frequency (unit is hertz, Hz) |
| Planck's constant | energy can be gained or lost only in whole number multiples (deltaE = nhv; h = Planck's constant (6.626 x 10^-34 J/s), v = frequency of the radiation); a quantum is a packet of energy, this means radiation waves behave like distinct particles (photons) |
| photoelectric effect | when light strikes a metal surface, electrons can be emitted from the surface; light must be of a certain minimal frequency |
| Einstein's theory of relativity | energy has mass, allowing us to calculate the apparent mass of a photon: m = E/c^2 (m = mass, E = energy, c = speed of light) |
| electrons dual nature | electrons behave as particles, but also as waves when they form a diffraction pattern (when light is scattered from a regular array of points or lines) |
| Bohr model | the electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits; model gives hydrogen atom energy levels consistent with the hydrogen emission spectrum |
| points about Bohr's model (1) | it correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electron |
| points about Bohr's model (2) | as the electron becomes more tightly bound, it's energy becomes more negative relative to the 0 energy reference state (corresponding to the electron being at infinite distance from the nucleus); as electron gets closer to nucleus, energy is released |
| limitations to Bohr's model | cannot be applied to atoms other than hydrogen, it's fundamentally incorrect |
| development of orbitals | Schrodinger's equation allows us to determine the wave functions of an atom in terms of the total energy (E) of the atom; each solution of the equation gives us a wave function or orbital that is characterized by a particular energy value |
| orbital | specific wave function; 1s orbital - wave function corresponding to the lowest energy for the hydrogen atom; wave function provides no info about the detailed path of an electron; orbital is a 3D region where an electron is most likely to be found |
| probability distribution | intensity of color is used to indicate the probability value near a given point in space; represents the square of a wave function |
| radial probability distribution | plots the total probability of finding an electron in each spherical shell versus the distance from the nucleus |
| hydrogen 1s orbital | maximum radial probability; this is most likely the place to find an electron but lots of time, an electron will be somewhere else; size of 1s orbital: radius of the sphere that encloses 90% of the total electron probability |
| principal quantum number | n, has integral values (1, 2, 3, etc.); related to the size and energy of the orbital; as n increases = orbital becomes larger and the electron spends more time farther from the nucleus; increase in n = higher energy because electron is less tightly bound |
| angular momentum quantum number | l, has integral values from 0 to n-1 for each value of n; related to shape of atomic orbitals; value of l for a particular orbital is commonly assigned a letter: l = 0 is s, l = 1 is p, l = 2 is d, l = 3 is f |
| magnetic quantum number | ml, has integral values between l and -l, including 0; the value of ml is related to the orientation of the orbital in space relative to the other orbitals in the atom |
| shape of s orbitals | one lobe/sphere |
| shape of p orbitals | two lobes separated by a node at the nucleus; one lobe has a positive sign, other lobe has a negative sign; occur in 3 orientations (-1, 0, 1 magnetic quantum number - x, y, and z orientations) |
| shape of d orbitals | first appear in n=3; off-axis lobes; lobes have different signs for mathematical function; occur in five orientations (-2, -1, 0, 1, 2 magnetic quantum number - rename to xz, yz, xy, x2-y2, and z2) |
| shape of f orbitals | first appear in n=4; more complex shapes |
| orbital energies | all orbitals with the same n value are degenerate (have the same energy) |
| nodes | areas of high probability separated by areas of zero probability; number of nodes increases as n increases |
| electron spin quantum number | ms, can be +1/2 or -1/2 to imply the electron can spin in two diff directions; in each orbital you can have two electrons and they must have opposite spins |
| Pauli exclusion principle | in an atom, no two electrons can have the same set of four quantum numbers |
| polyelectronic atom | atoms with more than one electron; forces- attraction to nucleus due to opposite charge, repulsion from other electrons due to same charge; since electron paths are unknown, the electron repulsions can't be calculated |
| orbitals in polyelectronic atoms | hydrogen-like orbitals, similar shapes; diff sizes and energies due to added electron repulsion; electrons prefer to be in lowest energy orbital available (spend time closer to nucleus) |
| Aufbau principle | as protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen-like orbitals |
| Hund's rule | lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by Pauli principle in a particular set of degenerate orbitals (unpaired electrons have parallel spins) |
| orbital diagrams | depiction of quantum number assignments for electrons; where electrons are (boxes with arrows); electron configuration = 1s2 2s2, etc. |
| valence electrons | the electrons in the outermost principal quantum level of an atom; involved in bonding; elements in the same group (vertical column) have the same valence electron configuration and show similar chemical behavior |
| noble gas electron configuration | to avoid writing inner-level electrons, can write electron configuration starting from the previous noble gas (ex. for oxygen, original: 1s2 2s2 2p4; noble gas: [He] 2s2 2p4) |
| ionization energy | energy required to remove an electron from a gaseous atom or ion; increases from left to right and decreases going down a group |
| electron affinity | energy change associated with the addition of an electron to a gaseous atom; decreases from left to right and increases going down a group |
| atomic radius | half the distance between the nuclei of atoms in a chemical compound; decreases from left to right and increases down a group |
| ionic bonds | force of attraction between oppositely charged ions, usually between metal and nonmetal; usually balance opposing charges evenly to reach a neutral molecule (cation - positive; anion - negative) |
| covalent bond | formed by sharing electrons between atoms |
| forces in ionic bonds | attraction to the protons in the nucleus |
| bond energy | the energy required to break a bond in a chemical compound |
| size of ions | positive ions- decrease in radius because you're getting rid of an orbital negative ions- increase in radius because you're adding on or filling an orbital |
| isoelectronic ions | ions containing the same number of electrons; size decreases as number of protons increases in a group of isoelectronic ions because electrons experience a greater attraction as the positive charge on the nucleus increases |
| forces in covalent bonds | electron is attracted to protons in both nuclei, there is electron-electron repulsion for both atoms but less for one atom |
| bond length | distance between the nuclei of two atoms in a bond; the position at which the energy associated with the compound is minimal |
| polar covalent bond | unequal sharing of electrons between the two atoms forming the covalent bond (one atom has partial negative charge, one has partial positive) |
| electronegativity | ability of an atom in a molecule to attract shared electrons to itself; differences in electronegativity affect the type of bond formed: no difference (<0.5) = pure covalent; small difference (<2) = polar covalent; large difference (>2) = ionic |
| electronegativity trend | increases left to right, decreases going down a group |
| dipole moment | molecules that have a center of positive charge and a center of negative charge are called dipolar; any diatomic molecule with a polar bond will have a dipole moment (arrow points to partial negative atom and end of arrow goes on partial positive atom) |
| lone pairs | those pairs of electrons to be localized on an atom (not bonded to other atom) |
| bonding pairs | those electrons found in the space between atoms |
| Lewis dot structure | a way of depicting atoms and compounds that shows location of all valence electrons; helps visualize octet rule and appropriate bonding configurations |
| drawing lewis structures | determine number of valence electrons, draw skeleton structure, draw a single bond between each set of atoms, fill in octets of the atoms, working from outside of the molecule in, until all remaining electrons are used up |
| formal charge | number of valence electrons on free atom - number of valence electrons assigned to the atom in the molecule; atoms try to have formal charges as close to 0 as possible, any negative formal charges are expected to reside on the most electronegative atoms |
| resonance | occurs when more than one valid Lewis structure can be written for a particular model; indicates that the resulting electron structure of the molecule is given by the average of the valid Lewis structures |
| VSEPR model | valence shell electron-pair repulsion; molecular structure- 3D arrangement of molecules in an atom; predicts the molecular structure of molecules formed from nonmetals |
| VSEPR main postulate | structure around a central atom is determined by minimizing electron-pair repulsion |
| location of lone pairs in VSEPR structures | lone pairs take up more space than bonding pairs so orientation will minimize electron repulsion; often compresses angle of bonded pairs; two lone pairs will point away from each other |
| VSEPR problem solving | draw lewis structure, determine number of lone pairs and bonds around central atom, number of lone pairs + bonds equals the electronic geometry, number of bonds determines molecular geometry |
| two bonded pairs | linear; pairs are 180 degrees from each other on opposite sides of central atom |
| three bonded pairs | trigonal planar; each pair is 120 degrees from each other |
| four bonded pairs | tetrahedral; 109.5 degree bond angle |
| three bonded pairs and one lone pair | trigonal pyramid; each pair is <120 degrees from each other |
| two bonded pairs, one/two lone pairs | bent; <120 degree bond angles |
| five bonded pairs | trigonal bipyramidal; 120 and 90 degree bond angles |
| six bonded pairs | octahedral; 90 degree bond angles |
| no single central atom | molecular structure can be predicted from the arrangement of pairs around multiple central atoms |
| four bonded pairs and one lone pair | see-saw; 90, 120, and 180 degree bond angles |
| three bonded pairs and two lone pairs | t-shaped; 90 and 180 degree bond angles |
| two bonded pairs and three lone pairs | linear; 180 degree bond angles |
| four bonded pairs and two lone pairs | square planar; 90 and 180 degree bond angles |
| five bonded pairs and one lone pair | square pyramidal; 90 and 180 degree bond angles |
Created by:
jpawloski