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OAT General Chem
Atomic and molecular sturcture, periodic properties
| Term | Definition |
|---|---|
| atom | basic building block of matter - the smallest unit of chemical element |
| Dalton's atomic theory | 1. all elements are composed of very small atoms. 2. all compounds are composed of atoms of more than one element. 3. a given chemical reaction involves separation, combination, or rearrangement of atom |
| protons | - carry a positive charge - has a mass of 1 amu - the atomic number of an element = the number of protons found in an atom of the element |
| neutrons | - carry no charge - slightly larger than protons - isotopes of an element have different numbers of neutrons but the same number of protons |
| electrons | - carry a negative charge - very small mass - valence electrons on the outer shell |
| ion | - a positive or negative charge on an atom due to a loss or gain of electrons - a positive charge is a loss of electrons - a negative charge is a gain of electrons |
| mass number | total number of protons and neutrons |
| molecular weight | weight in grams per one mole |
| mole | unit used to count particles and represented by Avogadro's number (6.02 x 10^23) |
| isotopes | - all elements exist as a collection of two or more isotopes - same number of protons different number of neutrons |
| quantum theroy | - developed by Max Planck in 1900 -proposing that energy emitted as electromagnetic radiation from matter comes in discrete bundles (quanta) -the energy value of a quantum is the equation E=hf |
| planck's constant | proportionally constant known as (h) = 6.626 x 10^-34 J*s |
| frequency | (f) sometimes designated (v) is the frequency of the radiation |
| bohr model | - niels bohr (1913) developed a model of electronic structure of the hydrogen atom - the model used the quantum theory, placed conditions on the value of the angular momentum (L = nh/2pi) - bohr equated the values of angular momentum to energy E= Rh/n2 |
| rydberg energy | experimentally determined constant, 2.18 x 10ˉ¹⁸ J/electron known as R(H) |
| quantized | electrons can only exist at specific energy levels, separated by specific intervals |
| orbital radius | the smaller radius, the lower the energy state of the electron smallest orbit an electron can be n=1 |
| ground state | electron is in its lowest energy state |
| electromagnetic energy of photons | E = hc/wavelength c= velocity of light in a vacuum |
| line spectrum | the spectrum composed of light at specific frequencies, each line on the emission spectrum corresponds to a specific electronic transition |
| atomic emission spectrum | the set of frequencies of the electromagnetic waves emitted by atoms of the element - can be used as a fingerprint |
| balmer series | spectrum of light when an electron drops to energy level n=2 - visible and UV |
| lyman series | group corresponding to transitions between upper levels n=1 - UV light |
| paschen series | infrared spectra that are released n>3 |
| absorption spectrum | the wavelengths of absorption correspond directly to the wavelength of emission since the energy difference between levels remains unchanged |
| heisenberg uncertainty prinicple | impossible to simultaneously determine with perfect accuracy the momentum and position of an electron |
| quantum numbers | n, l, m(l), and m(s) |
| pauli exclusion prinicple | no two electrons in a given atom can process the same set of four quantum numbers |
| energy state | the position and energy of an electron described by its quantum number - the value of n limits the values of l which limits the values of m(l) |
| prinicipal quantum number | - first quantum number (denoted by the letter n) - take on any positive integer - the maximum n that the electrons of an element at its ground state corresponds with that element's period in the periodic table - the max # of e- = 2n2 |
| shell | an electron is present in an atom |
| azimuthal quantum number | - second quantum number (denoted by the letter l) - tells us the shape of the orbitals (subshells; s,p,d,f) - the value l can be any integer in range of 0 to n-1 - 4l +2(first two columns = s subshell, right block = p subshell, columns 3-12 = d block) |
| magnetic quantum number | - third quantum number (denoted by the letter m(l)) - describes the orientation of the orbital in space - values all integers from l to -l including 0 |
| spin quantum number | - fourth quantum number (denoted by m(s)) - the spin of a particle (two spin orientation are designated +1/2 and -1/2) - electrons in different m(l) values with the same m(s) are parallel spin - electrons with different m(s) in the same m(l) is paired |
| electron configuration | the pattern by which subshells are filled and the number of electrons with each principal level and subshell |
| aufbau prinicple | subshells are filled from lowest to highest energy and each subshell will fill completelu before electrons begin to enter the next one |
| n+l rule | -used to rank subshells by increasing energy - the lower the sum of the firts and second quantum numbers the lower the energy of the subshell |
| hund's rule | within a given suborbital, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins - electrpns 'prefer' empty orbitals to half-filled ones |
| paramagnetic | has unpaired electrons, a magnetic field will align the spins of the electrons and weakly attract the atom to the field |
| diamagnetic | have no unpaired electrons and are slightly repelled by a magnetic field |
| valence electrons | - group 1 and 2 = s electrons outermost - groups 3-8 = s and p electrons outermost - transition = s and d outermost subshells - lanthanide and actinide = s, d, and f outmost subshells |
| periodic law | the chemical properties of the elements are dependent in a systematic way upon their atomic numbers |
| periods | 7 rows - representing the principal quantum numbers n=1 to n=7 |
| groups | columns - represent elements that have the same electronic configuration in their valence = similar chemical properties |
| representitve elements (A elements) | - either have s or p sublevels as therir outermost orbitals - in groups 1A - 7A, all of which have incompletely filled s or p subshells of the highest prinicipal number |
| nonrepresentitive elements (B elements) | - transition elements, partly filled d sublevels - the lanthanide and actinide, partly filled f sublevels |
| inert (noble gases) | - group 8A - stable, fully-filled formations |
| effective nuclear charge (Zeff) periods | a trend from left to right, protons are added one at a time and the electrons of the outermost shell experience an increasing degree of nucleat attraction becoming closer and tightly bound to the nucleus (net positive charge from the nucleus) |
| effective nuclear charge (Zeff) groups | from top to bottom, the outermost electrons become less tightly bound to the nucleus, the number of filled principal energy levels are shielded from the attraction by the nucleus increases downward |
| atomic radii | - decreases across a period from left to right and increases down a given group - largest radii are located at Group 1A at the bottom towards the left |
| factors that effect atomic radii | - altering the electron cloud changes the radius - the more positive Zeff, the smaller the radius (left to right) - valance electrons that are farther from the nucleus will be more negative charge, increasing atomic radii (top to bottom) |
| ionic radius | radius of a cation or an anion - will affect the physical and chemical properties of an ionic compound - cations (+) will be smaller; anions (-) larger radius |
| ionization energy | the energy required to completely remove an electron from a gaseous atom or ion (endothermic) - the closer and more tightly bound an electron to the nucleus, the higher the ionization energy - increases from left to right (period) - decreases up->down |
| first ionization energy | energy required to remove one valence election from the parent atom |
| second ionization energy | the energy needed to remove a second valence electron from the univalent ion to form the divalent ion - second ionization is usually greater than first ionization energy |
| electron affinity | the energy change that occurs when an electron is added to a gaseous atom and it represents the ease with ehich the atom can accept an electron - the higher the nuclear charge, the higher the electron affinity |
| positive and negative electron affinities | - positive electron affinity value represents energy release when an electron is added to an atom - negative electron affinity value represents a release of energy |
| electron affinity equation | X (g) + e- => X- (g) - X is an atom of a given element in the gaseous state |
| electron affinity in groups | - in group 2 (alkaline earth metals) have low electon affinity values - in group 7 (halogens) have high electron affinity because the addition of an electron to an atom results in a complete filled shell - in group 8 (noble gases) have zero affinity |
| electronegativity | a measure of the attraction an atom has for electrons in a chemical bond - the greater the electronegativity, the greater the attraction for bonding electrons - increases from left to right (periods) - decreases from top to bottom (group) |
| pauling electronegativity scale | the most common electronegativity scale, where the values range from 0.7 for the most electropositive elements (ex: cesium) to 0.4 fro the most electronegative element (fluorine) |
| electronegivity to effective nuclear charge | - elements with low Zeff will have low electronegativities do not attract electrons strongly - elements with high Zeff will have high electronegativities the strong pull the nucleus has on electrons |
| metals | - located on the left side and in the middle of the periodic table - shiny solids at room temperature and have high melting points and high densities - large atomic radius, low ionization, low electronegativity - good conductors of heat and electricity |
| nonmetals | - located on the right side - brittle in solid state - high ionization and electronegativity - poor conductors of heat and electricity - share the ability to gain electrons easily - partially filled p orbitals |
| metalloids | - found along a diagonal line between metals and nonmetals - densities, boiling points, and metling points fluctuate - electronegativities and ionization energies lie between metals and nonmetals possessing characteristics of both |
| alkali metals | - Group 1A - physical properties common to metals, lower density of other metals - one valence electron - largest atomic radii - highly reactive due to low ionization energy - form univalent cations - low electronegativities |
| alkaline earth metals | - Group 2A - many characteristically metallic properties - dependent on the ease with they lose electrons - two valence electrons - smaller atomic radii than group 1 - divalent cations (form +2) - low electronegativities and positive electron affin |
| carbon group | - Group 4A (containing carbon) - have 2 electrons in their outermost p subshell, configuration that is distant from a noble gas - most stable electron sharing with 4 covalent bonds |
| pnictogens | - Group 5A (contains nitrogen) - mixture of nonmetals (N and P), metalloids (As and Sb), and a metal (Bi) - forms 3 covalent bonds - nitrogen holds a positive charge in organic reactions |
| chalcogens | - Group 6A (contains oxygen) - requiring 2 additional valence electrons to be stable - fairly electronegative - form -2 anions - 2 covalent bonds and two nonbonded pairs |
| halogens | - Group 7A - highly reactive nonmetals with 1 valence electron less than the closest noble gas - form -1 anions - want to donate electrons - reactive towards groups 1 & 2 |
| noble gases | - Group 8A - completely nonreactive - high ionization energies - no electronegativity - low boiling points - gases at room temp |
| transition elements | - Group 1B to 8B - high melting point and boiling point - left to right 5 d prbitals - low ionization and oxidation states |
| hydration complexes | dissolved ions formed with molecules of water |
| subtraction frequencies | frequencies not absorbed give the complexes their characteristic colors |
Created by:
Jalisa.bland
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