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53 Cards in this Set

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  • Back
electrons (location)
outside of the nucleus in specialized regions called orbitals
mass number (A)
atomic number (Z)
number of protons and neutrons in atom
number of protons in atom
A X
Z
atomic mass unit
equivalent to the mass of one proton (one positively charged particle in the nucleus)
the farthest electrons from the nucleus
valence electrons; these participate during bonding
electrons (charge and weight)
magnitude is the same as protons but opposite. weight is 1/1,800 of a proton's.
The farther away the valence electrons,
the more likely they are to be influenced by forces since there is a weaker pull from the nucleus. the farther the valence electrons from the nucleus, the more reactive the element.
a loss or gain of an electron results in
an ion; an atom with a positive or negative charge.
atomic weight standard
carbon is arbitrarily assigned 12 atomic mass units
atomic weight (units)
g/mol; atomic weight is the weight of 6.022 x 10^23 atoms (1 mole) of an element
1 mole
represents 6.022 x 10^23 particles: atoms, molecules, or atomic mass units.
molecular weight
the weight of 6.022 x 10^23 molecules
isotopes
because most elements in nature exist as several isotopes (compounds with the same number of protons and electrons but with different numbers of neutrons), atomic weights for atoms are averages which depend on the relative frequencies of each isotope. this is why most atomic weights are not whole numbers but include decimals.
quantum theory
max planck (1900)theorized that electromagnetic radiation emmitted from matter comes in discrete bundles called quanta. E=hf, energy equals planck's constant times the frequency of the radiation.
Bohr's equation of electric-proton attraction
E=-R/n^2

R=2.18 x 10^-18 J/electron
n=quantum number
the energy or an electron is directly related to its
orbital; the smallest orbital has a quantum number (n) of 1. the smaller the orbital, the lower its energy state.
at normal temperatures, the electrons in an atom are in the
ground state; when excited by energy or heat, they jump to higher levels and subsequently return to their ground state, emmitting energy in the form of light (photons).
Bohr's equation of electromagnetic energy of photons
E=hc/wavelength

h=planck's constant
c=speed of light 3.00 x 10^8 m/s
atomic emmission spectrum
a linear spectrum of light is observed for each atom when its electrons are excited to higher energy levels. because each atom have distinct energy levels, each element possesses a unique atomic emmission spectrum.
Bohr's model of the atom VS. Heisenberg Uncertainty Principle
While Bohr's model describes electrons as following circular orbits at fixed distances around the nucleus, the Heisenberg uncertainty principle described electrons as always being in a random state of motion within regions of space (orbitals); an orbital is a representation of the probability of finding an electron in that area. The principle states that it is impossible to know simultaneously the momentum (speed) and the position of a given electron at a given time. As it becomes easier to calculate its momentum, it is more difficult to locate its position and vice versa.
Modern atomic theory states that any electron can be described by four quantum numbers:
n-principal quantum number
l-azimuthal quantum number
m(l)-magnetic quantum number
m(s)-spin quantum number

n=1,2,3
l= 0 to n-1
m(l)= -l to +l
m(s)= +1/2 or -1/2
principal quantum number: n
integer value; the larger the value of n, the higher the energy level and radius of the electron's orbit.
azimuthal quantum number: l
refers to the subshellls that occur within each principal energy level: s,p,d,f correspond to 0,1,2,3
magnetic quantum number: m(l)
specifies the particular orbital within the subshell in which the electron is highly likely to be found. (x,y, or z)
spin quantum number: m(s)
electrons in the same orbitals must have opposite spins, positive and negative one half.
pauli exclusion principle
no two electrons in a given atom can possess the same set of four quantum numbers.
how many possible values for "l" are there for principal quantum number "n"?
n values for l for n.
how many possible values are there for m(l), for a value of "l"?
how many possible values for m(l) for principal quantum number n?
2l + 1
n^2
how many possible electrons are there for principal quantum number n?
2n^2; two per orbital
parallel vs. paired electrons
paired electrons are electrons in the same orbital with opposite spins. electrons that are in different orbitals with the same spin quantum value have parallel spins.
electron configuration 2p^4...what does each term denote?
2=energy level (n)
p=subshell (l)
4=how many electrons are in the subshell
subshells are filled in what order?
how are subshell energies determined?
which will fill first, 3d or 4s subshell?
from lowest to highest energy; each subshell will completely fill before the next begins filling.

(n + l); the higher this value, the higher the energy level

3d
Hund's rule
orbitals are filled such that therea re a maximum number of half-filled orbitals with parallel spins. Electrons prefer empty orbitals to half-filled orbitals.
order of electrons filling, on periodic table
1s, 2s, 2p, 3s,3p,4s,3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
paramagnetic
has unpaired electrons and weakly attracts an electromagnetic field which aligns the electrons
diamagnetic
has no unpaired electrons and are slightly repelled by a magnetic field.
valence electron
electrons in outermost shell (meaning the ones included after the last [noble gas]). the group number (number at top of column) tells how many valence electrons there are for that element.
representative elements (A)
nonrepresentative elements (B)
groups with s or p subshells as their outermost groups; groups with partially filled d or f orbitals.
elements with partially filled f orbitals
lanthanide and actinide series
going across the periodic table from left to right and down to up...
increasing electronegativity, increasing ionization energy (energy required to remove/add an electron), decreasing atomic radius
noble gases
inert gases; group VIII elements which have stable octet formations
explain the decrease in atomic radius along a period from left to right
as protons are added going from left to right across a period, electrons that are added cannot shield each other from the increasing atomic pull of the protons since they are in the same shell. therefore, the atomic radius decreases across a period
first ionization energy
second ionization energy
the second ionization energy is always larger than the first ionization energy.
which group has the lowest first ionization energy?
the Group I elements have the lowest first ionization energies because by removing exactly one electron, they form a stable octet formation.
which group has the highest electron affinity
the affinity that an atom has to the addition of an electron is highest for the group VII elements.
Electronegativity and ionization energy
the higher the ionization energy (the energy required to remove an electron), the higher the electronegativity (the affinity it has for the electron).
malleability
ability of a metal to be hammered into shapes
ductility
ability of a metal to be drawn into wires
nonmetals are located
on the right side of the periodic table
metals
are good conductors of heat and electricity; have low ionization energies and low electronegativities. also, have luster, high densities and boiling points, and have good malleability and ductility.
nonmetals
are poor conductors of heat and electricity; usually brittle and have high ionization energies and high electronegativities.
most reactive metals
group IA and IIA metals
metalloids
possess characteristics of both metals and nonmetals: BORON, SILCON, GERMANIUM, ARSENIC, ANOTIMONY, TELLURIUM
transition metals
all transition metals are considered metals. have high boiling and melting points. have low ionization energies and are thus found in several oxidation states; have the ability to form complexes. because electrons held loosely, highly malleable and ductile.