Ch. 1: Atomic Structure (Complete) Flashcards
CH 1.3
defn: quanta
discrete bundles of energy emitted as electromagnetic radiation from matter
eqn: energy of a quantum
E = hf
h = Planck’s constant = 6.626 x 10 ^ - 34 Js
f = frequency of radiation
eqn: Bohr’s angular momentum of an electron orbiting a hydrogen nucleus
n = principal quantum number
h = Planck’s constant
changes in quantized amounts
when you see a formula on test day, what should you do?
focus on ratios and relationships –> this simplifies our calculations to a conceptual understanding which is usually enough to lead us to the right answer
the MCAT tends to ask how one variable might affect another variable, rather than a plug-and-chug application of complex equations
eqn + what does it tell us conceptually: Bohr’s energy of the electron
R(H) = Rydberg unit of energy = 2.8 x 10 ^ -18 J/electron
changes in quantized amounts
the energy of an electron increases (becomes less negative) the farther out from the nucleus it’s located (increasing n)
what does the negative sign in Bohr’s energy of the electron equation represent?
the electron in any of its quantized states in the atom will have an attractive force toward the proton
defn: orbit
the defined pathway around a proton forming a dense core, around which a single electron revolves
defn: ground state of an atom
the state of lowest energy, all electrons are in the lowest possible orbitals
defn: excited state of an atom
when at least one electron has moved to a subshell of higher than normal energy
are electrons restricted to specific pathways as Bohr positied?
no, but they tend to be localized in certain regions of space
mnemonic: what happens to electrons as they go from a lower to higher energy level?
they get AHED
Absorb light
Higher potential
Excited
Distant from the nucleus
at room temperature, are the majority of atoms in a sample in the ground state or excited state?
ground state
why do electrons return rapidly to ground state and what happens when they do?
WHY: the lifetime of an excited state is brief
WHAT: there is an emission of discrete amounts of energy in the form of photons
the wavelength of the photon is characteristic of the specific energy transition it undergoes
eqn: electromagnetic energy of emitted photons from ground state transition
h = Planck’s constant
c = speed of light in a vacuum
lambda = wavelength of radiation
expln (2): atomic emission spectrum
- emissions from electrons dropping from an excited to ground state are quantized into photons and give rise to fluorescence –> we see the color of the emitted light –> so the spectrum is composed of light at specified frequencies
- each line on the emission spectrum corresponds to a specific electron transition
why does each element have its own unique atomic emission spectrum?
because each element can have its electrons excited to a different set of distinct energy levels
defn: Lyman series, Balmer series, Paschen series
Lyman: the group of hydrogen emission lines corresponding to transitions from energy levels n ≥ 2 to n = 1
Balmer: the group corresp. to transitions from energy levels n ≥ 3 to n = 2
Paschen: the group corresp. to n ≥ 4 to n = 3
defn and concept: the energy of an emitted photon
E = hc/lambda = Rh (1/ni^2 - 1/nf^2)
The energy of the emitted photon corresponds to the difference in energy between the higher-energy initial state and the lower-energy final state
expln (2): absorption spectrum
- When an electron is excited to a higher energy level, it must absorb exactly the right amount of energy to make the transition –> exciting the electrons of a particular element results in energy absorption at specific wavelengths
- wavelengths of absorption correspond exactly to wavelengths of emission bc the energy level difference is the same
identification of elements in what phase of matter requires absorption spectra?
the gas phase
what are the 4 key takeaways of atomic emission and absorption spectra?
- each element has a characteristic set of energy levels
- for electrons to move from a lower energy level to a higher energy level, they must absorb the right amount of energy to do so
- electrons absorb energy in the form of light
- when electrons move from a higher energy level to a lower energy level, they emit the same amount of energy in the form of light
CH 1.1
where are protons found? what is the charge of a proton? what is the mass of a proton?
the nucleus of an atom
1.6 x 10^-19 = e (written as + 1 or +1e)
1 amu (atomic mass unit)
what is the atomic number (Z)?
what is the mass number (A)?
the atomic number of an element is equal to the number of PROTONS found in an atom of that element
the mass number of an element is the sum of protons and neutrons in atom’s nucleus
which of the following is a unique identifier for each element? atomic number or mass number? why?
Atomic number because elements are defined by the number of protons they contain
what is the charge of neutrons? is the mass of neutrons greater than or less than the mass of protons?
neutral (no charge)
mass = slightly larger than that of the proton
for a given element, is the number of protons constant or variable? is the number of neutrons constant or variable?
is atomic number variable or constant? is mass number variable or constant?
PROTONS and ATOMIC NUMBER: constant
NEUTRONS and MASS NUMBER: variable
defn + char(1): isotope
atoms that share the same atomic number but have different mass numbers (same number of protons, different number of neutrons)
generally exhibit similar chemical properties
what is convention for showing the mass number and atomic number of an atom?
char (4): electrons
- move through the space surrounding the nucleus
- associated with varying levels of energy –> move around the nucleus at varying distances, which correspond to varying levels of electrical potential energy
- charge = -1 (also written as - 1 e)
- mass approximately 1/2000 of a proton
what is the implication of the fact that subatomic particles’ masses are so small?
the electrostatic force of attraction between the unlike charges of the proton and electron is far greater than the gravitational force of attraction based on their respective masses
do electrons close to the nucleus have higher or lower energy? what about those further out from the nucleus?
CLOSE to the nucleus: lower energy levels
FAR from the nucleus (in higher electron shells): have higher energy
func + char (3): valence electrons
- electrons that are farthest from the nucleus
- have the strongest interactions with the surrounding environment
- have the weakest interactions with the nucleus
- they determine the reactivity of an atom
what does the sharing or transferring of valence electrons in bonds allow for?
for elements to fill their highest energy level to increase stability
why are valence electrons much more likely to become involved in bonds with other atoms?
because they experience the least electrostatic pull from their own nucleus
defn: cation v. anion
cation = a positively charged atom
anion = a negatively charged anion
CH 1.2
which is reported on the periodic table: atomic mass or atomic weight?
atomic weight! –> constant for a given element
atomic mass varies from one isotope to another
is the number of protons in a neutral atom equal to the number of neutrons? or the number of electrons?
electrons!
what other number is atomic mass almost equal to? why are they slightly different?
mass number (sum of protons and neutrons)
some mass is lost as binding energy
what differentiates different isootopes?
they have varying mass numbers which happens due to a varying number of NEUTRONS, the number of protons is consistent
what are the names for the 3 isotopes of H?
- protium (1 proton, atomic mass = 1 amu)
- deuterium (1 proton, 1 neutron, atomic mass = 2 amu)
- tritium (1 proton, 2 neutrons, atomic mass = 3 amu)
defn: atomic weight
the weighted average of the different isotopes
the number reported on the periodic table
half-life is a marker of stability, what implication does this have on the abundance of different isotopes?
longer-lasting isotopes are more abundant because half-life is a marker of stability
why is atomic weight have such utility?
it represents both the mass of the “average” atom of that element AND the mass of one mole of the element
defn: Avogadro’s number
6.02 x 10^23
defn: mole
a number of “things” equal to Avogadro’s number
CH. 1.4
defn: orbital
regions of space around the nucleus that electrons are localized within and move rapidly within
defn: Heisenberg uncertainty principle
It is impossible to simultaneously determine with perfect accuracy the momentum and the position of an electron
defn: Pauli exclusion principle
no two electrons in a given atom can possess the same set of four quantum numbers (n, l, ml, ms)
defn: an electron’s energy state
the position and energy of an electron described by its quantum numbers
analogy: the 4 quantum numbers
is like an address!
one lives in a certain state (n)
in a certain city (l)
on a certain street (ml)
at a particular house number (ms)
defn: principle quantum number (n)
larger n = higher energy level and radius of the electron’s shell
within each shell, there is a capacity to hold a certain number of electrons given by 2n^2 where n is the principal quantum number
the difference in energy between two shells decreases as the distance from the nucleus … decreases or increases? as a function of what?
INCREASES bc the energy different is a function of
[1/ni^2 - 1/nf^2]
defn: azimuthal (angular momentum) quantum number (l)
refers to the shape and number of subshells within a given principal energy level (shell)
important implications for chemical bonding and bond angles
what is the relationship of l and n?
the value of n limits the value of l:
for any given value of n, the range of possible values for l is 0 to (n-1)
the n-value tells you the number of possible subshells
defn: spectroscopic notation
the shorthand representation of the principal and azimuthal quantum numbers
principal quantum number remains a number, azimuthal quantum number is designated by a letter
l = 0 = s
l = 1 = p
l = 2 = d
l = 3 = f
eqn: maximum number of electrons within a subshell
4l + 2
where l is the azimuthal quantum number
describe the pattern of energies of the subshells
the energies of the subshells increase with increasing l value, but the energies of subshells from different principal energy levels may overlap
defn + possible values: magnetic quantum number (ml)
specifies the particular orbital within a subshell where an electron is most likely to be found at a given moment in time
each orbital can hold a maximum of two electrons
the possible values of ml are the integers between -l and l including 0
how are the shape of atomic orbitals determined? what are the shapes of the s and p orbitals?
dependent on the subshell in which they are found
s subshell orbitals: spherical
p subshell orbitals: dumbbell-shaped along x, y, and z axes
defined in terms of probability density (the likelihood that an electron will be found in a particular region of space)
defn: spin quantum number (ms)
an electron has two spin orientations: +1/2 and -1/2
defn: paired
whenever two electrons are in the same orbital, they must have opposite signs
defn: parallel spins
electrons in different orbitals with the same ms values
defn + setup (3): electron configuration
for a given atom or ion, the pattern by which subshells are filled, as well as the number of electrons within each principal energy level and subshell
use spectroscopic notation
- 1st number = principal energy level
- letter = subshell
- superscript = # of electrons in the subshell
Think about how electron subshells are filled – explain
defn + aka: Aufbau principle
aka: the building-up principle
electrons fill from lower to higher energy subshells and each subshell will fill completely before electrons begin to enter the next one
func + defn: n + l rule
used to rank subshells by increasing energy
the lower the sum of the values of the first and second quantum numbers (n + l), the lower the energy of the subshell
what happens if if two subshells possess the same n + l value?
the subshell with the lower n value has a lower energy and will fill with electrons first
The f block is not usually presented in its actual position, where is its actual position and how does this affect determining electron configurations?
the f block fits between the s and d blocks in the periodic table
so when we are trying to find the electron configuration, and we want to read across the periodic table, we need to put the f block in the correct position
how can we abbreviation electron configurations?
by placing the noble gas that precedes the element of interest in brackets
how do we write electron configuration of an ion (both anion and cation)?
anions: have additional electrons that fill according to the same rules as neutral atoms
cations: start with the neutral atom, remove electrons from the subshells with the highest value for n first
- if multiple subshells are tied for the highest n value, then electrons are removed from the subshell with the highest l value among these
defn: Hund’s rule
within a given subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins
orbitals in what type of subshells fill according to Hund’s rule?
the orbitals in subshells that contain more than one orbital (such as the 2p subshell with 3 orbitals)
what is the basis behind Hund’s rule?
electron repulsion = electrons in the same orbital tend to be closer to each other and thus repel each other more than electrons placed in different orbitals
what is an important corollary from Hund’s rule? what two notable exceptions does this create?
- half-filled and fully filled orbitals have lower energies (higher stability) than other states
exceptions:
1. chromium (and other elements in its group)
2. copper (and other elements in its group)
explain why chromium is an exception to Hunds rule (3)
it should have electron configuration [Ar]4s23d4
however moving one electron from the 4s subshell to the 3d subshell allows for the 3d subshell to be half-filled: [Ar]4s13d5
although moving the 4s electron up to the 3d orbital is energetically unfavorable, the extra stability from making the 3d subshell half-filled outweights that cost
explain why copper is an exception to Hunds rule
it should have electron configuration [Ar]4s23d9
but instead has electron configuration [Ar]4s13d10 because a full d subshell outweighs the cost of moving an electron out of the 4s subshell
are similar shifts seen in other groups? why or why not?
similar shifts can be seen with f subshells, but are NEVER observed for the p subshell because the extra stability doesn’t outweigh the cost
defn: paramagnetic vs. diamagnetic
PARAMAGNETIC: materials composed of atoms with unpaired electrons will orient their spins in alignment with a magnetic field, and the material will thus be weakly attracted to the magnetic field
PARAmagnetic means that a magnetic field will cause PARAllel spins in unpaired electrons and therefore cause an attraction
DIAMAGNETIC: materials consisting of atoms that have only paired electrons will be slightly repelled by a magnetic field
given sufficiently strong magnetic fields beneath any object, any diamagnetic substance can be made to levitate
what is the background concept behind paramagnetism and diagmagnetism?
the presence of paired or unpaired electrons affects the chemical and magnetic properties of an atom or molecule
defn: valence electrons
those electrons that are in the outermost energy shell, are most easily removed, and are available for bonding
the “active” electrons of an atom that to a large extent dominate the chemical behavior of an atom
