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Chem #1 Flashcards

1
Q

what did rutherford contribute to knowledge of atomic structure?

A

described a dense, positive nucleus.

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2
Q

describe a proton

A

1 charge, found in the nucleus of an atom, fundamental unit of charge: 1.6 x 10-19 C.
o Mass: 1 amu (atomic mass unit)

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3
Q

nucleus

A

the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford based on the 1909 Geiger–Marsden gold foil experiment

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4
Q

what is atomic number

A

protons

the number of protons found in an atom of that element.
o Unique identifier for each type of atom.

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5
Q

mass number

A

the sum of the protons and neutrons in the atom’s nucleus. P+N

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6
Q

isotope

A

atoms that share an atomic number but have a different mass number (same element)

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7
Q

neutrons

A

neutral and have no charge, located in the nucleus.
o Mass: just barely larger than a proton
o N + P = almost the entire mass of an ato

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8
Q

if there are 6 protons in an electrically neutral atom, how many electrons are there?

A

6

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9
Q

electron

A

move through the space surrounding the nucleus and are associated with varying levels of energy. Varying distances = varying levels of electrical potential energy.
o Charge: same as a proton but negative
o Mass: 1/2000 that of a proton

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10
Q

electron shells

A

the different energy levels that electrons are held in

 Closer to the nucleus is known as lower energy levels.

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11
Q

valence electrons

A

the electrons that are farthest from the nucleus for any given element.
 Weakest interactions with the nucleus and most interaction with the environment. These are the electrons that participate in bonding.
 Determine the reactivity of an atom.

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12
Q

cation

A

: lose electrons, positively charged atom

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13
Q

anion

A

gain electrons, negatively charged atom

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14
Q

1 amu

A

1.66 x 10^-24 g

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15
Q

atomic mass

A

synonymous with mass number

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16
Q

atomic weight

A

: the weighted average of the different isotopes for an element, the number reported on the periodic table.
o Longer lasting isotopes are generally more abundant, based on half-life.
o There are no atoms with an actual mass of exactly the atomic weight #.
o It is the mass of the average atom of the element, and the mass of one mole of the element, in grams.

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17
Q

what is avogadro’s number

A

6.02 x 10^23

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18
Q

quanta

A

discrete bundles of electromagnetic radiation

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19
Q

what is the planck relation energy of a quantum (relates energy to frequency of an electron)

A

E = hf
o h: planck’s constant, 6.626 x 10-34 Jxs
o f: (sometimes greek nu) frequency of the radiation

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20
Q

what does velocity equal

A

frequency x wavelength

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21
Q

the energy of an electron _____ as it gets further from the nucleus

A

increases

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22
Q

when an electron falls to a lower energy level (closer to the nucleus), what happens?

A

it emits a photon of energy E = hf

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23
Q

ground state

A

the orbit with the smallest, lowest-energy radius, n = 1

o The state of lowest energy in which all electrons are in the lowest possible orbitals

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24
Q

excited state

A

when at least one electron has moved to a subshell of higher than normal energy

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25
energy difference between energy levels closer to the nucleus is ____
greater
26
atomic emission spectra
fingerprint for an element. Each element can have its electrons excited to a different set of distinct energy levels. o Can be used to determine the elements on planets and stars.
27
bohr model of hydrogen
atomic emission spectrum of hydrogen, simplest of all elements.
28
lyman series
the group of hydrogen emission lines corresponding to transitions from energy levels n>/=2 to n=1 o Has larger energy transitions than the Balmer series.
29
balmer series
the group corresponding to transitions from energy levels n>/=3 to n=2 (includes 4 wavelengths in the visible region)
30
paschen series
: transitions from n>/=4 to n=3. | • Equation for energy of released photon for specific drop in energy levels (see specific handwritten notes)
31
compare lyman, balmer, paschen
lyman: to 1 balmer: to 2 paschen: to 3
32
absorption spectrum
characteristic to each element as they can only absorb energy at specific wavelengths. o Correspond exactly to the wavelengths of the emission spectra.
33
heisenberg uncertainty principle
it is impossible to simultaneously determine, with perfect accuracy, the momentum and the position of an electron orbiting an atom.
34
pauli exclusion principle
no two electrons in a given atom can possess the same set of four quantum numbers.
35
what do quantum numbers describe?
o Quantum numbers describe the electrons energy state: position and energy.  Give info on shape, size, and orientation of the orbitals.
36
what is the order of increasing specificity of quantum numbers?
o n  l  ml  ms
37
Principal quantum number
n can take on any positive integer value. The higher the value of n, the higher the energy level and radius of the electron’s shell o Max # of electrons within a shell: 2n2 o The difference in energy between two shells decreases as the distance from the nucleus increases.
38
Azimuthal (angular momentum) Quantum Number
the shape and number of subshells o Important for chemical bonding and bond angles. o l = 0 to (n-1) o Spectroscopic notation: the shorthand representation of the principal and azimuthal quantum numbers.  l=0: s (contains 2 e-, 2 elements)  l=1: p (contains 6 e-, 6 elements)  l=2: d (contains 10 e-, 10 elements)  l=3: f (contains 14 e-, 14 elements) o Maximum number of electrons within a subshell: 4l + 2 o Energies of subshells will increase with increasing l BUT the energy levels of subshells from different principal E levels may overlap
39
what does l = 2 mean?
in the d subshell
40
magnetic quantum number
particular orbital within a subshell where an electron is likely to be found. o Each orbital can hold a max of 2 electrons. o Integers between -l and l and includes 0. o Value of ml represents the number of orbitals in that particular subshell. o Shapes of orbitals are defined in terms of a concept called probability density, the likelihood that an electron will be found in a particular region of space.
41
spin quantum number
o The electron has two possible spin orientations: +1/2 and -1/2. o Paired electrons: electrons in the same orbital must have opposite spins. o Parallel spins: electrons in different orbitals with the same ms value
42
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. o Ex: 2p4: principal energy level 2, p subshell, 4 electrons in it  Assumes the 1s and 2s are filled.
43
aufbau principle
electrons are filled from lower-to-higher energy subshells o Each subshell will completely fill before electrons begin to enter the next one. o n + l rule: can be used to rank subshells by increasing energy (the lower the sum, the lower the energy of the subshell)  if two subshells have the same value for this, the lower n value will fill first.
44
hunds rule
within a subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins. o The basis for this is repulsion. o Half-filled and fully filled orbitals have lower energies (higher stability) than other states. o EXCEPTIONS  Chromium (and other elements in its group): [Ar]4s13d5 • Energetically unfavorable but extra stability  Copper (and other elements in its group): [Ar]4s13d10  Similar shifts seen for the f subshell, but NEVER for the p subshell
45
how to fill and remove electrons from subshells short tricks
o N + L trick to see which subshell gets filled first.  6s2, then 4f o For removing electrons for cations  Start by removing electrons from highest n value first, if tied, remove from highest L value first
46
what are the exceptions to Hunds rule?
o EXCEPTIONS  Chromium (and other elements in its group): [Ar]4s13d5 • Energetically unfavorable but extra stability outweighs this cost.  Copper (and other elements in its group): [Ar]4s13d10  Similar shifts seen for the f subshell, but NEVER for the p subshell
47
paramagnetic
materials composed of atoms with unpaired electrons will orient their spins in alignment with a magnetic field and will thus be weakly attracted to it. o Magnetic field causes parallel spins in unpaired e- and therefore causes attraction.
48
diamagnetic
: materials consisting of atoms that have only paired electrons will be slightly repelled by a magnetic field.
49
valence electrons
o Helps us understand an atom’s properties o The electrons in the atom’s outermost energy shell, most easily removed, and available for bonding. o Groups 1 and 2, highest s subshell o Groups 13-18, highest s and p subshell o Transition elements: highest s and d subshell o Lanthanide and actinide: highest s and f subshell o All elements in period 3 (starting with sodium) and below (larger) may accept electrons into their d subshell, which allows them to hold more than 8 in their valence shell. VIOLATE octet rule.
50
which elements can violate the octet rule?
o All elements in period 3 (starting with sodium) and below (larger) may accept electrons into their d subshell, which allows them to hold more than 8 in their valence shell. VIOLATE octet rule.
51
what are the periods and groups on a periodic table?
period: rows groups: columns
52
which electrons have the greatest potential energy?
valence electrons
53
if the charge on an atom is 0, then _____
there is the same number of electrons as there are protons.
54
name the isotope of hydrogen with: :one proton and 0 neutron :one proton and 1 neutron :one proton and 2 neutrons
protium deuterium tritium
55
no one isotope will have a _______ exactly equal to the element's atomic weight.
mass | For example, the atomic weight of bromine is 79.9 amu but there is no isotope of bromine with this mass.
56
what two things does the atomic weight represent?
the mass of the average atom of an element | the mass of one mole of the element
57
electrons _____ restricted to specific pathways, but tend to be _______
are not | localized in certain regions of space
58
why did Bohr's model fail?
It failed to take into account the repulsion between multiple electrons surrounding the nucleus. Bohr postulated that electrons follow a clearly defined circular pathway or orbit at a fixed distance from the nucleus, whereas modern quantum mechanics has shown that this is not the case (electrons rather move rapidly and are localized within regions of space called orbitals).
59
principal quantum number: The higher the value of n, the ____ the energy level and radius of the electron’s shell
higher
60
the difference in energy between shells ____ as the distance from the nucleus increases.
decreases
61
the l = 0 subshell is which subshell?
s
62
to remove electrons for electron configuration of a cation, first remove from the _________ and if tied then remove from the highest _____ value first.
highest principal quantum number (n) and then the l value (azimuthal)
63
ferromagnetic
already aligned before a magnetic field is even introduced.
64
all elements in period ____ and below can accept electrons into their ____ and thus violate the ___ rule
3 (starting with sodium), d orbitals, octet
65
what is the name of the 4f subshell
lanthanide series
66
what is the name of the 5f subshell
actinide series
67
what is the name of the 5f subshell
actinide series
68
for an electron to jump from a lower energy level to a higher one, it must absorb an amount of energy ______ to the energy difference between the two levels
precisely equal
69
The Rutherford model described ______
a dense, positively charged nucleus
70
____ spins of electrons means their spins are in the same direction
parallel

MCAT Chem (12 decks)

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