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MCAT > Gen Chemistry Ch. 1: Atomic Structure > Flashcards

Gen Chemistry Ch. 1: Atomic Structure Flashcards

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MCAT flashcards
1
Q

Proton

A

Positive charge, mass around 1 amu

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2
Q

Neutron

A

No charge, mass around 1 amu

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3
Q

Electron

A

Negative change, negligible mass

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4
Q

Nucleus of an atom

A

Contains the protons and neutrons, electrons move around the nucleus

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5
Q

Atomic number (Z)

A

Number of protons in a given element, unique identifier because elements defined by number of protons

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6
Q

Mass number (A)

A

Sum of an element’s protons and neutrons, can be different within an element because of isotopes

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7
Q

Atomic mass

A

Essentially equal to the mass number, the sum of an element’s protons and neutrons

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8
Q

Isotopes

A

Atoms of a given element (same atomic number) that have different mass numbers, differ in number of neutrons

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9
Q

Isotopes of hydrogen

A

Protium, deuterium, and tritium

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10
Q

Atomic weight

A

Weighted average of the naturally occurring isotopes of an element, listed by periodic table

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11
Q

Rutherford model

A

The atom has a dense, positively charged nucleus that makes up only a small fraction of the volume of the atom

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12
Q

Bohr model of the atom

A

A dense, positively charged nucleus is surrounded by electrons revolving around the nucleus in orbits with distinct energy levels

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13
Q

Planck model

A

The energy difference between energy levels is called a quantum

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14
Q

Quanta

A

Discrete bundles of energy that are emitted as electromagnetic radiation from matter

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15
Q

Quantization

A

There is not an infinite range of energy levels available to the electron; electrons can exist only at certain energy levels

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16
Q

Atomic absorption spectrum

A

Unique to each element, for an electron to jump from a lower energy level to a higher one, it must absorb an amount of energy precisely equal to the energy different between the two levels

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17
Q

Atomic emissions spectrum

A

Unique to each element, when electrons return from an excited state to the ground state, they emit an amount of energy that is exactly equal to the energy difference between the two levels, sometimes the EM energy emitted corresponds to a frequency in the visible light range

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18
Q

Quantum mechanical model

A

Electrons do not travel in defined orbits but rather are localized in orbitals

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19
Q

Orbital

A

A region of space around the nucleus defined by the probability of finding an electron in that region of space

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20
Q

Heisenberg uncertainty principle

A

It is impossible to know both an electron’s position and momentum exactly at the same time

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21
Q

Quantum numbers

A

Completely describe any electron in an atom

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22
Q

Four quantum numbers

A

n, l, ms, ml

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23
Q

Principal quantum number

A

n - described the average energy of a shell

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24
Q

Azimuthal quantum number

A

l - aka angular momentum quantum number, describes the subshells within a given principal energy level (s, p, d, or f), values range from 0 to n-1, s = 0, p=1, d = 2, f =3

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25
Q

Magnetic quantum number

A

ml - specifies the particular orbital within a subshell where an electron is likely to be found at a given moment in time, values from -l to l

26
Q

Spin quantum number

A

ms - indicates the spin orientation of an electron in an orbital (either +1/2 or -1/2)

27
Q

Electron configuration

A

The pattern by which subshells are filled, as well as the number of electrons within each principal energy level and sub shell, uses spectroscopic notation

28
Q

Spectroscopic notation

A

Combining the n and l values as a number and letter, respectively to designate the location of electrons

29
Q

n+1 rule

A

Determines how electrons fill the principal energy levels and subshells according to increasing energy, the lower the sum of n+l, the lower the energy, if n+l is equal than the sub shell with the lowest n value will fill first

30
Q

Hunds rule

A

Describes how electrons fill orbitals, states that subshells with multiple orbitals (p, d, and f) fill electrons so that every orbital in a subshell gets one electron before any of them gets a second

31
Q

Paramagnetic materials

A

Have unpaired electrons that align with magnetic fields, attracting the material to a magnet, causes electrons to have parallel spins

32
Q

Diamagnetic materials

A

Have all paired electrons, which cannot easily be realigned, they are repelled by magnets

33
Q

Valence electrons

A

Electrons in the outermost shell available for interaction (bonding) with other atoms, least electrostatic pull from nucleus

34
Q

Fundamental unit of charge value

A

e = 1.6 e -19 C

35
Q

Attraction between nucleus and electrons

A

Masses so small that electrostatic force of attraction is far greater than the gravitational force of attraction

36
Q

Cation

A

Positively charged atom

37
Q

Anion

A

Negatively charged atom

38
Q

Mole

A

Number of “things” equal to avogadros number

39
Q

Avogadro’s number

A

Na = 6.02 E 23

40
Q

Energy of a quantum

A

Described by Planck relation, E = h*f

41
Q

Plancks constant value

A

6.626 E -34 J*s

42
Q

Bohr angular momentum equation

A

L = n*h/2/pi

43
Q

Bohr energy of an electron equation

A

E = -Rh/(n^2)

44
Q

Rydberg unit of energy value

A

Rh = 2.18 E -18 J/electron

45
Q

Ground state

A

An atom in the state of lowest energy

46
Q

Excited State

A

When at least one electron has moved to a sub shell of higher than normal energy

47
Q

Electromagnetic energy of a photon equation

A

E = h*c/lambda

48
Q

Relation between wavelength and frequency equation

A

C = 3E8 = f*lambda

49
Q

Line spectrum

A

Spectrum where each line on the emission spectrum corresponds to a specific electron transition

50
Q

Lyman series

A

Group of hydrogen emission lines corresponding to transitions from energy levels n>=2 to n=1, highest energy

51
Q

Balmer series

A

Group of hydrogen emission lines corresponding to transitions from energy levels n>=3 to n=2, includes 4 wavelengths in the visible region

52
Q

Paschen series

A

Group of hydrogen emission lines corresponding to transitions from energy levels n>=4 to n=3, highest wavelength

53
Q

Energy associated with energy change equation

A

E = hc/lambda = -Rh[1/ni^2 - 1/nf^2]

54
Q

Pauli exclusion principle

A

No two electrons in a given atom can possess the same set of four quantum numbers

55
Q

Maximum number of electrons within a shell equation

A

num = 2*n^2

56
Q

Maximum number of electrons within a sub shell equation

A

num = 4*l + 2

57
Q

Paired electrons

A

Electrons in the same orbital that has opposite spins

58
Q

Parallel spins

A

Electrons in different orbitals that have the same spine

59
Q

Aufbau principle

A

aka building-up principle; electrons fill from lower to energy subshells and each upswell will fill completely before electrons begin to enter the next one

60
Q

Electron configuration of cations

A

Start with neutral atom, remove electrons from the subshells with the highest n first; if multiple are tied with highest n value, remove from the highest l among them

61
Q

Electron configuration exceptions

A

Chromium and copper groups because half filled and fully filled orbitals are more stable so they will take 1 electron from the s orbitals to create either a half filled or fully filled d orbital

MCAT (67 decks)

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