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MCAT General Chemistry > Chapter 1: Atomic Structure > Flashcards

Chapter 1: Atomic Structure Flashcards

1
Q

Atomic Number

A

number of protons found in an atom of an element (Z)

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2
Q

Mass Number

A

Sum of protons and neutrons in nucleus (A)

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3
Q

What is the convention for nuclear notation?

A
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4
Q

Isotopes

A

Same atomic number, different mass number

Almost all elements exist as 2+ isotopes, which are usually present in the same proportions in any sample of a naturally occurring element

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5
Q

Unit of Charge

A

1.6 x 10-19 C

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6
Q

Protons

A

Found in nucleus of atom, +1e charge, mass of 1 amu

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7
Q

Neutrons

A

No charge, slightly larger mass than proton, can vary in number among atoms of an element

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8
Q

Electrons

A

Move through space around nucleus, different energy levels, -1e charge, mass = 1/2000 of proton

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9
Q

Attraction of subatomic particles

A

Electrostatic attraction > gravitational force

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10
Q

Electron Shell

A

Electrons in outer shells have higher average energy and travel farther from the nucleus than those in inner shells

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11
Q

Valence Electrons

A

Electrons farthest from nucleus; have strongest interactions with environment. They determine the reactivity of an atom.

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12
Q

What allows atoms to increase stability?

A

The sharing or transferring of valence electrons in bonds, which allows elements to fill their highest energy level

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13
Q

How many proton and electrons are present in an atom’s neutral state?

A

Equal numbers of protons and electrons

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14
Q

Cation

A

Positively charged atom

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15
Q

Anion

A

Negatively charged atom

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16
Q

Atomic Mass

A

Mass number (different for each isotope)

Different isotopes have different atomic masses

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17
Q

Atomic Weight

A

Weighted average of the different isotopes of an element, which have their own atomic mass.

Represents both the mass of the “average” atom of that element, in amu, and mass of one mole of the element (in grams)

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18
Q

What is the mass of 1 proton?

A

About 1 amu

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19
Q

What is the size of 1 amu?

A

1/12 the mass of a Carbon-12 atom

1.66 x 10-24 g

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20
Q

Mole

A

Avogadro’s number = 6.02 x 1023

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21
Q

Which electrons have the lowest energy level?

A

The ones closest to the nucleus

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22
Q

Which is stronger, the electrostatic attraction or gravitational force between subatomic particles?

A

Electrostatic attraction, since their mass is really small

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23
Q

What was Rutherford’s idea?

A

Atoms have a small, dense, positively charged nucleus

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24
Q

What did Planck theorize about the emission of energy from atoms?

A

Energy is emitted as EM radiation from matter comes in discrete bundles called quanta

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25
What is the equation for **Planck relation**? (E1.1)
E = hf h = Planck's constant = 6.626 x 10-34 J⋅s f = nu (v) = frequency of radiation
26
What is the equation of speed of light?
v = fλ = c v = velocity f = frequency λ = wavelength c = 3 x 108 m/s
27
What is Bohr's model?
electrons orbit a postively charged, dense nucleus in specific allowable paths called orbits
28
What is the equation for the **angular momentum** of an electron? (E1.2)
n = principal quantum number h = Planck's constant = 6.626 x 10-34 J⋅s Angular momentum changes in discrete amounts with respect to the principal quantum number (since all the variables are constants)
29
What is the equation for the **energy of an electron**? (E1.3)
RH = Rydberg unit of energy = 2.18 x 10-18 J/electron This value also changes discretely.
30
How much energy is there when the proton and electron are completely separated?
None. Zero.
31
Does the energy of an electron increase or decrease as we move farther out from the nucleus?
Increase! As principal quantum (shell) number increases, energy increases
32
What type of path does an electron take, according to Bohr?
An electron revolves in a defined pathway (**orbit**) at a discrete energy value around the dense core proton(s)
33
Ground State
The smallest, lowest energy orbit; all electrons are in the lowest possible orbitals
34
Excited State
At least 1 electron has moved to a subshell of higher than normal energy
35
What has changed between the original ideas of electron localization between then and now?
Electrons are not restricted to specific pathways but tend to be localized in certain regions of space
36
How is energy emitted going from excited to ground state?
It is emitted in discrete amounts of energy in the form of photons
37
What is the equation for the **energy of photons**? (E1.4)
h = Planck's constant c = speed of light in vacuum = 3.00 x 108 m/s λ = wavelength of radiation This equation combines E = hf and c = fλ
38
What is an atomic emission spectrum?
Each element has its electrons excited to a specific set of distinct energy levels
39
Each line on an emission spectrum corresponds to a what?
A specific electron transition
40
What are the 3 series of the hydrogen emission spectrum? (LBP)
1. Lyman series 2. Balmer series 3. Paschen series
41
What energy levels does the Lyman series comprise?
Energy levels n ≥ 2 to n = 1 Has larger transitions than Balmer, since the photon wavelength is shorter in the UV region
42
What energy levels does the Balmer series comprise?
Energy levels n ≥ 3 to n = 2 Includes 4 wavelengths in the visible region
43
What energy levels does the Paschen series comprise?
Energy levels n ≥ 4 to n = 3
44
What is the relationship between energy and wavelength?
Inversely proportional
45
What is the equation of energy for emitted photons in terms of shell level? (E1.5)
Corresponds to the difference in energy between higher energy initial state and lower energy final state Why is it initial minus final? Because the atom emits a photon, therefore there is a negative value for energy, which denotes a decrease
46
What happens when you excite an electron of a particular element?
Energy absorption occurs at specific wavelengths; this is why every element has a specific absorption spectrum
47
Conservation of Energy
∆Eabsorption = ∆Eemission (for any 2 energy levels)
48
Orbitals
Electrons move rapidly and are localized within regions of space around the nucleus called orbitals
49
Can we pinpoint the exact location of an electron?
No, but we can calculate the probability of finding an electron within a region of space surrounding the nucleus
50
Heisenberg Uncertainty Principle
It is impossible to simultaneously determine with perfect accuracy, the momentum and position of an electron
51
What are quantum numbers?
These are descriptors for the position and energy of an electron. There are 4 quantum numbers: n, l, ml, ms (these are listed with increasing specificity)
52
Pauli Exclusion Principle
No 2 electrons in a given atom can possess the same set of quantum numbers
53
Principal Quantum Number
Denoted by "n" Takes any positive integer value Greater n, greater **shell** energy level/radius
54
What is the maximum number of electrons a particular shell can hold? (E1.6)
2n2
55
Azimuthal (angular momentum) quantum number
Denoted by "l" Refers to the shape and number of subshells within a principal energy level (aka shell)
56
How many possible values for l are there for a given value of n?
0 to (n - 1) e.g. for n = 2, the possible values of l are 0 and 1
57
What is the **spectroscopic notation** for principal/azimuthal quantum numbers?
The principal quantum number is denoted by number, the azimuthal is denoted by letter l = 0 = s l = 1= p l = 2 = d l = 3 = f
58
What is the maximum number of electrons that can be held by a particular subshell? (E1.7)
4l + 2
59
Can subshell energy levels overlap?
Yes! Energies of subshells from different energy levels may overlap, e.g. 4s has a lower energy state than 3d
60
What is the relationship between subshell energy and l?
Subshell energy increases as l increases
61
Magnetic quantum number
Denoted by ml Specifies the orbital within a subshell Each orbital holds 2 electrons max
62
What are the possible values of ml given l?
-l to +l including zero e.g. subshell d, with l of 2, has ml (-2 to +2)
63
What does orbital shape depend on?
The type of subshell that is specified, e.g. s is spherical, p is dumbbell-shaped
64
Spin Quantum Number
Denoted by ms Takes on either +1/2 or -1/2 2 electrons in one orbital MUST have opposite spins
65
How do we describe 2 electrons in one orbital with opposite spins?
"paired spins"
66
How do we describe 2 electrons in different orbitals with the same ms values?
"parallel spins"
67
Aufbau principle
Electrons fill from lower to higher subshells; each subshell has to be filled completely before going onto the next one
68
n + l rule
This is what we used to rank subshells by increasing energy. If 2 subshells have the same n + l value, the subshell with the lwoer n value has a lower energy
69
Hund's Rule
In a given subshell, orbitals fill so that there are a max # of half-filled orbitals with parallel spins (due to electron repulsion)
70
What is the stability of half/fully-filled orbitals?
They are more stable / lower energy
71
What are the 2 exceptions for electron configuration?
Chromium: should be [Ar]4s23d4, but it's actually [Ar]4s13d5, so that 3d can be half-filled Copper: should be [Ar]4s23d9, but it's actually [Ar]4s13d10 This happens sometimes with f subshell, but NEVER p subshell
72
How do we describe an element that has unpaired electrons present?
"paramagnetic"
73
How do we describe an element in which only paired electrons are present?
"diamagnetic"
74
What do valence electrons do?
They comprise the outermost energy shell and are most easily removed; they can form bonds by transferring/sharing electrons
75
Where are the valence electrons for Groups 1 and 2?
Only in the s subshell
76
Where are the valence electrons for Groups 13-18?
Only in s and p subshells
77
Where are the valence electrons for the transition elements, Groups 3-12?
Only in s and d subshells
78
Where are the valence electrons for the Lanthanide and Actinide (f group) series?
s and f subshells
79
Octet Rule
Atoms tend to bond in a way that gives them 8 valence electrons
80
Which elements can violate the octet rule?
Those in Period 3 (starting with Na) and below can accept electrons into their d subshell (to hold more than 8 electrons)

MCAT General Chemistry (12 decks)

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