Ch. 1: Atomic Structure (Complete)

Question 1

CH 1.3

Question 2

defn: quanta

Answer 2

discrete bundles of energy emitted as electromagnetic radiation from matter

Question 3

eqn: energy of a quantum

Answer 3

E = hf

h = Planck’s constant = 6.626 x 10 ^ - 34 Js
f = frequency of radiation

Question 4

eqn: Bohr’s angular momentum of an electron orbiting a hydrogen nucleus

Answer 4

n = principal quantum number
h = Planck’s constant

changes in quantized amounts

Question 5

when you see a formula on test day, what should you do?

Answer 5

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

Question 6

eqn + what does it tell us conceptually: Bohr’s energy of the electron

Answer 6

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)

Question 7

what does the negative sign in Bohr’s energy of the electron equation represent?

Answer 7

the electron in any of its quantized states in the atom will have an attractive force toward the proton

Question 8

defn: orbit

Answer 8

the defined pathway around a proton forming a dense core, around which a single electron revolves

Question 9

defn: ground state of an atom

Answer 9

the state of lowest energy, all electrons are in the lowest possible orbitals

Question 10

defn: excited state of an atom

Answer 10

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

Question 11

are electrons restricted to specific pathways as Bohr positied?

Answer 11

no, but they tend to be localized in certain regions of space

Question 12

mnemonic: what happens to electrons as they go from a lower to higher energy level?

Answer 12

they get AHED

Absorb light
Higher potential
Excited
Distant from the nucleus

Question 13

at room temperature, are the majority of atoms in a sample in the ground state or excited state?

Answer 13

ground state

Question 14

why do electrons return rapidly to ground state and what happens when they do?

Answer 14

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

Question 15

eqn: electromagnetic energy of emitted photons from ground state transition

Answer 15

h = Planck’s constant
c = speed of light in a vacuum
lambda = wavelength of radiation

Question 16

expln (2): atomic emission spectrum

Answer 16

1. 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
2. each line on the emission spectrum corresponds to a specific electron transition

Question 17

why does each element have its own unique atomic emission spectrum?

Answer 17

because each element can have its electrons excited to a different set of distinct energy levels

Question 18

defn: Lyman series, Balmer series, Paschen series

Answer 18

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

Question 19

defn and concept: the energy of an emitted photon

Answer 19

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

Question 20

expln (2): absorption spectrum

Answer 20

1. 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
2. wavelengths of absorption correspond exactly to wavelengths of emission bc the energy level difference is the same

Question 21

identification of elements in what phase of matter requires absorption spectra?

Answer 21

the gas phase

Question 22

what are the 4 key takeaways of atomic emission and absorption spectra?

Answer 22

1. each element has a characteristic set of energy levels
2. 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
3. electrons absorb energy in the form of light
4. 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

Question 23

CH 1.1

Question 24

where are protons found? what is the charge of a proton? what is the mass of a proton?

Answer 24

the nucleus of an atom

1.6 x 10^-19 = e (written as + 1 or +1e)

1 amu (atomic mass unit)

Question 25

what is the atomic number (Z)?

what is the mass number (A)?

Answer 25

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

Question 26

which of the following is a unique identifier for each element? atomic number or mass number? why?

Answer 26

Atomic number because elements are defined by the number of protons they contain

Question 27

what is the charge of neutrons? is the mass of neutrons greater than or less than the mass of protons?

Answer 27

neutral (no charge)

mass = slightly larger than that of the proton

Question 28

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?

Answer 28

PROTONS and ATOMIC NUMBER: constant

NEUTRONS and MASS NUMBER: variable

Question 29

defn + char(1): isotope

Answer 29

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

Question 30

what is convention for showing the mass number and atomic number of an atom?

Answer 30

Question 31

char (4): electrons

Answer 31

1. move through the space surrounding the nucleus
2. associated with varying levels of energy –> move around the nucleus at varying distances, which correspond to varying levels of electrical potential energy
3. charge = -1 (also written as - 1 e)
4. mass approximately 1/2000 of a proton

Question 32

what is the implication of the fact that subatomic particles’ masses are so small?

Answer 32

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

Question 33

do electrons close to the nucleus have higher or lower energy? what about those further out from the nucleus?

Answer 33

CLOSE to the nucleus: lower energy levels

FAR from the nucleus (in higher electron shells): have higher energy

Question 34

func + char (3): valence electrons

Answer 34

1. electrons that are farthest from the nucleus
2. have the strongest interactions with the surrounding environment
3. have the weakest interactions with the nucleus
4. they determine the reactivity of an atom

Question 35

what does the sharing or transferring of valence electrons in bonds allow for?

Answer 35

for elements to fill their highest energy level to increase stability

Question 36

why are valence electrons much more likely to become involved in bonds with other atoms?

Answer 36

because they experience the least electrostatic pull from their own nucleus

Question 37

defn: cation v. anion

Answer 37

cation = a positively charged atom

anion = a negatively charged anion

Question 38

CH 1.2

Question 39

which is reported on the periodic table: atomic mass or atomic weight?

Answer 39

atomic weight! –> constant for a given element

atomic mass varies from one isotope to another

Question 40

is the number of protons in a neutral atom equal to the number of neutrons? or the number of electrons?

Answer 40

electrons!

Question 41

what other number is atomic mass almost equal to? why are they slightly different?

Answer 41

mass number (sum of protons and neutrons)

some mass is lost as binding energy

Question 42

what differentiates different isootopes?

Answer 42

they have varying mass numbers which happens due to a varying number of NEUTRONS, the number of protons is consistent

Question 43

what are the names for the 3 isotopes of H?

Answer 43

1. protium (1 proton, atomic mass = 1 amu)
2. deuterium (1 proton, 1 neutron, atomic mass = 2 amu)
3. tritium (1 proton, 2 neutrons, atomic mass = 3 amu)

Question 44

defn: atomic weight

Answer 44

the weighted average of the different isotopes

the number reported on the periodic table

Question 45

half-life is a marker of stability, what implication does this have on the abundance of different isotopes?

Answer 45

longer-lasting isotopes are more abundant because half-life is a marker of stability

Question 46

why is atomic weight have such utility?

Answer 46

it represents both the mass of the “average” atom of that element AND the mass of one mole of the element

Question 47

defn: Avogadro’s number

Answer 47

6.02 x 10^23

Question 48

defn: mole

Answer 48

a number of “things” equal to Avogadro’s number

Question 49

CH. 1.4

Question 50

defn: orbital

Answer 50

regions of space around the nucleus that electrons are localized within and move rapidly within

Question 51

defn: Heisenberg uncertainty principle

Answer 51

It is impossible to simultaneously determine with perfect accuracy the momentum and the position of an electron

Question 52

defn: Pauli exclusion principle

Answer 52

no two electrons in a given atom can possess the same set of four quantum numbers (n, l, ml, ms)

Question 53

defn: an electron’s energy state

Answer 53

the position and energy of an electron described by its quantum numbers

Question 54

analogy: the 4 quantum numbers

Answer 54

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)

Question 55

defn: principle quantum number (n)

Answer 55

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

Question 56

the difference in energy between two shells decreases as the distance from the nucleus … decreases or increases? as a function of what?

Answer 56

INCREASES bc the energy different is a function of
[1/ni^2 - 1/nf^2]

Question 57

defn: azimuthal (angular momentum) quantum number (l)

Answer 57

refers to the shape and number of subshells within a given principal energy level (shell)

important implications for chemical bonding and bond angles

Question 58

what is the relationship of l and n?

Answer 58

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

Question 59

defn: spectroscopic notation

Answer 59

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

Question 60

eqn: maximum number of electrons within a subshell

Answer 60

4l + 2

where l is the azimuthal quantum number

Question 61

describe the pattern of energies of the subshells

Answer 61

the energies of the subshells increase with increasing l value, but the energies of subshells from different principal energy levels may overlap

Question 62

defn + possible values: magnetic quantum number (ml)

Answer 62

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

Question 63

how are the shape of atomic orbitals determined? what are the shapes of the s and p orbitals?

Answer 63

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)

Question 64

defn: spin quantum number (ms)

Answer 64

an electron has two spin orientations: +1/2 and -1/2

Question 65

defn: paired

Answer 65

whenever two electrons are in the same orbital, they must have opposite signs

Question 66

defn: parallel spins

Answer 66

electrons in different orbitals with the same ms values

Question 67

defn + setup (3): electron configuration

Answer 67

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

Question 68

Think about how electron subshells are filled – explain

Answer 68

Question 69

defn + aka: Aufbau principle

Answer 69

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

Question 70

func + defn: n + l rule

Answer 70

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

Question 71

what happens if if two subshells possess the same n + l value?

Answer 71

the subshell with the lower n value has a lower energy and will fill with electrons first

Question 72

The f block is not usually presented in its actual position, where is its actual position and how does this affect determining electron configurations?

Answer 72

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

Question 73

how can we abbreviation electron configurations?

Answer 73

by placing the noble gas that precedes the element of interest in brackets

Question 74

how do we write electron configuration of an ion (both anion and cation)?

Answer 74

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

Question 75

defn: Hund’s rule

Answer 75

within a given subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins

Question 76

orbitals in what type of subshells fill according to Hund’s rule?

Answer 76

the orbitals in subshells that contain more than one orbital (such as the 2p subshell with 3 orbitals)

Question 77

what is the basis behind Hund’s rule?

Answer 77

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

Question 78

what is an important corollary from Hund’s rule? what two notable exceptions does this create?

Answer 78

1. 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)

Question 79

explain why chromium is an exception to Hunds rule (3)

Answer 79

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

Question 80

explain why copper is an exception to Hunds rule

Answer 80

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

Question 81

are similar shifts seen in other groups? why or why not?

Answer 81

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

Question 82

defn: paramagnetic vs. diamagnetic

Answer 82

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

Question 83

what is the background concept behind paramagnetism and diagmagnetism?

Answer 83

the presence of paired or unpaired electrons affects the chemical and magnetic properties of an atom or molecule

Question 84

defn: valence electrons

Answer 84

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