Atomic Structure

Question 1

Scientific models

Answer 1

Representations of objects, systems, or events—use familiar objects to represent unfamiliar things; can help scientists communicate their ideas, understand processes, and make predictions.

Question 2

Development of atomic model

Answer 2

Dalton - all matter made of invisible atoms (1803)
Thomson - negative and positive charges (1904)
Rutherford - nucleus (1911)
Bohr (incorporated the subatomic particle theory of matter, which was starting to develop in the late 1800s) - energy levels (1913)
Schrödinger - electron cloud model (1926)
Our knowledge is still developing; we will refine it as we ex. better our technology

Question 3

Thomson’s cathode ray experiment

Answer 3

• Created “cathode rays”—which were charged particles—and sent them through oppositely charged plates and through a magnetic field
• Showed that the “cathode ray” was deflected away from the negatively charged plate (so it was composed of negatively charged particles which could separate from the atom), determined that the mass of the particles was much smaller than any known atom and was constant for different elements: this disapproved Dalton’s theory that atoms are indivisible

Question 4

Rutherford’s gold foil experiment

Answer 4

• Fired positively charged alpha particles (two protons and two neutrons [identical to a He nucleus]) at gold foil
• Most went through the empty space of the gold atom, but some hit the nucleus and were deflected
• Results suggested that an atom was made mostly of empty space with the positive charges grouped together in a very small, dense nucleus

Question 5

Bohr’s quantized shell model

Answer 5

• Theorized that Rutherford’s model had a problem w/ the placement of the electrons (if the electrons were stationary, then they would be attracted to the nucleus, and if they were spinning randomly around the nucleus, they would lose energy and spiral into it [which they don’t] b/c a charged particle moving on a curved path emits electromagnetic radiation)
• His theory fixes this problem by requiring that the electrons move in “permitted orbits” where they don’t lose energy (the energy of the electron depends on the size of the orbit and is lower for smaller orbits; radiation can occur only when the electron jumps from one orbit to another; the atom will be completely stable in the state w/ the smallest orbit since there’s no orbit of lower energy into which the electron can jump)

Question 6

Current model (Schrödinger - “electron cloud” model)

Answer 6

• Based on mathematical wave functions and describes the regions in space, or orbitals, where electrons are most likely to be found
• Describes the probability that an electron can be found in a given region of space at a given time; tells us where an electron might be (not where it is)
• Allows the electron to occupy three-dimensional space

Question 7

Atomic structure

Answer 7

An atom consists of a positively charged dense nucleus composed of protons and neutrons; negatively charged electrons occupy the space outside the nucleus in 3-D orbitals

Question 8

Mass #, atomic #, relative mass + charge #’s

Answer 8

A = the mass # (# of protons + # of neutrons)
Z = atomic # (# of protons in nucleus of atom)
- An atom is neutral, so # of protons (atomic #) = # of electrons; each element has a fixed # of protons
- (Actual masses in data book) Relative mass of protons and neutrons is 1 (electrons, 1/1836 [negligible—we don’t worry about their mass b/c they don’t really contribute anything])
Relative charge of proton is +1; neutrons, 0; and electrons, -1

See nuclear notation (conventionally, atomic # on bottom)

Question 9

Isotopes

Answer 9

Isotope - atoms of the same element w/ the same atomic # and different mass #’s (all just isotopes [not ex. “isotopes of each other”])

Question 10

What determines chemical and physical properties?

Answer 10

• Chemical properties determined by # of electrons in the highest energy level (outer shell): isotopes have the same electron config, so same chemical properties
• Physical properties differ (rates of diffusion; nuclear abilities such as radioactivity and the ability to absorb neutrons; slightly different boiling points)—any that depend on mass
• Radioisotopes (radioactive isotopes) are used in nuclear medicine (for diagnosis and treatment), in biochem research (as tracers and w/ 3-D imaging to detect cancer), and for historical dating

Question 11

What is relative atomic mass?

Answer 11

Ar is an average based on the proportion of isotopes present in nature (abundance): the weighted average (takes into account how much of each one there is) of an atom taking into account the masses of all the isotopes of an element relative to 12C (decided standard [gas would be hard to isolate, work w/]: figured out mass, divided by 12 to get mass of protons, so if something has six protons, they divide it by six)
Ar = (mass of isotope 1)(% isotope 1)/100 + (mass of isotope 2)(% isotope 2)/100 + …
Whichever the # is closer to is the more abundant; no unit b/c it’s relative to something else (for atomic mass, g/mol)

Question 12

Mass spec

Answer 12

Can show the presence of isotopes of elements in a sample and their abundance (allowing calculation of relative atomic mass [splits stuff up according to their mass])
Has five processes that take place: vaporization, ionization, acceleration, deflection, detection
*see notes for details

Question 13

Using mass spec data to calculate RAM

Answer 13

Relative intensity is how much comes out at one mass and how much comes out at another
Add the heights up for full percentage (each height out of full to get your percentages [you could do it w/ a ruler])
m/z basically just mass b/c charge is just +1
Mass spec can provide info for calculation of relative atomic mass
If Ar and isotope masses are known, one can be x, the other (100 - x)
*Ask yourself: does it make sense logically?
Anions have more electrons than protons, so negative charge; cations have fewer electrons than protons, so positive charge

Question 14

Continuous spectrum (ex. visible light spectrum)

Answer 14

Raging Martians Invade Venus Using X-ray Guns
Long wavelength means can transmit really long distances (low wavelength = low energy)
Energy has different wavelengths (frequency is how many peaks go by in a second, and the properties are inversely proportional)
A continuous spectrum shows a spread of energies and a smooth gradual change from one energy to another; rainbows show a continuous spectrum of colors as white light is split into all its different frequencies (in visible light spectrum, violet closest to ultraviolet [highest energy in visible light spectrum] and red closest to infrared)

Question 15

Absorption (excitation), emission

Answer 15

When atoms absorb energy, they become “excited” and some of their electrons jump up to higher energy levels
For gasses, this is achieved by passing an electrical discharge through the gas at low pressure (gives it enough energy [ex. neon lights]); for metals, can be observed by heating them in a Bunsen flame
When the electrons return to lower energy levels, the extra energy is released as EM radiation (ΔE = hf [light frequency of f, h = Planck’s constant = 6.626 x 10-34 J-s])

Question 16

Convergence

Answer 16

Energy levels not equally spaced (get closer to each other as n increases and eventually converge: point is called the convergence, the limit of which [n∞—highest point, when everything converges] indicates the energy required to completely remove the electron from the atom and ionize it—so if you give it enough KE, you could send it across the n∞ level and ionize it [overcome attractive force])
Energy levels further away from the nucleus have more energy

Question 17

Why/How fall back down?

Answer 17

Energy levels still exist even when empty
Again, lowest energy is most stable (also, we’re looking at hydrogen, so… [?])
Put that much energy, not a stable state, so falls back down (transmits back, emits light photon equal to energy put in)
Doesn’t have to fall all the way down to the bottom
Due to the way energy levels are spaced, transitions from n = 2 to n = 1 have greater ΔE than transitions from n = 3 to n = 2 (and ∞ to 2)

Question 18

Relationship between ΔE, frequency, and wavelength

Answer 18

The frequency of the EM radiation emitted when an electron transitions from a higher to a lower energy level is directly proportional to ΔE
Radiation is sometimes defined by its wavelength (𝛌) rather than its frequency (these properties of a wave are related by the equation c = 𝒗𝛌, where c is the velocity of EM radiation (and also speed of light): solve for the frequency and apply it in the other equation
c/𝛌 = 𝒗
ΔE = hc/𝛌 (now, if you know the change in energy, you can find the wavelength and vice versa)
The greater the frequency, the shorter the wavelength

Question 19

Why hydrogen?

Answer 19

Since hydrogen atoms have only one electron, good element to study
When an electric current is passed through a glass tube that contains hydrogen gas at low pressure, the tube gives off blue light (visible EM radiation [also gives off UV and IR radiation)

Question 20

Hydrogen emission spectrum

Answer 20

Sample has lots of atoms (at first [pre-energy addition], all start at bottom [go at different energy levels—b/c there are a lot—, fall back down to various])
If they could exist anywhere they wanted to, would all be in random places
Anything to n = 1 gives off photon in UV (shorter wavelength, higher energy) b/c diff between 1 and 2 is really big compared to others
Down to n = 3 is infrared
Down to n = 2 is visible (gives off less energy than even shortest one to n = 1 [UV])
Transitions from n = 3 down to n = 2 have a greater ΔE than transitions from n = ∞ down to n = 3

When the visible light is passed through a prism, four narrow bands of bright light are observed from hydrogen (each element has its own atomic emission spectrum which appear as lines in different areas of the EM spectrum)–color you see is mixture of these four

*Always a transition down

Question 21

Line spectrum (discrete lines) produced

Answer 21

A line spectrum is produced:
- Shows the specific energies released as electrons transition down to lower energy levels; has discrete energy values (or colors) at fixed points
- B/c only certain energies are allowed within an atom, there are a limited # of amts of energy (ΔE) that an electron can lose: this means that only certain frequencies can be emitted, leading to a line spectrum
- A continuous spectrum would imply that electrons inside atoms could have any energy

Question 22

Why are the lines converging?

Answer 22

Lines converging as frequency increases b/c energy levels are converging
In every hydrogen, these differences in energy are constant
By studying the frequencies of the lines in the emission spectrum of an element, the energies of the energy levels in atoms can be determined (the presence of sublevels can make the spectra quite complex)

Question 23

What to say

Answer 23

Describe emission:
- It is a (discrete) line spectrum
- The lines result when energy is released from the transition of electrons from higher to lower energy levels
- The lines converge at higher energies
Evidence:
- Gives evidence that electrons reside in discrete energy levels

Question 24

Arrangement of electrons

Answer 24

The most stable, lowest energy levels (shells) are closest to the nucleus
Lower energy levels are filled first w/ electrons
Electronic structure refers to the arrangement of electrons in shells (ex. Ar (18) 2, 8, 8)
Period # = # of shells
Group # = # of valence electrons

Question 25

Energy levels, sublevels, and orbitals

Answer 25

Electron shells are called energy levels and have a quantum # n: they can hold 2n^2 electrons
Each level is further divided into sublevels, which have the names s, p, d, f (each level has as many sublevels as the level #)
Each sublevel is further split into orbitals (regions of space where the electron is most likely to be within the sublevel; have different shapes and hold a maximum of two electrons, which have opposite spin [more of a sign—positive, negative?]), which contain the electrons
Seven energy levels (??)
s has one orbital (max two electrons), p has three (max six electrons), d has five (max ten electrons), and f has seven (max 14 electrons)
Many charged particles, so create a magnetic force (have opposite effects [spins])

Question 26

n = 1 energy level

Answer 26

Closest to nucleus
Contains only one sublevel and orbital
Orbital has spherical symmetry (orbitals of this shape are referred to as “s” orbitals)

Question 27

n = 2 energy level

Answer 27

Has two sublevels (an “s” and a “p” [contain three orbitals, which have a figure eight shape along the three axes x, y, and z])
Since each orbital can hold two electrons, the p sublevel holds a total of six electrons
The p orbitals are slightly higher in energy than the s orbitals b/c of electron repulsion

Question 28

Orbital shapes

Answer 28

W/ the figure eight, chance in the middle is very small (where they don’t tend to be in p orbitals is where they tend to be in s)
Three p orbitals align themselves on [x, y, z] plane perpendicular to each other
Only two in first level b/c repelling each other and close to nucleus (so not a lot of space)
Repelling, so can’t occupy same space as first; want to be as far apart as possible
Moving all over and all in a volume (if you attached a tracer to see, would trace the orbital as it moved)

Question 29

Using the Aufbau Principle to determine electronic config

Answer 29

Electronic configuration can be determined from the Aufbau Principle (means “building up”)
Orbitals w/ the lowest energy are filled first (then work yourself up from there)
In the p orbitals, one electron goes in each orbital before the start to double up
Energy level, sub level, # of electrons in orbital (superscript)
Even if ex. final p orbital empty, put there just to show
Sum of superscripts has to add up to # of electrons

Question 30

n = 3 energy level

Answer 30

Contains three sublevels (s, p, and d)
The d sublevel contains five orbitals w/ complex shapes and can hold a total of ten electrons
The energies of the sublevels in the third level do increase slightly from 3s to 3p to 3d
However there is an added complication: the 3d sublevel is slightly higher in energy than the 4s orbitals in the next level, so the 4s is filled before 3d
3d goes before when 42 is full (only then do you start filling it?)? Grouped together (convention: okay to write is at it’s meant to be)?
Once the 4s is filled, it becomes higher in energy than the 3d (so when 4s subscript is 2, 3d goes before): exceptions include Cr and Cu (4s^1 [last]—this stabilizes the 3d subshell and stabilizes the atom
Turns out, having the d be partially full and having one in 4s is a lower energy state (rather than having two squished in 4s and one empty): same idea for both (and remember to fill each one at a time—being ⅗ full is apparently okay [almost the i])
“Crucial”

Question 31

Interpreting electronic configuration notation

Answer 31

Reads ex. 2py2 (designated according to their axes)
For condensed form, take noble gas in period above
Look at final term (ykwIm): first # is energy level, second (subscript) is valence electron count (can match to periodic table)
We can lump ex. p orbitals up if we have something after (???—I think maybe it’s okay? Ns)
Within four lobes, electrons spread out
W/ p, spread out in the three b/c, again, repelling (put one in each, double up after)—within the same sublevel
B/c of the way they’re laid out (different shape of orbitals causes blue color),
Need to know up to Kr (Z = 36)—last is 4p^6s
s-block is first two groups, d-block is that middle chunk, p-block is the last, and f-block is bottom two periods (ex. Kr in p-block [last sublevel needs to be a p])

Question 32

Positive and negative ions

Answer 32

Electrons are removed from the highest energy level first
In the case of transition metals, the 4s electrons are removed before the 3d (higher level means less energy is required, so that’s the one that’s removed [things don’t switch until you have a 4s]: 4s fills first, buy also removes first, then 3d)
Draw whole thing first and then erase starting always w/ 4s if it’s there (?)
W/ negative ions, electrons are added to the highest energy level (no tricks)

Question 33

The Pauli Exclusion Principle

Answer 33

Only two electrons w/ opposite spin can occupy an atomic orbital

Question 34

Hund’s Rule

Answer 34

Electrons prefer parallel spins in separate orbitals in sublevels; in other words, electrons fill each and all orbitals in the subshell before they pair up w/ opposite spins

Question 35

Building-Up Principle (Aufbau Principle)

Answer 35

Electrons are added by successively filling the subshells w/ electrons closest to the nucleus (lowest energy levels) first

Question 36

Diagrams

Answer 36

Handy diagram to remember how to sort out order in which electrons fill sublevels is to write them all out first,
1s
2s 2p
3s 3p 3d
…
and then draw arrows southwest spearing them
Sometimes, boxes and arrows are used to show the electrons (sublevels should be filled first w/ one electron in each orbital, with one electron in each orbital, with the same spin direction, before adding the second electron to each orbital [starting w/ up is convention: starting w/ down is fine, but the first ones would have to all be going down])
Boxes joined if same sublevel (otherwise, spaced)

Question 37

Note (E)

Answer 37

Electron-electron repulsion = higher energy
W/ data, when multiplying, quote your ans to same # of sig figs as data provided (unless otherwise stated)
Metals formc

Question 38

What is IE?

Answer 38

Ionization energy is the energy needed to remove one mole of electrons from 1 mole of atoms in the gaseous state, forming gaseous ions w/ a single positive charge (endothermic process)

*(n-1)+, n+, -

Question 39

Going down group 2

Answer 39

The radius of the atom increases, so the distance between the nucleus and the outer electron increases
Therefore the force of attraction between the nucleus and outer electron is reduced and so less energy is needed to remove the outer electron
As the # of protons in the nucleus increases going down group 2, the increased shielding and distance have a greater influence of the value of the first ionization energy

Question 40

Going across period 3

Answer 40

General increase across the period
There are more protons in each nucleus, so the nuclear charge in each element increases—therefore the force of attraction between the nucleus and outer electrons is increased
There is a negligible increase in shielding b/c each successive electron enters the same main energy level
So more energy is needed to remove the outer electron

Question 41

Why gaseous?

Answer 41

Metals can be ionized, but only concerns gas

ex. K from a solid to a gas also takes an input of energy (we’d be accounting for this energy as well as that needed to remove an electron—would add separately)
Sometimes not obvious until you look at the electron arrangement (repulsion between paired electrons makes it easier)

Question 42

Successive IE’s

Answer 42

Successive IE’s of atoms refer to the energy needed to remove electrons progressively from the same atom

IE increases as successive electrons are removed b/c net nuclear charge increases, which pulls electrons in closer; a reduction in electron-electron repulsion also leads to the electron being held more strongly

The larger increases in IE are due to electrons being removed from the next energy level down (closer to the nucleus)

Electron is being removed from increasingly positive species

Question 43

Calculating IE w/ emission spectrum

Answer 43

In the emission spectra of hydrogen, the lines converge at higher energies; at the limit of convergence, the lines merge to form a continuum

If the electron in hydrogen is given enough energy to make a transition from the principle quantum # n = 1 to n = ∞, then it can escape the atoms (energy required to do this is called the ionization energy)

Highest energy line basically from n = ∞; magnitude essentially the same (one emitting, one absorbing)

If you know 𝝺 of this, you could find ΔE for one electron falling from just shy of n = ∞ to n = 1 (𝝺 can be measured using equipment w/ emission spectrum)

Question 44

Formula (w/ note)

Answer 44

ΔE = hc/𝝺

NOTE: IE calculated has units of joules (J) and is for 1 electron. To calculate the IE (kJ/mol), we need to divide it by 1000 to convert J to kJ and multiply by NA (6.022 x 10^23 [Avogadro’s 3])

Also,
1 m = 10^9 nm (or 1nm = 10^(-9) m)
*nano = 9

Question 45

Why does it become harder?

Answer 45

More protons; big jump is always after removing when valence has been removed

Question 46

Why log?

Answer 46

When large, easiest way to plot is to take natural log (increment getting bigger, so jump is humongous)

Question 47

Note (lang)

Answer 47

Shortest wavelength = highest energy (from n = ∞)
In UV region = down to n = 1

Question 48

NOTE

Answer 48

Going downwards in the graphs isn’t a thing (would mean that as you removed more electrons, it gets easier and easier)
“What kind of ion would it form normally?”
Less energy required to remove from a higher-energy sublevel
Are there enough electrons?
Only removing one electron each time
“How much energy do I need to put in to make this process happen?”
*Relatively speaking
3rd: K(g)2+ → Li3+(g) + e-
*If we’re removing, had to have been 2+ before
Circle large jumps in IE on graph?
Up to or beyond
On the emission spectrum, you don’t want to push it all away b/c it wouldn’t have anything to fall back to
When demonstrating, probably not best to go to infinity and down (gray area—is it removed or not?) [ykwim]