Chemistry Unit 3: Atomic Theory

electromagnetic radiation

energy that travels in waves

1) ex: microwaves, light, colors, etc.

properties of waves

1) frequency

2) wavelength

3) speed/velocity

wavelength

the distance between crest₁ and crest₂ (or trough₁ and trough₂)

1) symbol = λ

crest

the highest point of a wave

trough

the lowest point of a wave

frequency

the number of waves that pass a specific point per period fo time

1) symbol = ν (Nu)

speed of light

the speed of all electromagnetic radiation

1) = 3.00 × 10⁸ m/s

2) symbol = c

wave equation

c = λv

1) shows relationship between the speed of light, wavelength, and frequency

electromagnetic spectrum

= arrangement of different types of electromagnetic radiation from longest to shortest wavelength

1) shows that as wavelength increases, frequency decreases

2) radio waves → microwave → infrared → ultraviolet → x-rays → Gamma Rays

3) might wanna get a visible representation for the 10^power markers

Relationship between wavelength and frequency

as one goes up, the other goes down

1) = an inverse relationship

Relationship between frequency and energy

as frequency goes up, energy goes up with it (meaning they both are inversely related to wavelength)

radio waves

used for tv, cellular phones, radio

infrared radiation (IR)

the heat we feel from sunlight and infrared lamps in restaurants

1) also is in remotes

visible light

700-400nm wavelengths

1) = the only radiation our eyes can detect

2) violet = the shortest wavelength of visible light

ultraviolet light (UV)

can cause sunburn in sunlight

x-rays

can pass through soft substances to allow us to see images of the bones and teeth in body

continuous spectrum

spectrum containing all wavelengths of light in a specific range

1) ex: white light after passing through a prism

photons

packets of energy emitted by heated atoms

1) travel at the speed of light

line spectrum

a discontinuous series of wavelengths

atomic spectrums

line spectrums produced by atoms when heated

1) each atom has a unique atomic spectrum

2) = identification factor

light

a form of electromagnetic energy

spectrum

the entire range of electromagnetic energy

wave energy

measured in Joules

energy equation

E = (h)(ν)

1) h = Planck’s constant

Planck’s Constant

the number multiplied by the frequency to give you the energy

1) = 6.626 × 10^-34 Js

2) Js cancels out frequency’s 1/s units

violet

the visible light wave with the smallest wavelength

red

the visible light wave with the largest wavelength

electron light emission explanation

electrons absorb heat, get excited, then jump away from the nucleus

1) coming back down releases energy in the form of light

atomic spectrum (explained)

= light emissions from excited electrons returning to lower levels

Bohr’s discovery

used line spectrum of hydrogen to infer that electrons = organized into energy levels

1) discovered that lines on spectrum correspond with quantized energy changes

2) proposed idea of designated “orbits”

quantized energy meaning

only certain values of energy are permitted

1) = like an elevator; can’t just stop in between

Rutherford vs Bohr

1) Rutherford likely expected to see rainbow spectrum because lacked concept of quantized energy levels

2) Bohr expected to see line spectrum because he learned that you can’t just jump in between

Shrodinger

maker of the Electron Cloud Model

Electron Cloud Model

1) electrons are located in orbitals (not orbits)

2) allowed for atomic probability maps

atomic probability maps

show the likeliness of electrons to be located in a region of an atom

1) areas with high density = most likely to have an electron

Atomic Spectra Series

the quantized energy level transitions that give off light energy

1) Lyman Series

2) Balmer Series

3) Paschen Series

Balmer Series

quantized energy level transitions that give off visible light

distance of higher electron levels

higher electron levels = closer together

excited state

when an electron moves to a higher energy level after an atom absorbs energy

1) falls down after, emitting light energy

2) greater the distance of falling, the more energy is needed/+ released

principle energy levels

the quantized energy levels of an atom

1) = the “n-value” (n = 1,2,3,4…)

effect of n on electron energy (& vice versa)

lower n → lower energy

1) can’t excite themselves and move away

2) = Nate

higher n → higher energy

1) clearly had the energy to move away; more energy is required to keep them relevant within an atom

2) = Nicholas

sub-levels

principle energy levels = divided into sub-levels

principle energy sublevels (names)

1) s

2) p

3) d

4) f

sub-levels and orbitals

each sub-level contains orbitals

orbitals

the 3D regions in which an electron is most likely (90%) to be located

1) = probability maps essentially

# of orbitals per principal energy level

n²

maximum # of electrons per orbital

2

1) each electron is differentiated by having an opposite spin (clockwise or counter clockwise)

s

the lowest starting sub-level

1) has a single orbital

2) in shape of sphere

p

the second lowest starting sub-level

1) has 3 orbitals

2) = in shape of dumbbell

d

second highest sub-level

1) has 5 orbitals

2) = in shape of clover

f

highest sub-level

1) has 7 orbitals

electron capacity for each principal energy level

2n²

3 Principals of Electrons

1) Aufbau Principle

2) Pauli Exclusion Principle

3) Hund’s Rule

ways of observing electron location & movement

1) emission spectroscopy*

2) ionization energy

emission spectroscopy

forcing electrons into higher energy states and observing what happens when they return

1) = a way to observe electrons

ionization energy observation method

measuring the amount of energy required to create a positive ion (removing electrons from the atom)

ionization energy

the minimum amount of energy required to remove one electron from an atom in its gaseous phase

1) = measured in KJ/mol

2) increases as an atom loses electrons (because p to e ratio increases, increasing pull and therefore increasing energy needed to pull it away)

STUDY

LABS

atomic radius

the distance from the nucleus to a valence electron

atomic radius trends

increases ← and down (round the corner)

1) ← because more electrons on the same ring scrunches atom in with positive charge

2) down because more rings = more size

ionic radius

increases ← and down (round the corner)

1) ← because in ion form, entire period has the same number of electrons; so fewer protons means less scrunch

2) down because everyone either gained or lost the same amount; now just who started with more

ionic radius for the same element

the more electrons/negative charge, the larger the ionic radius

1) electrons outnumbering protons allows them to spread out more on a single ring, increasing radius

Ionization energy trends (periodic table)

increases going up and to the right → (turn out of Clayton)

1) increases going up because fewer rings means tighter pull

2) increases going to the right because more electrons on same ring means tighter pull

ionization energy irregulars

2A > 3A and 5A > 6A

1) happens because if you look at electron composition