Atomic Theory Notes

I. Development of the AtomThere were many scientists that contributed to discovering and determining the internal structure of an atom.

1) John DaltonProposed his atomic theory based on ideas of particles in matter at the time

- Atoms are a neutral piece of mattero All elements are composed of tiny particles called atomso Atoms cannot be broken down furthero All atoms of the same element have identical properties,

atoms of different elements have different propertieso Atoms combine in whole number ratios to form compounds

2) J.J. ThomsonDiscovery of charged particles with his “cathode ray experiment”

- Cathode ray was deflected by a magnet- Cathode ray was attracted to a positive charge and repelled by a negative charge

o Opposite charges attract, so cathode ray was negatively chargedo Called these negative charges “electrons”

- Changed the model of an atom to be a ball of positively charged matter with negative electrons embedded (stuck) throughout

3) Ernest RutherfordGold Foil Experiment led to two major discoveries

- Beam of alpha particles (radioactive particles) towards a piece of gold foil- Expected alpha particles to pass through and hit a fluorescent screen on opposite side

o Most of particles passed through in a straight path, but….o Some particles bounced off to the side or backwards

Two major conclusions:a) Atom is mostly empty space (where the negative electrons travel)b) Small, dense region at the atom’s center called a nucleus, where

positive protons are found

4) Neils BohrWhen studying light, proposed that electrons are placed around the nucleus in different levels, called energy levels

- When an electron gains energy, it moves to a higher energy level

- When the electron drops back down to original level, it releases the same amount of energy in the form of light

o The different levels electrons can drop back to produce different colors of light (ROYGBIV)

- His model is referred to as the planetary model, because the electrons travel in circular paths (orbit) around the nucleus in their designated, original energy levels

5) QuantumThis model describes the probable location of electrons around the nucleus because it is mathematically and scientifically impossible to predict when an electron is at a given time

- The probable location is a region of space around the nucleus- Electrons can be found in any part of this region, which are known as electron clouds

o Electrons are more likely to be found closer to the nucleus due to their attractions to the positive charge there

o Less likely to find electrons the farther away from the nucleus you are

II. Subatomic Particles1) Identity of the subatomic particles

The atom is made up of three main particles: protons, electrons and neutrons.a) Protons (p+) – positive particles found in the nucleus of an atom

∑ The number of protons is the identifying factor of an element∑ Have an approximate mass of 1 amu∑ Number of protons is the nuclear charge of the atom

b) Neutrons (n0) – neutral particles found in the nucleus of an atom∑ Have an approximate mass of 1 amu

c) Electrons (e-) – negative particles found in the empty space around the nucleus∑ In a neutral atom, the number of protons = the number of electrons∑ Have an approximate mass of 0 amu

Subatomic particle

Proton Electron Neutron

Symbol p+ e- n0

Location Inside the nucleus

Outside the nucleus

Inside the nucleus

Mass 1 amu 0 amu 1 amu

2) Determining subatomic particles in a NEUTRAL atomThe periodic table is necessary to determine the number of protons, neutrons and electrons

a) Protons = the atomic number on the periodic table

b) Electrons = add up the numbers in the electron configuration - Will be the same as the number of protons (atomic number)

c) Neutrons = mass number – atomic number = # of neutronso Mass number is the protons + neutrons combinedo Mass number the atomic mass rounded to a whole number

3) Determining subatomic particles in an IONThe periodic table is necessary to determine the number of protons, neutrons and electrons in an ion. However, an ION is when an atom has lost or gained electrons

- Atoms can lose or gain electrons in chemical bonding

a) Protons = the atomic number on the periodic table

b) Neutrons = mass number – atomic number = # of neutronso Mass number is the protons + neutrons combinedo Mass number the atomic mass rounded to a whole number

c) ElectronsIn a neutral atom, the # of protons = the # of electrons

If an atom has lost electrons, it will have more protons than electrons- the extra protons will give the ion a positive charge = known as a cation

ex) Ca vs. Ca+2

p+ = 20 p+ = 20n0 = 20 n0 = 20e- = 20 e- = 1820p+ = 20e- = 0 charge 20p+ ≥ 18e- = +2 charge

If an atom has gained electrons, it will have more electrons than protons- the extra electrons will give the ion a negative charge = known as a anion

ex) O vs. O-2

p+ = 8 p+ = 8n0 = 8 n0 = 8e- = 8 e- = 108p+ = 8e- = 0 charge 8p+ ≤ 10e- = - 2 charge

III. IsotopesIsotopes are atoms of the same element with the same number of protons and electrons, but a different number of neutrons (and therefore mass number)

Ex) 16O817O8

18O8 Ex) 1H12H1

3H1

p+=8 p+=8 p+=8 p+=1 p+=1 p+=1e-=8 e-=8 e-=8 e-=1 e-=1 e-=1n0=8 n0=9 n0=10 n0=0 n0=1 n0=292% 1% 1% 99% 0.6% 0.4%

Each of these isotopes makes up certain percent abundance amount out of 100% here on earth These percent abundances are looked up and cannot be changed.

ALL elements on the periodic table have isotopes, which each have different percent abundances. We can use this information to determine the weighted average atomic mass (WAAM) for each natural element.

WAAM = (mass of isotope #1 x % abundance) + (mass of isotope #2 x % abundance) + …100 100

*This formula MUST be memorized and you must know how to use it to determine the WAAM of an element** answer should reflect the most abundant isotope and should be in the range of the original isotope masses!!

Sample Problems:1) There are two isotopes of Fe. Fe-56 makes up 90.3% of all Fe and has a mass of 55.996 grams,

and Fe-58, which makes up 9.7% of all Fe and has a mass of 57.938 grams. What is the weighted average atomic mass?

WAAM = (55.996 g x 90.3%) + (57.938 g x 9.7 %) = 50.564 g + 5.62 g = 56.184 grams100 100

Answer falls between 56 and 57WAAM is closer to 56 because there is a greater percent abundance of that isotope!

2) There are three isotopes of Ne. Ne-20 makes up 90.5% of all Ne and has a mass of 19.992 grams, Ne-21, which makes up 0.27% of all Ne and has a mass of 20.993 grams, and Ne-21, which makes up 9.23% of all Ne and has a mass of 21.991 grams. What is the weighted average atomic mass?

WAAM = (19.992 g x 90.5%) + (20.993 g x 0.27 %) + (21.991 g x 9.23 %)100 100 100

= 18.093 g + 0.0567 g + 2.0298 g = 20.1765 grams

Answer falls between 20 and 22WAAM is closer to 20 because there is a greater percent abundance of that isotope!

IV. The Bohr Model1) Contribution to the Model of the Atom

Protons and neutrons are in the nucleus- Electrons are placed in orbits or energy levels around the nucleus

* Energy level (n) = fixed pathway (orbit) around the nucleus, each with a designated amount of energy assigned to it

maximum number ofelectrons in each energylevel = 2n2, where n is the energy level

** lowest-energy energy level is the one closest to the nucleus** highest-energy energy level is the one farthest from the nucleus

When electrons are placed around the nucleus:a) Begin filling lower energy levels first (those closest to the nucleus)b) Must completely fill lowest energy level before beginning next energy level

2) Electron ConfigurationsAn electron configuration shows the number of electrons that reside in each energy level around the

nucleus. The electron configurations on the periodic table are written in ground stateGround State: when electrons are in the lowest energy configurations possible

- Shows number of electrons in each energy level

Valence Electrons (v.e.-): the last number in the electron configuration; used to determine chemical properties and for chemical bonding

ex) Carbon (C) ex) Phosphorus (P) ex) Copper (Cu)2 – 4 2 – 8 – 5 2 – 8 – 18 – 1

* has 4 valence electrons * has 5 valence electrons * has 1 valence electron

V. Ground State vs. Excited State1. Ground State (GS)

- stable electron arrangement around the nucleus- electrons are in their lowest possible energy locations in their designated energy levels- all electron configurations given on the periodic table are in ground state

ex) Al = 2 – 8 – 2 (13e- total)Ca = 2 – 8 – 8 – 2 (20e- total)Kr = 2 – 8 – 18 – 8 (36e- total)

2. Excited State (ES)If energy is added to an atom (light, electrical, heat), electrons in ground state will absorb that energy and jump to higher energy levels

- temporary moment at which an electron has absorbed energy and has jumped up to a higher (farther from the nucleus) unfilled energy level = this is known as EXCITED state

- unstable state- the electron configuration in excited state will NOT match the configuration on the Periodic

Table and will not follow the rule of 2 – 8 – 18 – 32 ….

ex) What element is: 2 – 1 – 1 – 1 ex) What element is: 2 – 6 – 2 – 2

- add up total # of e- = 5 - add up total # of e- = 12- determine what element on PT - determine what element on PT

has 5e- = Boron (B) has 12e- = Magnesium (Mg)- GS of Boron is 2 – 3 - GS of Mg is 2 – 8 – 2

so this is ES of Boron so this is ES of Magnesium

ex) Provide an excited state configuration of Phosphorus (P).

- Ground State is 2 – 8 – 5, so any configuration that has 15e- but doesn’t match this configuration in GS will work

o Just cannot overfill an energy level (follow rules of max electrons for each)ß 2 – 7 – 6ß 1 – 8 – 6 ß 2 – 8 – 4 – 1 ß Possibilities are endless!

3. Movement between Ground State and Excited StateWhen energy is added to an atom (light, electrical, heat), electrons in ground state will absorb that energy and jump to higher energy levels [Excited State]

- Being in excited state is unstable, and electrons will return back to their original energy levels

o When electrons return back to original levels [Ground State], the same amount of energy that was initially absorbed is now released

o This energy is released as light particles, and the colors ROYGBIV are seenß Color emitted by an electron depends on what two energy levels it

moved between

a) Energy comes in and is absorbed by an electron. c) Electrons don’t stay in ES for very long.b) Electron jumps from ground state to a higher d) Electron returns to lower energy level (GS)

energy level (excited state) and emits energy as a photon of light

VI. Visible Light and Waves1) Visible Light

The color produced by an electron depends on the distance the electron moves between fromexcited state back to ground state.

- Red: lowest energy drop (small distance between excited state energy level and ground state energy level)

- Orange

- Yellow IncreasingEnergy

- Green distance electronfalls back down

- Blue to Ground State

- Indigo

- Violet highest energy drop (large distance between excited state energy level and ground state energy level)

To the naked eye, the light photons given off by electrons returning to ground state is seen as a one of the colors of the continuous (visible light) spectrum (ROYGBIV).

However, we can take an instrument called a SPECTROSCOPE, which is basically a PRISM, and use it to scatter the light being produced. The pattern that will be seen is then called BRIGHT LINE SPECTRUM. These patterns of spectral lines are UNIQUE for each element (just like a human fingerprint is unique to each person).

- Much research has been done to identify all of the spectral band patterns for each unique element, and bright line spectrum have been produced for these elements, indicating at what wavelength each band of color is seen.

- Therefore, if given an unknown sample, we can use spectral lines to identify which element (or elements) we have (each element gives off its own characteristic bright line spectra.)

Answer: Gas A and Gas D (all the spectral lines match up)

Li gives off red light Na gives off yellow light

K gives off blue light Rb gives off purple light

2) Waves and Visible LightVisible light waves are the only electromagnetic waves we can see. We see these waves as the colors of the rainbow. Each color has a different wavelength. Red has the longest wavelength and violet has the shortest wavelength. When all the waves are seen together, they make white light.

R O Y G B I V

A) Properties of Waves

Wavelength (λ) = the distance between two neighboring peakso Units are meters (m) or nanometers (nm)o 1 nm = 10-9 m

Frequency (f) = the number of complete wavelengths that pass a given point each secondo Unit is 1/second or s-1

= Low energy wave,larger wavelength

= High energy wave,smaller wavelength

B) Calculation of LightFormula 1 Formula 2E = hf c = λf

E = energy (Joules) c = speed of lighth = Planck’s Constant (3.0 x 108 m/s)

(6.626 x 10-34 J∙s) λ = wavelength (m or nm)f = frequency (1/s) f = frequency (1/s)

Joules = (J∙s)(1) m = (m)(1)S s s

Example Problems:

1. Calculate the wavelength in m of the electromagnetic radiation whose frequency is 7.5x1014s-1.

c = λf 3.0 x 108 m/s = λ(7.5x1014s-1) 3.0 x 108 m/s = λ(7.5x1014s-1) λ = 4 x 10-7 m 7.5x1014s-1 7.5x1014s-1 or

400 nm

2. Determine the frequency of light with a wavelength of 4.257 x 10-7 cm.

c = λf Unit Conversion for λ: 4.257 x 10-7 cm x 10-2 m = 4.257 x 10-9m1 1 cm

3.0 x 108 m/s = (4.257 x 10-9m)f f = 7.05 x 1016 1/s or s-1

4.257 x 10-9m 4.257 x 10-9m

3. Determine the energy, in Joules, of a photon whose frequency is 3.55 x 1017 s-1.

E = hf E = (6.626 x 10-34 J∙s)(3.55 x 1017 s-1) E = 2.35 x 10-16 J

4. What is the frequency of a radio wave with an energy of 1.55 x 10-24 J?

E = hf 1.55 x 10-24 J = (6.626 x 10-34 J∙s)f f = 2.34 x 109 1/s6.626 x 10-34 J∙s 6.626 x 10-34 J∙s

5. What is the energy of a photon of ultraviolet light with a wavelength of 4.25 x 10-8 m?

Part 1: c = λf 3.0 x 108 m/s = (4.25 x 10-8m)f f = 0.706 x 1016 1/s

4.25 x 10-8m 4.25 x 10-8m

Part 2: E = hf E = (6.626 x 1034 J∙s)(0.706 x 1016 1/s) E = 4.68 x 10-18 J

VII. Quantum (Wave-Mechanical) Model- Estimates the probability of finding an electron in certain regions of space called electron clouds

The location of electrons in these “clouds” are described by:

A. Energy Level (n) : 1, 2, 3, and 4

B. Sublevel: s, p, d and f

C. Orbitals:

s p d f

* each orbital can hold a maximum of 2e-

D. Electron Spin: Up ↑ or down ↓

A) ENERGY LEVELS (n)- the first number describes the MAJOR ENERGY LEVEL of the electron and is called the

principle quantum number - this is the same as the number of the energy level that contains the electrons

o higher the “n” the more energyo n = 1, 2, 3, 4, etc… 7

B) SUBLEVELS - Each energy level has one or more sublevels associated with it- Each principle energy level contains as many sublevels as the number of the level

Ex) n = 1 has 1 sublevel n = 2 has 2 sublevels n = 3 has 3 sublevels n = 4 has 4 sublevels

The first (1st) sublevel is designated the s sublevel

The second (2nd) sublevel is designated the p sublevel

The third (3rd) sublevel is designated the d sublevel these twoshapes are

The fourth (4th) sublevel is designated the f sublevel very complex

* Sublevels are described by using the principal energy level together with the letter designation of each sublevel- 1s - 2s 2p - 3s 3p 3d - 4s 4p 4d 4f

C) ORBITALS- an orbital is a location inside the atom where an electron is most likely to be found

- Orbitals are all the same size; they are represented as boxes and can fit up to two e- in them

“s” orbital – electrons are found in 1 orbital with spherical shapes

“p” orbitals- electrons are found in 3 orbitals with dumbbell shapes

“d” orbitals - electrons are found in 5 orbitals with dumbbell shapes

“f” orbitals - electrons are found in 7 orbitals with dumbbell shapes

D) SPIN OF AN ELECTRON - This number indicates that each orbital can contain two electrons spinning in opposite directions- Spin is a quantum mechanical property of electrons and may be thought of as up spin or down spin.

o Up Spin ____ ** Each orbital can only contain two electrons spinning in the

opposite direction

o Down Spin ____

In an atom….

Energy Level Sublevel(s) Orbital(s) # of e-

1 s 1s 1s2 2e-

2 s 2s2

2s8e-

p 2p 2p6

3 s 3s2

3s

p 3p 3p6 18e-

d 3d 3d10

·

‚

Three rules when filling sublevels and orbitals:1) Aufbau principle: electrons enter orbitals of lowest energy first2) Pauli Exclusion Principle: An atomic orbital can contain a maximum of 2 electrons3) Hund’s rule: When electrons occupy orbitals, one electron enters each orbital with the same spin

until all orbitals contain one electrons, then pairs can be made with an electron of opposite spin

Examples:

Atomic # Bohr Notation Sublevel Notation Orbital Notation

3 2 – 1 1s22s1 ↑↓ ↑_1s 2s

8 2 – 6 1s22s22p4 ↑↓ ↑↓ ↑↓ ↑_ ↑_1s 2s 2p

11 2 – 8 – 1 1s22s22p63s1 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_1s 2s 2p 3s

17 2 – 8 – 7 1s22s22p63s23p5 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_1s 2s 2p 3s 3p

20 2 – 8 – 8 – 2 1s22s22p63s23p64s2 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓1s 2s 2p 3s 3p 4s

29 2 – 8 – 18 – 1 1s22s22p63s23p64s13d10

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑↓ ↑↓ ↑↓ ↑↓1s 2s 2p 3s 3p 4s 3d10