1. Structure of Matter

2. Important Discoveries About the Atom
Law of Conservation of Matter
In 1774, Antoine Lavoisier performed experiments & measurements that led to this law.
In chemical reactions, matter cannot be created nor destroyed
Law of Constant Composition
In 1700, Joseph Proust mad measurements on chemical reactions and compounds to develop this law
Each pure chemical compound always has the same percentage composition of each element by mass.
These laws led John Dalton to develop his theory.

3. Important Discoveries About the Atom
Dalton’s Atomic Theory states that:
All matter is composed to tiny, indivisible particles, called atoms, that cannot be created or destroyed.
Each element has atoms that are identical to each other in all of their properties, and these properties are different from the properties of all other atoms.
Chemical reactions are simple rearrangements of atoms from one combination to another in small whole-number ratios.

4. Important Discoveries About the Atom
Dalton’s atomic theory led him to propose the …
Law of Multiple Proportions
When 2 elements can be combined to made 2 different compounds, and if samples of these 2 compounds are taken so that the masses of one of the elements in the 2 compounds are the same in both samples, then the ratio of the masses of the other element in these compounds will be a ratio of small whole numbers.

5. Law of Multiple Proportions

6. Important Discoveries About the Atom
In 1834, Michael Faraday showed that an electric current could cause chemical reactions to occur, demonstrating the electric nature of the elements.
In the 1870s the cathode ray tube was developed by Sir William Crookes.
He mistakenly thought that the cathode rays were negatively charge molecules instead of electrons.

7. Important Discoveries About the Atom
In 1897, J.J. Thomson determined that cathode rays were a fundamental part of matter he called electrons.
Thomson also discovered the electron’s charge to mass ratio, (e/m = × 108 coulombs per gram) by measuring the deflection of the cathode rays in the presence of electric and magnetic fields.

8. J.J. Thomson

9. Important Discoveries About the Atom
Oil Drop Experiment
In 1909, Robert Millikan performed this experiment. He determine the charge of an electron to be × coulomb.
Using Thomson’s charge to mass ratio, he then calculated the mass of an electron= × 10-28g.

11. Important Discoveries About the Atom
From those experiments, the plum pudding model was developed, which has electrons swimming in a sea of positive charges.

12. Important Discoveries About the Atom
Gold Foil Experiment
Performed by Ernest Rutherford in 1910.
He was interested in radioactive materials, alpha & beta particles had already been discovered.
Rutherford’s Nuclear Model

13. Important Discoveries About the Atom
In 1919, Rutherford also discovered the proton with a mass of 1.67×10-24g.
It’s 1836 times bigger than an electron
His student, James Chadwick, discovered the nucleus in It has almost the same mass as a proton.

14. Subatomic Particles Name Symbol Absolute charge (coulombs)
Absolute mass (g)
Relative charge
Relative mass
Electron
e or e-
-1.602×10-19
9.109×10-28 -1
5.486×10-4
Proton
p or p+
+1.602×10-19
1.673×10-24 +1
1.0073
Neutron
n or n0
1.675×10-24
1.0087

15. Important Discoveries About the Atom
While all of this stuff was happening, other physicists were interested in the interaction between light and matter. (mid-1800s)
One thing they discovered was that each element, when heated or sparked with electricity, gives off characteristic colors.
A spectroscope was used to show these colors consist of discrete wavelengths of light (line spectra), not the uniform rainbow.
The line spectra for most elements & compounds can be complex but hydrogen’s is relatively simple.

16. Important Discoveries About the Atom

17. Important Discoveries About the Atom
In 1885, Johann Balmer found a mathematical relationship between the wavelengths of the lines in the visible light region of the spectrum.
Similar lines were discovered in the infrared (Paschen series) and ultraviolet (Lyman series) regions.
Johannes Rydberg extended Balmer’s equations so that all of the wavelengths could be predicted.

18. Important Discoveries About the Atom
Planetary or Solar System Model
Proposed by Niels Bohr in He assumed that electrons move around the nucleus in only certain circular orbits.
Max Planck described light as packets or quanta of energy called photons.

19. Important Discoveries About the Atom
In 1924, Louis de Broglie suggested that if light can be considered as particles, then the small particles (electrons) may also have the characteristics of waves!
In 1927, Erwin Schrödinger applied mathematical equations for waves to electrons in the atom and began the wave-mechanical model of the atom.
For the hydrogen atom, the results are very similar to Bohr’s model of the atom except the electron doesn’t follow a precise orbit. The position of the electron is described as a probable location.
Werner Heisenberg developed the uncertainty principle in the 1920s that says the position & momentum of any particle cannot be both known exactly at the same time. As you know one more precisely, the other becomes less certain.

20. Atomic Structure
An atom usually exists in the lowest possible energy state – called the ground state.
An atom that has more energy than the ground state is said to be in an excited state.
When at atom loses energy in going from an excited state to the ground state, that energy is emitted as light.

21. Atomic Structure

22. Atomic Structure Wavelength, Frequency, & Energy of Light
All EM radiation may be considered as waves defined by wavelength (λ) & frequency (ν).
Wavelength – distance between 2 repeating points on a sine wave.
Frequency – the # of waves that pass a point in space each second

23. Atomic Structure
Wavelength and frequency are inversely proportional to each other.
(wavelength)(frequency) = speed of light
λν = c
c = 3.00×108m/s
λ= meters
ν = s-1 or 1/s

24. Atomic Structure
Max Planck discovered that the energy of the EM waves is proportional to the frequency & inversely proportional to the wavelength
hν = E Planck’s constant (h)
h = E h = 6.63×10-34Js λ

25. Atomic Structure
The Bohr Model- requires the electrons in the atom be confined to specific, allowed orbits. Using physics, the energy of an orbit with the number, n, is
En = -2πme4 = ×10-18joule
n2h2 n2
Where m = mass e-, e = charge on e-, h = Planck’s constant, & n = principal quantum number
n also represents the number of each orbit in Bohr’s model, starting with the one closet to the nucleus

26. Atomic Structure Bohr’s orbits
Energy, in the form of light, is emitted from an atom when an electron moves from an orbit to a lower-numbered orbit.
When an electron is promoted to a higher numbered orbit, energy must be added.
The energy difference between 2 orbits is constant so you know how much energy is being released when the electron drops down to its original orbit. You use the equation in the last slide to calculate the energy difference.

27. Atomic Structure
Hydrogen
Emitting light
Another representation

28. Atomic Structure
Bohr also thought that the momentum (mass × velocity) of the electron be related to the size of the electron’s orbit.
Mv = nh
2πr
For hydrogen (n = 1), the radius of the electron was calculated to be 53pm, called the Bohr radius. The radii of other orbits are all whole-number multiples of it.
Gave chemists a theoretical value for the size of a hydrogen atom.
Bohr won a Noble Prize for all this. Go Bohr!

29. Atomic Structure Wave-Mechanical Model of the Atom
After Bohr’s achievement, Louis deBroglie ascertained that electrons could also act as waves, not just particles.
1 way of looking at this duality, is to look at the equation for the energy of a wave…
E = hv = h = mc2 λ
Rearranging this equation, you can see the relationship between the mass of the electron and the wave (frequency) … = h = mc2

30. Atomic Structure
Describing the motion of an electron, however, requires complex “wave equations” and higher level calculus (differential equations) than is taught in high school. But understanding the results of these wave equations can be done:
Wave equations require 3 numbers, called quantum numbers, in order to reach a solution.
The principal quantum number, n
The azimuthal quantum number, l
The magnetic quantum number, ml
And a 4th, unique, quantum number, the spin quantum number, ms
There are specific rules for applying these numbers to electrons but it isn’t required on the AP test.

31. Atomic Structure
The wave equations changed the picture of the atom completely. The fixed orbits of the Bohr model are replaced with a cloud of electrons around the nucleus. The modern orbital is the region of space in which there is the greatest (90%) probability of finding an electron.
Probability plot with 90% of electrons within circle
Bohr orbit as fixed rings
Probability plot of electrons in the 1st orbit

32. Atomic Structure
The circular orbits of the Bohr theory are replaced with spherical electron clouds. The wave equations have shown that most electron clouds have shapes that are more complex than Bohr’s orbits but are still simple geometric shapes.
The arrangement of electrons deduced from the wave equations agrees well with the periodic table. Many physical & chemical properties of elements & compounds are more fully understood with knowledge gained about the electronic structure & orbital shape.

33. Atomic Structure
The results of the wave equations agree completely with the Bohr model. Specifically, the energy change for an electron moving from one electron cloud to another. Also the 53pm radius found for the electron in the hydrogen atom still holds true.
The Heisenberg uncertainty principle is fundamental to the wave mechanical model. The exact location and momentum of an electron cannot be known at the same time.

34. Atomic Structure Principal energy levels (shells)
Another term for principal quantum number
There are a maximum of 7 shells in an atom.
1st shell is closest to the nucleus, the 7th is farthest away.
As they become larger, the further away from the nucleus, the more electrons they can hold.

35. Atomic Structure Sublevels (subshells)
Each principal energy level contains one or more sublevels, or the azimuthal quantum number, l.
l can never be greater than n-1.
How many sublevels can each principal energy level hold?
The same number as the value of n for that energy level. Ex. n = 3? It has 3 sublevels. n = 6? It has 6 sublevels. However, for the 118 known elements, only 4 of the sublevels are used.

36. Sublevels Principal level, n Sublevel number, l Sublevel letter 1 s 2s 2
0, 1
s, p 3
0, 1, 2
s, p, d 4
0, 1, 2, 3
s, p, d, f 5* 6*
0, 1 ,2
s, p ,d 7*
*Only sublevels used by known elements are shown here.

37. Atomic Structure Orbitals
Each sublevel may contain one or more electron orbitals, a region of space that has a high electron density.
Each orbital may hold a maximum of 2 electrons.
In order to share the orbital, the electrons must have opposite spins.
The number of orbitals a sublevel can have depends on its azimuthal quantum number, l, and is equal to 2l +1.

38. Orbitals Sublevel number, l Sublevel letter Number of orbitals 2l + 1
Number of electrons per Sublevel s 1 2 p 3 6 d 5 10 f 7 14
Each orbital is given a magnetic quantum number, ml. Values range from –l to +l.
orbital
ml. values s p
-1, 0, +1 d
-2, -1, 0, +1, +2 f
-3, -2, -1, 0, +1, +2, +3

39. Orbital shapes

40. Electronic Structure of the Atoms p d f 7s 7p 6s 6p 6d 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s
start
Orbital filling diagram

41. Electron Configurations
Some rules to abide by:
Aufbau Principle – electrons fill the lowest available energy level before moving to a higher one.
Some exceptions:
Element
Electron Configuration
Copper, Cu
1s22s22p63s23p64s13d10
Silver, Ag
1s22s22p63s23p64s23d104p65s14d10
Gold, Au
1s22s22p63s23p64s23d104p65s24d105p66s14f145d10
Chromium, Cr
1s22s22p63s23p64s13d5
Molybdenum. Mo
1s22s22p63s23p64s23d104p65s14d5

42. Abbreviated Electron Configurations- Use noble gases1 2 3 4 5 6 7 s p d f
Example
Fe = 1s22s22p63s23p64s23d6
Fe = [Ar]4s23d6

43. Valence Electrons
Many times, chemists are only interested in the outermost electrons in an atom, the valence electrons.
Only s and p electrons are valence electrons.

44. Orbital Diagrams Hund’s Rule
p, d, or f orbitals in a sublevel must all be filled with one electron each before a 2nd electron is allowed to pair in any orbital.

46. Quantum Numbers Principal quantum number, n = 1-7
Azimuthal quantum number, l = n-1 = 0-6
Magnetic quantum number ml = -l -0-+l
Spin quantum number, ms = + ½ or – ½
Lowest possible values for each quantum number are used 1st.
Pauli exclusion principle- no 2 electrons can have the exact same quantum number
While writing the quantum numbers is not on the AP test, knowing what they look like according to the rules is.
See handout.

47. Quantum numbers, what do they mean?
Principal quantum number, n – represents the average distance of the electron from the nucleus, or the size of the principal energy level.
Azimuthal quantum number, l – represents the shape(s) of the orbitals within the sublevel.
Magnetic quantum number ml – represents the oreientation of each orbital in space on an x, y, z axis.
Spin quantum number, ms – represents the “spin” of the electrons.