1. Elements, atoms, & the discovery of atomic structure

2. Early models of the atom
Chapter 4
Early models of the atom

3. What is an atom?
The smallest particle of an element that can keep the same properties of the element.
What is the smallest number of carbon atoms that have all the properties of a carbon atom?
One

4. Democritus (Ancient Greece, 440 B.C.)
Democritus thought atoms to be indivisible and indestructible.
He had no experimental evidence for this claim.
He stated that atoms are the smallest particles of all matter.
Who was the first person to use the word atom?
Democritus

5. Atomic Structure- outline
Dalton’s Atomic Theory
Thomson’s model vs. Rutherford’s model of the atom
Discovery of subatomic particles:
Protons, neutrons, and electrons
Isotopes
Quantum mechanical model

6. Early model of the atom John dalton (1766-1844)
English chemist and school teacher
Came up with a Theory in 1803 based on experimental evidence.
Dalton’s Model was a dense solid sphere; indivisible and unchanged in chemical reactions.
Where did he get his ideas?

7. Berzelius’s Experiment Law of Definite Proportions
aka- law of constant composition: a compound always contains elements in definite proportions.

8. Proust’s experiments Law of Definite Proportions
aka- law of constant composition

9. Dalton’s Atomic Theory (pg 103)
All matter is made of indivisible atoms; they can be neither created nor destroyed during chemical reactions
All atoms of a given element are identical in their physical and chemical properties; they differ from atoms of every other element
Atoms of different elements combine in simple whole-number ratios to form compounds (can form more than one compound together)
Chemical reactions consist of the combination, separation, or rearrangement of atoms
Which of these are no longer valid?

10. Can we see an atom? We can now view individual atoms!
Scanning electron microscope (SEM)

11. J.J. Thomson’s Experiments 1897
Used a cathode ray tube: metal is placed at the positive end (anode) and the negative end (cathode).

12. J.J. Thomson’s Experiments
Rays produced from the cathode end.
The beam bends toward positive plates.
A small paddle wheel spins when hit by the cathode rays.

13. What happened to the Dalton model?
Conclusions-
Particles were bent by the charged plates
particles are charged.
Particles set the wheel in motion
particles have mass.
Particles were the same size no matter what metal was used
particles are the same, no matter the element.
Atoms are neutral, but are made of negatively charged particles
atoms are made of negative and positive charge.
What happened to the Dalton model?

14. The Plum Pudding Model
An early and now obsolete attempt to describe the interior structure of atoms
Electrons scattered throughout positively charged matter
Electrically neutral
electron
sphere of positive charge

15. Rutherford – Thomson’s student

16. Rutherford's gold foil experiment
Alpha particles (positively charged) bombarded foil of various metals.
A fluorescent screen was placed around to detect the particles as they passed through the metal.

17. Rutherford's expected vs. actual results
Animated Tutorial
Rutherford's expected vs. actual results
Animation
Rutherford expected α-particles to pass undeflected through atoms.
But, he observed that a small fraction of the α-particles were deflected
Evidence that the positively charged part of the atom consisted of a tiny, dense object at the atom's center.
He proposed the nuclear model of the atom.
Alpha particles: positively charged particles. Move extremely fast.

18. Rutherford model- nuclear
Rutherford did propose that the electrons moved around the nucleus like the planets orbiting the sun.
What is the problem with this model?
Charge and mass of atom did not work out!

19. Chadwick- discovers neutrons
Act as a kind of glue to hold the nucleus together.
Positively charged protons are in a very confined space but shouldn’t because they repel each other.
Protons and neutrons are all attracted to each other as a result of another force - the strong nuclear force.
The neutrons don't contribute any repulsive effects, so having more around can help to hold the nucleus together.

20. Subatomic particles: summary
Name
Location
Charge
Mass
Discovered
Proton
Neutron
Electron
Chemical reactions involve changes
Nuclear reactions involve changes

21. Subatomic Particles
Mass of nucleus comes from the mass of protons and neutrons (= the nucleus).
Atomic mass unit: (amu)
1amu = 1/12 of the mass of carbon-12 (1.67 x 10-24g); 12 protons + 12 neutrons

22. The nuclear atom
Nearly all of the atom's mass is in the nucleus, which consists of protons and neutrons. Nearly all of the atom's volume is supplied by the electrons, which exist outside the nucleus. Students can be challenged to explain how the electrons, with negative charge, remain outside of the nucleus, with its positive charge.

23. How small is an atom? FUN FACT
An atom is so small, a single water droplet contains about 5 sextillion(1021) atoms
Electrons are on the outside of the atom with very little mass. Most of the mass of the atom is in a central nucleus.
Therefore, an atom is mostly empty space
You can think of it as being like a marble in the middle of a football stadium.
the marble is the nucleus-on the 50-yard line; spectators are the electrons.
First, a drop of water is between 45 and 50 microliters (µl), if dispensed from a dropper; I will use 50 µl for the calculation. 50 µl of water should have a mass of 50 milligrams (mg) at Standard Temperature and Pressure. Water is H2O, which has a molecular mass of g/mole, so 50 mg of water contains millimoles (mM). According to Avogadro's number, one mole of any compound contains x 1023 molecules, so mM of water has 1.67 x 1021 H2O molecules. There are 1.67 sextillion molecules of H2O in a drop of water.
Or, since you asked for atoms, and each water molecule is composed of three atoms:
There are about 5 sextillion atoms in a drop of water.

24. Size of an atom

25. Atomic Number and Mass Number
Atomic Number: number of protons in the nucleus of one atom - number of electrons in a neutral atom
Mass Number: total number of protons and neutrons in an atom’s nucleus.

26. Atomic Mass
Average Atomic Mass: average mass of all known atoms of an element.
Unit: amu (atomic mass unit)

27. Atomic Mass

28. Isotopes
Naturally occurring isotopes
Atoms of the same element that contain different numbers of neutrons.
What is the chemical symbol? What is their atomic number? What is the mass number of the atom on the left?

29. Stable vs. Unstable isotopes
Radioactive Isotopes: unstable atoms
due to a nucleus with too many or too few neutrons
No amount of neutrons can hold a nucleus together once it has more than 82 protons. Elements with an atomic number greater than 82 have unstable isotopes.
Unstable atoms emit energy in the form of radiation when they break down (decay)
Large nucleus (unstable)  nucleus + energy
Reaction gives off LOTS of energy (= nuclear energy)

30. Discoveries lead to more about atomic theory
X-rays given off from anode when cathode operating (light energy)
Radioactivity- , , , rays
“Quantization of Energy” – 1900 Max Planck. E = hv
1905 Light as a wave and particle
Einstein's Ideas about Light

31. Electromagnetic spectrum

32. Waves
Calculating the frequency of a repeating event is accomplished by counting the number of times that event occurs within a specific time interval, then dividing the count by the length of the time interval. For example, if 71 events occur within 15 seconds, the frequency is: 4.7 Hz (or 4.7 s-1)
If 1= 4s-1 = ?

33. Electromagnetic Spectrum
Speed of light
c = speed of light (3.0 x 108 m/s)
Types of light energy:
 = wavelength  = frequency E = energy
“speed of light = c” only when in a vacuum.
Wavelength= m/cycle
Frequency= cycles/second
Speed of light= meters/second
Energy= joules or calories
c =  

34. Electromagnetic spectrum

35. Diffraction grating/prism
Note: A light bulb is an example of blackbody radiation (continuous spectrum). Most densely
packed solids will emit a continuous spectrum when heated to a certain temperature.

36. Absorption or Emission of light
The atom can absorb or emit light.
Examples of absorption –
the color of shirt.
Photosynthesis
Examples of emission
Gas discharge tubes
Flame tests
Neon lights
Lamps

37. Excited Electrons and Spectra
Line spectrum - can be used to identify an element – it is a characteristic property of that element.
Examples of practical use: determine the chemical make-up of the stars and plants’ atmospheres.
FIREWORKS! SIMILAR CONCEPT TO OUR
FLAME TEST
Different metal will burn different colors.
-What metallic elements do you think are in these fireworks?

38. Continuous vs. Line Spectrum

39. Hydrogen’s line spectrum

40. Another great student… Niels Bohr (student of Rutherford)
Revised Rutherford’s model to include newer discoveries about how an atom could absorb or emit light!
Here’s his thoughts:
Electrons are found in “distinct energy levels”. This means electrons can’t be found in-between these levels.
Like Rutherford he proposed e- orbited the nucleus.

41. Bohr Model
Electrons absorb energy and move to outer energy levels. When they relax, they give off energy.
“Your theory is crazy, but it's not crazy enough to be true”. Niels Bohr

42. Quantum Theory
Classical Theory
Quantum Theory
vs.

43. Bohr Model Energy of photon depends on the difference in energy levels6
Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1

44. Bohr Model
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen! 
Did not agree with classical physics. 

45. Electrons and energy
An electron’s P.E. & K.E both change when it relaxes (down orbital/s) or is in an excited state (up orbital/s)
EXCITED STATE:
Absorbs a photon or quantum of energy elevates to higher energy level
GROUND STATE:
Electrons in their lowest energy levels

47. Atomic structure- Bohr model
Energy level=n
Lowest energy state is closest to nucleus-attracted to the protons
When one energy level is filled, electrons are found at higher levels.
Each energy level can hold a maximum number of electrons (2n2 electrons)
First shell = two electrons
Second shell = eight electrons
Third shell = eighteen electrons

48. Quantum Mechanical Model
Electrons have properties of waves and light (De Broglie)
It is impossible to know both the position and momentum of an electron (Heisenberg)
The probability of finding in electron in a certain area around the nucleus. (Schrödinger)
Sublevels- defined by energy level/distance from nucleus
Orbitals- mathematical function corresponding to a region within atom
each with a maximum of 2 e- with opposite spin

49. Quantum Mechanical Model
Determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus of the atom.

50. Heisenberg Uncertainty Principle
Where is an electron?
Heisenberg Uncertainty Principle
– It is impossible to know both the position and momentum of an electron.

51. S Orbitals

52. P orbitals

53. D orbitals

54. F orbitals

55. An orbital is a mathematical (3D) graph of the solution to the quantum mechanical wave equation. It defines a region of space that has a high probability of containing up to 2 e-. Movie visual

56. How do concepts of energy levels and orbitals fit together?
Each energy level is made of 1 or more sublevels:
Each sublevel is made of 1 or more orbitals:

57. Orbitals are filled from lowest to highest energy, in order of the periodic table

58. Electron Configurations
Aufbau Principle
Electrons fill from lowest energy to highest energy.

59. Electron Configurations
Pauli Exclusion Principle
Paired electrons must have opposite spins.
Each orbital holds 2 electrons.

60. Electron Configurations
Hund’s Rule
Electrons must be unpaired before they are paired in a sublevel.
“Make sure that everyone gets a helping!”
WRONG
RIGHT

61. Abbreviated Configurationsp 1 2 3 4 5 6 7
d (n-1)
f (n-2) 6 7
© 1998 by Harcourt Brace & Company

62. Abbreviated Configurations
Example - Germanium
[Ar]
4s2
3d10
4p2

63. Abbreviated Configurationsp 1 2 3 4 5 6 7
d (n-1)
f (n-2) 6 7
© 1998 by Harcourt Brace & Company

64. Chapter 6
tHE pERIODIC tABLE

65. Periodic Table of Elements
Sing-a-long

66. Names and symbols Symbol Name Symbol Name H Hydrogen He Helium Li
Lithium Be
Beryllium B
Boron C
Carbon N
Nitrogen O
Oxygen F
Fluorine Ne
Neon Na
Sodium Mg
Magnesium Al
Aluminum Si
Silicon P
Phosphorus
Symbol
Name S
Sulfur Cl
Chlorine Ar
Argon K
Potassium Ca
Calcium Fe
Iron Co
Cobalt Cu
Copper Zn
Zinc Ag
Silver Sn
Tin I
Iodine Au
Gold Hg
Mercury Pb
Lead
First letter always capitalized and second/third lowecase

67. Universe’s elements

68. Earth’s elements

69. Human Body Elements

70. Diatomic elements
Label your PT

71. Metals, nonmetals, metalloids
Label your PT

72. Periodic table organization
Groups or families = column, similar chemical properties
Alkali metals
Alkaline earth metals
Halogens
Noble gases
Period = row, chemical and physical trends repeat
Other sections
Transition metals
Metalloids
Metals
Nonmetals
Lanthanide and actinide series (inner transition metals or rare earth)

73. Trends based on # of electrons
Groups (columns)
Elements in the same group have similar properties; why?
They all have the same # of outer electrons=
VALENCE ELECTRONS- Use the periodic table note valence electrons
Periods (rows)
Elements in a period have valence electrons in the same outer energy level.
They all have the same # of inner electrons=
CORE ELECTRONS- Use the periodic table note energy levels

74. Physical properties of elements
Physical state:
gas, solid, liquid
Conductivity:
Conductor, semiconductor
Physical qualities:
Label your PT