1. AP Chem Turn in Hydrate Lab Today: Atoms Review; back page = HW
Lab Thursday
You’re welcome to start coming in before school/during MM to do test corrections
Next test in ~ 2 weeks

2. Evolution of the Atomic Model
Since atoms are too small to see even with a very powerful microscope, scientists rely upon indirect evidence and models to help them understand and predict the structure of an atom.

3. Democritus & Leucippus (~400 BC)
Greek philosophers
first to propose that matter made up of tiny, indivisible particles called atomos, the Greek word for atoms

4. John Dalton (1803) Billiard Ball Model:
Dalton theorized that the basic unit of matter is a tiny particle called an atom
Dalton’s Theory:
All elements are composed of indivisible (can’t be broken down) atoms
All atoms of a given element are identical
Atoms of different elements are different;
Compounds are formed by the combination of atoms of different elements
Billiard Ball Model:
An atom is represented by a hard sphere

5. JJ Thomson (1897)
Used a cathode ray tube to show one of the smaller units that make up an atom
Because the cathode ray deflected towards the + charged plate when an electric field was applied, Thomson concluded that the ray was formed by particles and the particles were negatively charged

6. JJ Thomson (1897)
Thomson discovered that the atom is made up of small, negatively charged particles which he called electrons
Developed the Plum Pudding Model of the atom

7. Robert Millikan (1909)
Millikan’s Oil drop experiment allowed him to determine the mass and charge of an electron

8. Rutherford: Gold Foil Experiment

9. Ernest Rutherford (1911) Conducted the gold-foil experiment:
Directed alpha particles (small, positively charged particles), at a thin piece of gold foil
Most of the alpha particles passed straight through the foil and a few were slightly deflected
Some of the alpha particles were greatly deflected and bounced back  

10. Ernest Rutherford (1911)
Rutherford concluded that atoms have a dense, central core called the nucleus, while the remainder of the atom is essentially empty space
Positively charged particles known as the protons are found in the nucleus

11. James Chadwick First to prove the existence of the neutron
Provided explanation as to why the positively charged protons in the nucleus stayed intact and did not repel each other.

12. Neils Bohr (1913) Bohr Model
The nucleus contained the protons and neutrons
The electrons orbited around the nucleus (like planets orbiting the sun)  
Electrons are shown in concentric circles or shells around the nucleus
The 1st shell can hold a maximum of 2 electrons
The 2nd shell can hold a maximum of 8 electrons
The 3rd shell can hold a maximum of 18 electrons
Electrons in the outermost shell are called the valence electrons

13. Wave-Mechanical/Cloud Model (modern, present-day model)
Developed after the famous discovery that electrons can be viewed as both a wave and particle
Like planetary model, atom is pictured as having dense, positively charged nucleus
The difference in this model is how the electrons are pictured

14. Wave-Mechanical Model
Electrons have distinct amounts of energy and move in areas called orbitals (clouds)
An orbital is a region in which an electron is most likely (most probable) to be located
Value of l
Orbital (subshell)
Orbital Shape
Name* s
sharp 1 p
principal 2 d
diffuse 3 f
fine

15. What is important about the atomic number?
Unique for each element
= # protons in nucleus

16. How do you figure out the number of:
Protons in an atom?
# protons = the atomic number
Electrons in an atom?
# electrons = # protons
FOR NEUTRAL ATOMS, only!
Neutrons in an atom?
# neutrons = mass – # protons

17. What are Isotopes?
Isotopes are different forms of the same element that have a different mass
Isotopes have the same # protons but different # neutrons

18. What are Ions? Ions are atoms that have a positive or negative charge
In an Ion, the number of protons and electrons are not equal

19. Standard Nuclear Notation
Note: for ions, the charge (+ or -) is indicated in the upper right corner

20. What is the ONE THING that determines the identity of an atom?
# of protons (atomic #)

21. Practice

22. Average atomic mass on the periodic table is a weighted average of the masses of all naturally occurring isotopes

23. Average Atomic Mass 28.085 amu , 14 28 𝑆𝑖 151.9641 amu, 63 152 𝐸𝑢
75.91 % Cl-35, 24.08% Cl-37

25. Electromagnetic radiation
- a form of energy that has wavelike properties
- all forms found in the electromagnetic spectrum
Examples of electromagnetic radiation:
Xrays
Microwaves
Light waves
Radio waves – when you hear static on the radio when you aren’t getting a signal, that is the leftover microwaves from the big bang 15 billion years ago (it is cooling off so now the waves are the energy and wavelength of radio waves)

27. WAVELENGTH AND FREQUENCY ARE TWO IMPORTANT PROPERTIES OF WAVES.
DIFFERENT FORMS OF EMR TRAVEL IN WAVES.
WAVELENGTH AND FREQUENCY ARE TWO IMPORTANT PROPERTIES OF WAVES.

28. Wave Characteristics Amplitude Wavelength Wavelength
2 sec
0 sec
Amplitude
Wavelength
Wavelength
Frequency
Amplitude
The shortest distance between two equivalent
Points (meters)
How many waves pass a certain point per sec. (s-1 OR Hz)
Wavelength – the shortest distance between two equivalent points on a wave
Frequency – how many wave pass a certain point per second
Amplitidude –the height of the wave from crest to origin
State: Light travels as photons; photons have no mass but have energy
The energy a photon has is based on the equation E = hv
Energy is directly proportional to the frequency. Frequency has a unit of Hz which is = 1 𝑠 or s-1 or per second.
Wavelength is also related to frequency; how many people know the speed of light? Yes, what is it? EA 3 x 108 m/s
The speed of light is a constant that is known. Speed of light is based on the speed it travels over time. Wavelength is in meters and frequency is per second, so c=vλ
The height of a wave from crest to origin

29. THE HIGHER THE FREQUENCY, THE SHORTER THE WAVELENGTH
FREQUENCY = THE NUMBER OF WAVES THAT PASS BY A FIXED POINT PER SECOND.
THE HIGHER THE FREQUENCY, THE
SHORTER THE WAVELENGTH

30. THE HIGHER THE FREQUENCY, THE HIGHER THE ENERGY
FREQUENCY = THE NUMBER OF WAVES THAT PASS BY A FIXED POINT PER SECOND.
THE HIGHER THE FREQUENCY, THE HIGHER THE ENERGY

31. THE HIGHER THE ENERGY, THE SHORTER (SMALLER) THE WAVELENGTH

32. Practice

33. Photons - released during electromagnetic radiation
– tiny particles that have no resting mass which carry a quantum of energy
All photons travel at the speed of light in a vacuum
Photons behave like particles
quantum – the minimum amount of energy that can be absorbed or released from an atom
- cannot be any value but are in discrete energy levels (like bundles or packets of energy)

34. 𝑐=𝑣λ
3.00× 10 8 =𝑣×4.9× 10 −7
𝒗=𝟔.𝟏× 𝟏𝟎 𝟏𝟒 𝑯𝒛

35. 3.17 m
1.70 x J
8.28 x J
1.0 x 10-5 m

36. Electrons and Energy Levels
Each electron has a distinct amount of energy that is related to the energy level (shell) it is in
Electrons with the lowest energy are found in the shell closest to the nucleus
Electrons with the highest energy are found in the shell furthest from the nucleus
The greater the distance from the nucleus, the greater the energy of the electron

37. Ground vs. Excited State
Ground state=when the electrons occupy the lowest energy levels possible
Excited state= on or more electrons gain energy and move to a higher energy level or shell (leaving the lower energy levels to be not completely full)

38. Ground and Excited State
When an electron gains energy, it jumps to a higher energy level or shell
This is a very unstable condition
We call this condition the excited state
Very rapidly, an electron in the excited state will lose energy and move back to a lower energy level or shell
When excited electrons fall from an excited state to a lower energy level, they release energy in the form of light

39. Ground  Excited State
Electron gains energy from heat, light, electricity
Electron “jumped” to a higher energy level (shell)

40. Excited  Ground State Electron releases energy in the form of light
Electron “falls” back and returns to ground state (normal position)

41. Bright Line Spectrum
Electrons falling from an excited state down to the ground state give off visible light
Different elements produce different colors of light or spectra
These spectra are unique for each element (just like a human fingerprint is unique to each person)