1. The Atom

2. Law of Conservation of Mass/Matter
Matter cannot be created or destroyed
Total mass is constant in chemical reactions.
Originated with Antoine Lavoister (1700s)
Quantitative mass data of reactants and products in mercury oxide decomposition.

3. Law of Definite Proportions
Proposed by Joseph Proust (late 1700s)
Decompositions and research with copper carbonate
Compound composition and properties are fixed
All compound samples have the same composition
Same % of elements in the compound
Ex. H2O

4. Law of Multiple Proportions
2+ compounds with same 2 elements
Compositions of these compounds are related
Masses of elements related to each other in whole number ratios
Proposed by John Dalton in addition to his atomic theory.
Ex. CO2 (2:1), CO (1:1)

5. Terminology
Element– basic unit of a substance, contain only ONE type of atom, represented by symbol.
Example: Ag, only contains Ag atoms.
Atom—smallest particle of an element that still contains element properties.
Example: One atom of Au, cannot have a smaller particle of gold and still be gold.

6. Compound vs. Molecule Compounds: more than one element
elements combined in definite proportions
Molecule:
Smallest unit of a compound that still retains the properties of the compound.

7. How far back does the “atom” go?
Democritus
400 B.C.
Called the basic unit of matter an “atom”

8. The Atom and its Structure

9. Dalton Atomic Theory 1800s Atoms make up elements.
Atoms form compounds as a whole and cannot be divided. Compounds formed from atoms joining in FIXED proportions
English schoolteacher.
We now know this theory is not all the way true due to being able to divide an atom—nuclear chemistry.

10. Dalton Atomic Theory (cont.)
Atoms cannot be divided, destroyed, or created.
Atoms rearrange in chemical reactions.
All matter made of atoms
Atoms of an element have the same size, mass, etc.
Different atoms have various sizes, mass, etc.
Rearrange, combine, separate.
True: all matter made of atoms, different atoms differ in properties, atoms do rearrange

11. John Thomson 1897 Cathode-Ray experiments.
Discovered the electron particle and its possible charge (-).
Determined ratio between mass and charge of an electron
Rays move away from wire containing electric charge—most be negative charge.
-tub opposite cathode glows—due to stream of particles/cathode ray.
**Discovered electrons**--they were moving to opposite side and causing deflection., negative charge!
Cathode = positive.

12. Robert Millikan Oil drop experiments. 1909, American
Found the mass of an electron (VERY small) with Thompson’s data
Currently, mass of electron = x 10-31kg
Discovered electron charge
e = x C
Knew oil density, so watch the oil drop and measure mass/watch and note the particles radius.
Electric field helped find the charge of electron.
-Spray oil drops, electric field set up

14. Early Models of the Atom Thompson
Must be a balance between negative and positive charges
“Raisin-Pudding” model
Uniform distribution of positive charge
Positive cloud with stationary electrons

16. Early Models of the Atom Rutherford
How are electrons distributed in an atom?
Discovered alpha particles as 42He
Experiments with Au, Ag, and Pt foils bombarded with alpha particles

18. Early Models of the Atom Rutherford
Mostly empty space
Small, positive nucleus
Contained protons
Negative electrons scattered around the outside
Putting everything together.
Gold foil experiment. –bombard gold foil with alpha particles/positive charge, most particles passed through foil
-others deflected back!!! ---hit something in the atom/come in contact with something in atom, very small because only happened some of the time.
-also positive center due to being repelled.
Discovery of Nucleus

19. James Chadwick 1932 discovered neutrons contained in atom’s nucleus
No charge
Mass approximately same as proton mass

20. Early Models of the Atom Bohr
1913—hydrogen atom structure
Physics + quantum theory
Electrons move in definite orbits around the positively charged nucleus—planetary model
Does not apply as atoms increase in electron number
We will look more at this model later.
-Based on the hydrogen atom.

21. Erwin Schrödinger Quantum mechanics 1926---wave equation
Electrons behave more like waves than particles

22. Heisenberg’s Uncertainty Principle
Electron’s location and direction cannot be known simultaneously
Electron as cloud of negative charge

23. Modern Model of the Atom The electron cloud
Sometimes called the wave model
Electron as cloud of negative charge
Spherical cloud of varying density
Varying density shows where an electron is more or less likely to be

24. How did we discover electron arrangement in an atom?
ELECTROMAGNETIC
RADIATION ! ! !

25. Waves
Repeated disturbance through a medium (air, liquid) from origin to distant points.
Medium does not move
Ex. Ocean waves, sound waves

26. Characteristics of Waves
Wavelength
Distance between 2 points within a wave cycle
2 peaks
Frequency
# of wave cycles passing a point for a particular time unit
Usually seconds.

28. Wavelength and frequency are inversely proportional.

29. Electromagnetic Waves
Produced from electric charge movement
Changes within electric and magnetic fields carried over a distance
No medium needed

30. c = νλ c = speed of light, 3.0 x 10 8 m/s ν= frequency (s-1 or Hz)
Constant
ν= frequency (s-1 or Hz)
λ= wavelength (m)

31. Example 1:
Find the frequency of a green light that has a wavelength of 545 nm.
Frequency = 5.50x1014 Hz

32. Electromagnetic Spectrum
Contains full range of wavelengths and frequencies found with electromagnetic radiation
Mostly invisible, visible range (390 nnm -760 nm)
Different materials absorb/transmit the spectrum differently.

34. Types of Spectra What is a spectra? Continuous spectrum Line spectrum
Spectrum– white light/radiation split into different wavelengths and frequencies by a prism
Continuous spectrum
No breaks in spectrum
Colors together
Line spectrum
Line pattern emitted by light from excited atoms of a particular element
Aided in determining atomic structure

35. Line Spectrum
Pattern emitted by light from excited atoms of an element
Specific for each element
Used for element identification

39. Flame Tests Some atoms of elements produce visible light if heated
Each element has a specific flame color
Examples: Li, Na, Cs, Ca

40. A Bit of Quantum Theory……

41. Max Planck E = hν h= 6.626 x 10 -34 Js (Planck’s constant)
1900
Related energy and radiation
E = hν
h= x Js (Planck’s constant)
E = energy per photon (J)
Quantum---smallest amount of energy
Atoms can only absorb/emit specific quanta
Energy increases with frequency

42. Albert Einstein 1905 Added to Planck’s concept Photons—
Bundles of light energy
Same energy as quantum
Photons release energy and electrons gain energy
Threshold frequency– minimum amount of energy needed by photon to extract electron

43. Example 1
Calculate the energy found in a photon of red light with a wavelength of nm

44. THEREFORE ……… Light is in the form of electromagnetic waves
Photons can resemble particles
Gave raise to the possibility of thinking about wave AND particle qualities of subatomic particles (electron)

45. Atomic Structure Nucleus Electrons Protons Neutrons
Atoms are made of subatomic particles.

46. Atomic Structure Electrons Tiny, very light particles
Have a negative electrical charge (-)
Move around the outside of the nucleus

47. Atomic Structure Protons Much larger and heavier than electrons
Protons have a positive charge (+)
Located in the nucleus of the atom 

48. Atomic Structure Neutrons Large and heavy like protons
Neutrons have no electrical charge
Located in the nucleus of the atom 

49. Atomic Structure

50. Describing Atoms Atomic Number = number of protons
In a neutral atom, the # of protons = the # of electrons
Atomic Mass= the number of protons + the number of neutrons

51. Isotopes The number of protons for a given atom never changes.
The number of neutrons can change. 
Two atoms with different numbers of neutrons are called isotopes
Isotopes have the same atomic #
Isotopes have different atomic Mass #’s

52. Isotopes

53. Ions An atom that carries an electrical charge is called an ion
If the atom loses electrons, the atom becomes positively charged.
If the atom gains electrons, the atom becomes negatively charged
Positive charge—more protons than electrons.
Negative charge—more electrons so the negative charge is more than positive charge
**Think like a balance**
Balance of charges = neutral atom.

54. Ions The number of protons does not change in an ion.
The number of neutrons does not change in an ion.
So, both the atomic number and the atomic mass remain the same.

55. PEN Method for---
H O P
He F S
Li +1 Ne Cl -1
Be Na Ar
U-238 Mg +2 K