1. Chemistry Unit 3
Atomic Structure (Ch.3)

2. 3-1 Early Models of Atoms Democritus (450 BC)
Proposed that all matter was made of tiny indivisible particles.
He called these particles atomos (meaning indivisible).
We call them atoms.
Good looking guy!

3. Atom
An atom is the smallest particle of an element that retains the identity of that element.
If we repeatedly cut a piece of Al, the smallest possible piece is an atom of Al.
Classic model of an atom

4. Aristotle Didn’t agree with Democritus.
Believed matter was continuous and made up of only one substance called “hyle”
It wasn’t until the 1700’s when his ideas were reexamined.

5. Newton and Boyle (1600s)
Published articles stating their belief in the atomic nature of elements
They had no proof

6. Antoine Lavoisier (1770’s)
French, The “Father of Modern Chemistry”
Discovered the law of conservation of matter.
Matter is neither created nor destroyed.

7. Joseph Proust (1799) French Chemist
Developed The Law of Definite Proportions
Compounds always contain elements in the same proportion by mass.

8. Law of Definite Proportions
H20 (by mass is always)
88.9% Oxygen, 11.1% Hydrogen
If we had an 80g sample of H20 how much is O?
.889 x 80 = 71g
How much is H?
.111 x 80 = 9g

9. English schoolteacher
John Dalton (1803)
Proposed the atomic theory of matter, which is the basis for present atomic theory
John Dalton,
English schoolteacher

10. Atomic Theory of Matter
Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical, but differ from those of any other element.
Which element is this?

11. Atomic theory of matter
When elements unite to form compounds, they do so in a ratio of small whole numbers. This is called the Law of Multiple Proportions.
Ex: C and O can combine to form CO or CO2, but not CO1.8.

12. All matter is composed of tiny particles
Dalton’s Model of an Atom
All matter is composed of tiny particles

13. J.L. Gay-Lussac (early 1800s)
Observed that working with gas reactions at constant volume, temperature and pressure are directly related.
He named the discovery of this relationship Charles Law, which is represented by P1/T1=P2/T2.

14. Amadeo Avogadro (early 1800s) – Italian Physicist
Explained Gay-Lussac’s work using Dalton’s theory: Equal volumes of gases at the same temp/pres have the same number of gas molecules.

15. Michael Faraday (1839)
Suggested that atomic structure was related to electricity.
Atoms contain particles that have electrical charges.
Positive (+)
Negative (-)
Opposite charges attract
Like charges repel

16. William Crookes (1870’s) English Physicist
Developed the cathode ray tube to find evidence for the existence of particles within the atom.

17. J.J. Thomson (1896)
Used a cathode ray tube (CRT) to identify negatively charged particles, called electrons.
Determined the ratio of an electron’s charge to its mass.
Developed the “plum pudding” model of an atom.
Cathode ray bending toward a positive charge

18. Plum Pudding Model - + - + + - + - + + + - -
Atoms are composed of randomly arranged charged particles

19. Robert Millikan (early 1900s)
US Physicist
Used the oil drop experiment to prove the electron has a negative charge
Was able to determine the charge of the electron

20. Millikan’s Oil Drop Experiment

21. Bothe/Chadwick (early 1930s)
English
Found high energy particles with no charge with the same mass as the proton called neutrons.

22. Ernest Rutherford (1909)
Used the gold foil experiment to prove the atom is mostly empty space.
Rutherford concluded that all of an atom’s positive charge, and most of its mass is located in the center, called the nucleus. Analogy: thumb nail and the 50 yard line.

24. 98% of the particles passes straight through
2% of the particles deflected off at varying angles
0.01% of the particles bounced straight back

25. Rutherford’s Planetary Model of an atom- - + + + - +
Positive charge and majority of mass located in the nucleus.
Negatively charged electrons orbit the nucleus like planets. + + + - - -
Most of an atom is empty space!

26. Problem
He thought a moving electrical charge (-) in a curved path should lose energy (give off light).
If it did, it would fall into the (+) nucleus.
Why don’t the (-) electrons fall into the (+) nucleus?

27. Atom:The smallest particle of an element that has the properties of that element.
Make up of nucleus consists of protons and neutrons
Surrounded by an electron cloud
Electron cloud

28. Sub-Atomic Particles Protons Positively (+) charged
The number of protons in an atom refers to the atomic number (Z)
Composed of 3 quarks (2 up, 1 down)
Mass= x 10-27kg
Atomic mass  1 amu (µ)

29. Sub-Atomic Particles Neutrons Found in nucleus Neutral (no) charge
composed of 3 quarks (1 up, 2 down)
Atomic mass  amu (µ)
Isotopes- atoms of the same element that have a different number of neutrons.

30. Sub-Atomic Particles Electrons
Found in electron clouds surrounding the nucleus.
Negative (-) charge
Mass = x kg
1800 times smaller than protons & neutrons
Mass  0 amu (µ)

31. Sub-atomic particles Electrons
Orbit the nucleus at very high speed in energy levels (electron clouds).
Negatively (-) charged
Have no mass (when compared to protons and neutrons)

32. Atomic Number = Protons
The atomic number of an element is the number of protons an element has.
Located above the symbol of the element
The number of protons determines the identity of the element.
Each element has a different atomic number

33. Neils Bohr (1913)
Improved Rutherford’s work by saying electrons do not lose energy in the atoms so they will stay in orbit
Stated there are definite levels in which the electrons follow set paths without gaining or losing energy (Planetary Model)
Each level has a certain amount of energy associated with it and the electrons can only jump levels if they gain or lose energy
Could not explain why (-) electrons don’t fall into the (+) nucleus.

34. Energy Levels
In the ground state for an atom, electrons are at their lowest, most stable energy levels.
In the excited state, atoms require energy and electrons move to a higher energy level.

35. How many electrons are in an atom?
For an atom to have an overall neutral charge the number of electrons must equal the number of protons.
#Protons=#electrons
What element is this?

36. Mass number Mass = Protons + Neutrons
The Mass number of an atom is the sum of the mass of protons and neutrons
Located below the symbol of the element
Atomic mass is measured in amu’s, (atomic mass units)
Based on Carbon having a mass of 12
Mass = Protons + Neutrons

37. How many neutrons are in an atom?
Mass=Protons+Neutrons
195= 78 + Neutrons
195-78= Neutrons
Platinum has 117 Neutrons
Find the number of neutrons in:
Hydrogen Carbon
Helium Potassium
Boron Gold

38. Mass =Protons + Neutrons
Hydrogen (H) 1 =1 + Neutrons
Hydrogen has 0 neutrons
Helium (He) 4 = 2 + Neutrons
Helium has 2 neutrons
Boron (B) 11 = 5 + Neutrons
Boron has 6 neutrons
Carbon (C) 12 = 6 + Neutrons
Carbon has 6 neutrons
Potassium (K) 39 = 19 + N
Potassium has 20 neutrons
Gold (Au) 197 = 79 + N
Gold has 118 neutrons

39. Atomic Mass The average mass of all of the isotopes of an element.
Aka: average atomic mass number, or atomic weight.
Isotopes:Atoms of the same element with different masses.

41. Average Atomic Mass
Ne-20 has a mass of amu (u), and Ne-22 has a mass of amu (u). In any sample of 100 Ne atoms, 90 will be Ne-20. Calculate the average atomic mass of Ne.
.90 x =
.10 x =
avg mass = amu

42. Ions Na Na+ An atom that has gained or lost an electron.
It acquires a net electrical charge.
If an atom loses an electron (oxidation) it has more protons than electrons and has a net positive charge. (cation)
11 P
11 e-
11 P
10 e-
Na+

43. Ions
If an atom gains an electron (reduction) it has more electrons than protons and has a net negative charge.(anion)
7 valence e-
Full octet

44. Ionic Charges Charge of ion = # protons - # electrons
What is the charge of a magnesium atom that loses 2 electrons?
Number of protons 12
-Number of electrons 10
charge of ion +2
Mg2+ or Mg+2
Charge is written to the upper right of the symbol.

45. Representations of atoms
General form: (Elemental Notation)
X = Element Symbol
A = Atomic Mass (P + N)
Z = Atomic Number (P)
Ionic Charge
Charge X Z

46. What is the atomic structure?
Determine the number of:
P =
N =
e = 23 Na + 11

47. What is the atomic structure?
Determine the number of:
P = 11
N = 12
e-= 10 23 Na + 11

48. What is the atomic structure?
Determine the number of:
P =
N =
e- = I -
127 53

49. What is the atomic structure?
Determine the number of:
P = 53
N = 74
e- = 54 I -
127 53

50. Put into elemental notation?
Atomic # = 29
Atomic Mass = 64
Ionic charge = +2

51. Cu How many electrons? 64 2+ 29 Atomic # = 29 Atomic Mass = 64
Ionic charge = +2
# of electrons = 64 2+ 29

52. Put into elemental notation
37 Protons
48 Neutrons
36 Electrons ?

53. Rb + Put into elemental notation 37 Protons 48 Neutrons 85
36 Electrons 85 + 37

54. Max Planck (early 1900s)
Proposed Planck’s Theory which says that energy is given off in little packets or particles called quanta which is based on the particle nature of light
Each quantum of energy corresponds to the different energy levels for electrons.
Proposed the equation: E=hf, where E is energy, f is frequency, and h is Planck’s constant (6.63 x 10^-34 J/Hz)

55. De Broglie (1923)
Suggested that if waves can have a particle nature, particles can have a wave nature, known as the “wave-particle duality” principle
Wondered why the positive nucleus and negative electrons do not attract. Proposed that electrons moved so fast (speed of light) that they had properties of waves instead of particles.

56. The Study of Waves
Wave: a progressive disturbance propagated from point to point in a medium or space without progress or advance by the points themselves

57. Types of Waves
Wave Travel
Transverse: displacement of the medium is perpendicular to the direction of propagation of the wave.
Longitudinal: displacement of the medium is parallel to the direction of propagation of the wave
Mechanical: a wave that requires an energy source and an elastic material medium to travel.
Electromagnetic: a wave that does not require a material medium to travel; it propagates by electric and magnetic fields

58. Properties of Waves
Wavelength (ג): The distance between any part of the wave (peak) and the nearest part that is in phase with it (another peak). Standard unit is meters (m).
Frequency (f ): The number of peaks which pass a given point each second. Standard unit is cycles per second which is a hertz (Hz).
Amplitude (A): The maximum displacement of a vibrating particle from its equilibrium position. Standard unit is meters (m).
Velocity (v): the distance a wave (peak) travels in a given time. Standard unit is meters per second (m/s).
Energy (E): The energy of a single photon of radiation of a given frequency. Standard unit is the joule (J).

59. Some relationships between the properties of waves are represented by the equations: V=f*ג and E=h*f , where h=6.63x10^-34 J/Hz

60. Werner Heisenberg (1927)
Proposed his “Heisenberg uncertainty principle”, which says that the position and momentum of an electron cannot simultaneously be measured and known exactly.
The arrangement of electrons is discussed in terms of the probability of finding an electron in a certain location.

61. Erwin Schrodinger (1926)
Studied the wave nature of the electron and developed mathematical equations to describe their wave-like behavior.
The most probable location of the electrons can be found and the plot of this probability is called the charge cloud model.

62. The four quantum numbers
Principal Quantum Number (n)
Refers to the energy levels in the atom which is the distance from the nucleus and designated with a positive whole number (n=1,2,3,etc)
Wavelength of emitted photon is determined by the “energy jump” between energy levels
Energy levels (or shells) means electrons are contained in an area where the probability of finding the electron is 90%
Angular Momentum Quantum Number (l )
Refers to the sublevel (within an energy level) which is one or more “partitions” each with a slightly different energy.
The number of sublevels in a particular energy level is equal to the principal quantum number (n).
The types of sublevels include: s, p, d, f, etc.

63. The four quantum numbers (continued)
Magnetic Quantum Number (m)
Refers to the orientation in space of the electrons in a sublevel
Can have any whole number value from -1 to +1 which will tell how many orbitals are in a sublevel.
A maximum of 2 electrons per orbital.
Sublevel # of Orbitals Total # of electrons s p d f

64. Four Quantum Numbers (continued)
Spin Quantum Number (s)
Refers to the spin of the electron.
Pauli Exclusion Principle : if two electrons occupy the same orbital, they must have opposite spin.
Half-filled orbital: _____
Filled orbital: _____

65. Permissible Values of Quantum Numbers for Atomic Orbitals
n l m Orbital # of Subshells #of Orbitals Max # of Electrons s s
1 -1,0, p s
,0, p
2 -2,-1,0,1, d s
,0, p
,-1,0,1, d
3 -3,-2,-1,0,1,2,3 4f

66. Distribution of Electrons for Different Elements (Electron Configuration)
Electrons will occupy the lowest energy levels and sublevels first.
Notation:
Principal Quantum Number, n (energy level)
Number of electrons 2p 2 y
Orientation of Orbital
Type of Orbital (sublevel)

67. Long Notation: Pyramid Filling
“Rule of thumb” for filling electrons at the lowest energy level possible.

68. Give the long notation electron configuration for:Ca Ag

69. Give the short notationCa Ag

70. Orbital Diagrams Electron Dot Diagrams
Usually only done for the outer shell electrons, which always includes the s and p orbitals.
Electron Dot Diagrams
Shows the outer shell electrons for elements.