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DEVELOPMENT OF THE ATOMIC MODEL From Democritus to Rutherford

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1 DEVELOPMENT OF THE ATOMIC MODEL From Democritus to Rutherford

2 c.400 B.C. Ancient Greek Philosophy
Everything in the universe is made of one or more of the basic “elements:” Earth, Fire, Water, Air

3 What makes up the elements?
ARISTOTLE Matter is infinitely divisible; no matter how small a piece is, it can always be divided into smaller pieces DEMOCRITUS There exists a “smallest piece” of matter, which cannot be divided any further. These pieces are called “ατομοσ,” or “atoms”

4 Whose Argument Prevails?
Aristotle’s viewpoint enjoyed the support of most of the world because he was more well-known and because Democritus had no evidence to back up his claim since these “atoms” would be too small to see.

5 Antoine Lavoisier 1743-1794 Father of Modern Chemistry
Chemical Revolution – Stated the Law of Conservation of Mass Oxygen for combustion Decomposed Water into two gases, hydrogen and oxygen and then reformed the exact same amount of water.

6 1803: John Dalton Dalton was an English schoolteacher
Began teaching mathematics and chemistry at the age of 12 Revived the idea of Democritus’ “smallest piece” of matter

7 Dalton’s Atomic Theory
All matter is made of tiny particles called “atoms” Atoms are indivisible and indestructible Atoms of the same element are identical Atoms of different elements differ in some fundamental way Atoms combine in simple whole number ratios to form compounds Atoms are rearranged in chemical reactions but cannot be created or destroyed

8 Three Laws Explained by the Atomic Theory
Law of Conservation of Mass (Antoine Lavoisier) Law of Definite Proportions - compounds always contained the same mass ratio of one element to another. (Joseph Proust) Law of Multiple Proportions – When elements combine in different ratios, each new ratio is a unique compound. (John Dalton)

9 So what? These results could only be explained by assuming that matter was made of atoms – tiny building blocks – and that these atoms only came in certain sizes. Dalton’s View of an atom

10 1897: J.J. Thomson English physicist
Worked with Cathode-Ray Tubes (CRTs) Credited with the discovery of the electron

11 CRTs Mysterious particles emanated from the cathode end
These particles were deflected by magnetic and electric fields They were very small and negatively charged These particles were called “electrons” and were assumed to be a part of all matter


13 Thomson’s Atomic Model
Electrons Positively charged “goo” A.K.A. the “Plum-Pudding Model”

14 : Robert Millikan Set out to discover the charge of a single electron Famous experiment called the “oil-drop experiment” Using his results and the charge-to-mass ratio from Thomson, the mass of an electron was found

15 The Oil Drop Experiment
To view an animation of this experiment click below OIL DROP EXPERIMENT

16 1910: Ernest Rutherford Expert in radiation
Famous “Gold-Foil Experiment” Discovered the presence of the nucleus by firing alpha particles at a sheet of gold foil

17 The Gold Foil Experiment
To view an animation of this experiment click below GOLD FOIL EXPERIMENT

18 Rutherford’s Atom


20 Rutherford’s Atomic Model
Electrons Empty Space Nucleus Positively charged Made of “protons”

21 Gold Foil Conclusions The atom has a nucleus
The nucleus has a positive charge The nucleus is very small and very dense Most of the atom is empty space The electron resides in the region outside the nucleus

22 1932: James Chadwick A fellow researcher with Rutherford, Chadwick discovered years later that the nucleus was not made of only one particle – the proton – but of two particles. This second particle was called the “neutron” because it had no electrical charge

23 Subatomic Particles Electron (e-) – mass 1/1840 amu, charge -1, found in space around the nucleus Proton (p+) – mass of 1 amu, charge +1, found in the nucleus Neutron (n) – mass of 1 amu, no charge, found in the nucleus

24 The Atom Made up of 3 fundamental subatomic particles: protons, neutrons, and electrons Very small and very dense nucleus (nucleus make up over 99% of atom’s mass) Nucleus contains protons and neutrons Electrons occupy the empty space outside the nucleus # of protons = # of electrons

25 What’s in the atom? Nucleons – particles in the nucleus (protons and neutrons) The combined total of the protons and neutrons is called the mass number The number of protons is called the atomic number. The atomic number identifies the element. Electrons found in the space outside the nucleus Lots of empty space

26 Isotopes Particles with the same number of protons and electrons but different numbers of neutrons. Have different mass numbers. Have different masses. React the same chemically.


28 Rutherford’s Dilemma

29 If the electron is in orbit around the nucleus, it should be emitting radiation, but it is not.
What prevents the electron from being pulled into the nucleus?


31 Electromagnetic Properties
Wavelength – distance between consecutive waves. Frequency – the number of waves that pass a point in a given amount of time, usually 1 s.

32 Quantum Theory Proposed by Max Planck Two main ideas:
Energy changes are not continuous but rather occur in small increments called “Quantums”. The energy of a quantum is directly proportional to the frequency of the radiation.

33 ….. Quantum theory continued.
E = hν E = energy of a quantum h= Planck’s constant (6.63 x J.s) V = frequency

34 What is the energy of a quantum with a wavelength of 550 nm?

35 Photoelectric Effect

36 Smart Guy In explaining the photoelectric effect Albert Einstein showed that radiant energy, such as light, can posses particle-like properties.

37 Neils Bohr

38 Hydrogen’s Spectrum

39 Questions? Why do the lines always occur at the same place and why only 4 lines? Ans: Only specific energy changes are possible in an atom. The lines are representations of those energy changes.

40 Bohr reasoned that if an electron could occupy an infinite number of possible orbits, its jumps from these orbits should give rise to an infinite number of different energy radiations……….THUS A continuous spectrum

41 Bohr’s Interpretation
Electrons in atoms can only occupy certain fixed orbits or “energy levels”. These energy positions are quantized, meaning only certain values are possible within an atom. To move from one energy orbit to another one, an electron must absorb or emit a quantum of energy exactly equal to the energy difference between the two positions.

42 Electronic Transitions


44 Since Bohr . . . The Modern Model
The Quantum Mechanical Model Has a nucleus Electrons are in a “cloud” of negative charge. An electron “orbit” is an area where the electron is most “likely” to be.

45 Quantum Numbers Describe the properties of orbitals and electrons in those orbitals. Principal q.n. (n) – designates the main energy level or shell. (Bohr) Values: 1, 2, 3, 4, 5, …… n=1; means the electron is located in the first energy level, which has the lowest energy.

46 2. Angular momentum q.n. – (l) designates the shape of the atomic orbital. values: 0, 1, 2, 3, n-1 so if n=3, l can be 0, 1, or 2. if l = 0, (s) then it is spherical if l = 1, (p) then it is dumbbell if l = 2, (d) complex if l = 3, (f) complex

47 Orbital Shapes

48 3. Magnetic q. n. (ml) – designates the orbital’s orientation in space
3. Magnetic q. n. (ml) – designates the orbital’s orientation in space. values – (from –l through 0 to +l) ex. If l = 1 (p) then Ml can = -1, 0, corresponds to px 0 corresponds to py +1 corresponds to pz

49 4. Spin q. n. (ms) – describes the spin of the electron on its axis; clockwise or counter clockwise. values - +1/2 or – 1/2

50 Electron Configurations
Describes the arrangement of electrons in an atom. Each electron in an atom has a set of 4 q. n.

51 Rules that govern orbital filling
Aufbau principle – electrons enter orbitals of lowest energy first. Hund’s rule – when electrons enter orbitals of equal energy, degenerate orbitals, each orbital receives one electron, with parallel spins before any receive two. Pauli exclusion principle – no two electrons in an atom can have the same set of 4 q.n. They must have opposite spins.

52 Order of orbital filling

53 1869: Dmitri Mendeleev Russian chemist
Arranged elements in tabular form so that elements with similar properties were in the same column When listed in order by mass, elements generally repeat properties in groups of 8 (Law of Octaves)

54 The First Periodic Table
Most tables at the time listed elements by mass Mendeleev also arranged elements by mass, but left several “holes” in his table and occasionally reversed the order of elements to fit the properties of others in that column The “holes” were later filled in with newly discovered elements that had the properties predicted by Mendeleev’s table. The reason for the reversal of elements was explained later by Henry Moseley, who noted that the elements were in order by atomic number (number of protons) rather than by mass

55 Introducing the Elements
The Element Song Periods- Horizontal Rows (7 periods) Groups/families – Verticle columns (18 groups)


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