History and Progression of Atomic Theory

1. Historyand Progression of Atomic Theory

2. Structure of the Atom • An atom is the smallest particle of an element that retains the chemical properties of that element. • The nucleusis a very small region located at the center of an atom. • The nucleus is made up of at least one positively charged particle called a protonand usually one or more neutral particles called neutrons.

3. Surrounding the nucleus is a region occupied by negatively charged particles called electrons. • Protons, neutrons, and electrons are often referred to as subatomic particles.

4. Democritus 400 BC • This is the Greek philosopher Democritus who began the search for a description of matter more than 2400 years ago. • He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?

5. Democritus’ Atomic Theory 400BC Democritus asserted that space contained an infinite number of particles Named atomos, “not cutting” or "indivisible” Atoms are eternal and invisible; absolutely small, so small that their size cannot be diminished; totally full and incompressible. Atoms are homogeneous, differing only in shape, arrangement, position, and number

6. Atomos • To Democritus, atoms were small, hard particles that were all made of the same material but were different shapes and sizes. • Atoms were infinite in number, always moving and capable of joining together.

7. Dalton’s Theory (early 1800’s) • He deduced that all elements are composed of atoms. Atoms are indivisible and indestructible particles. • Atoms of the same element are exactly alike. • Atoms of different elements are different. • Compoundsare formed by the joining of atoms of two or more elements in specific ratios. • In chemical reactions, atoms are combined, separated, and rearranged • Explained the Law of the Conservation of Mass Matter is neither nor destroyed during ordinary chemical reactions or physical changes

8. Modern Atomic Theory • Atomic Theory has been modified • We now know that atoms are divisible into smaller particles • And that elements can have atoms with different masses (isotopes)

9. J.J. Thomson (English physicist1897) • Used Cathode Ray tube to determine the presence of – (electrons) and + (protons) particles.

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11. Current passed through a glass tube filled with various gases at low pressure from the cathode to the anode (terminals of the voltage source) • Stream of particles glowed (cathode ray) • Cathode rays deflected away from negatively charged objects • Led to the hypothesis that cathode ray particles are negatively charged • These negatively charged particles were named electrons

12. J.J. Thomson • Plum Pudding Model -- the structure of an atom is something like pudding. He assumed that the basic body of an atom is a spherical object containing electrons & protons randomly confined in homogeneous jellylike material. Positive charges cancel the negative charges. • This model was soon disproved

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14. Ernest Rutherford (Eng.1911) • Atoms have a central positive nucleus surrounded by negative orbiting electrons. • This idea was the result of his famous Gold Foil Experiment(see next slide). This experiment involved the firing of radioactive alpha particles through gold foil. • Alpha particles = + charged particles about 4x mass of H • This model suggested that most of the mass of the atom was contained in the small nucleus, and that the rest of the atom was mostly empty space. • Most of particles passed straight through the foil but approximately 1 in 8000 were deflected. • Conclusion- volume of the nucleus very small compared to the volume of the atom • To put the size or the nucleus in perspective, if the nucleus of an atom was the size of a marble, the atom would be the size of a football field!

15. Rutherford’s Gold Foil Alpha Scattering Experiment

16. Alpha scattering due to repulsion from the nucleus

17. Nucleus-Positive Charged Center The nucleus is far too large in this drawing.

18. Niels Bohr (Danish 1913) • The Bohr Model is probably familiar to us as the "planetary model" of the atom is used to symbolize atomic energy. • Electrons orbit the nucleus much like planets orbiting the Sun. • Electrons travel in certain orbits at specific distances from the nucleus with definite energy levels (energy shells or energy levels) • Lower energy state near the nucleus and higher energy state as you move away from the nucleus • Electrons gain or lose energy by jumping between orbits, absorbing or emitting electromagnetic radiation Think of rungs on a ladder, except the distance between the rungs is not constant!

19. Bohr Model of the Atom

20. Bohr Atom used as logo for Atomic Energy Commission

21. Chadwick (English 1932) • James Chadwick discovered a third type of particle, which he named the Neutron. • Neutrons help to reduce the repulsion between protons and stabilize the atom's nucleus. • Neutronsalways reside in the nucleus of atoms and they are about the same mass and size as protons. • Neutrons do not have any electrical charge; they are electrically neutral

22. Chadwick’s atom w/ P+, e- & N+/-

23. Table 1 (p.76)

24. Cloud Theory(1920’-1930’s) Quantum Mechanics

25. Quantum Model of the Atom • Louis de Broglie (French scientist 1924) • Hypothesized that electrons have wavelike properties • Electrons exist at certain frequencies corresponding to specific energies • Confirmed experimentally by investigating diffraction (bending of wave as it passes through small opening) and interference (waves overlap)

26. Cloud Theory • Based on the work of many scientists • Based on the mathematical approach of Quantum Mechanics • Electrons are assigned regions of space (Orbitals) not pathways (Orbits) • Electrons are moving around the nucleus rapidly in no predictable path producing a cloud of e-’s over time. Think of a rapidly moving fan blade.

27. Cloud Theory of Today Electron Cloud

28. Quantum Theory • Describes mathematically the wave properties of electrons and other very small particles • Schrodinger Wave Equation (Austrian 1926)- showed mathematically how electrons can exist as waves of certain energies and therefore certain frequencies

29. Heisenberg Uncertainty Principle(Heisenberg- German 1927) • Electrons act as particles and waves • How can you locate electrons in an atom? • Observe interference with photons • This knocks the electron off course • “It is impossible to determine simultaneously both the position and velocity of an electron or any other particle”

30. Atomic Number • Atoms of different elements have different numbers of protons. • Atoms of the same element all have the same number of protons. • The atomic number(Z) of an element is the number of protons of each atom of that element.

31. Atomic Number

32. Mass Number • The mass number is the total number of protons and neutrons that make up the nucleus of an isotope.

33. How to find # of Protons • Number of protons in an atom is ALWAYS equal to the Atomic Number

34. Number of Electrons (e-) • Atoms – Number of protons and electrons are equal (overall neutral charge) • Ions • Loss of electron(s) makes positive ions • Gain of electron(s) makes negative ions

35. Loss or Gain of Electrons • Atom Ion • Na  Na+1 + 1e- (LOSS) • +11 +11 • -11 -10 • 0 net +1 net charge • Cl2 + 2e-  2Cl- (GAIN) • +17 +17 • -17 -18 • 0 net -1 net

36. How to calculate # of Neutrons • Atomic Mass (rounded to integer) • - Atomic Number • ---------------------------------------------------- • Number of Neutrons in the nucleus • Atomic Mass – Atomic Number = # Neutrons

37. Atom Contents • Protons (p+) – always equal to Atomic # • Electrons (e-) • In Atoms – Same as the # of Protons • In Ions – Net charge after e-’s have been lost or gained in an attempt to become stable • Loss of e-’s = Positive charge • Gain of e-’s = Negative charge • Neutrons (n0) = Atomic Mass – Atomic # • Isotopes

38. Isotopes • Isotopes are atoms of the same element that have different masses. • The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons. • Most of the elements consist of mixtures of isotopes.

39. Designating Isotopes • Hyphen notation: The mass number is written with a hyphen after the name of the element. • uranium-235 • Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number.

40. The number of neutrons is found by subtracting the atomic number from the mass number. • mass number − atomic number = number of neutrons • 235 (protons + neutrons) − 92 protons = 143 neutrons • Nuclideis a general term for a specific isotope of an element.

41. Sample Problem A • How many protons, electrons, and neutrons are there in an atom of chlorine-37? • Solution: • atomic number = number of protons = number of electrons • mass number = number of neutrons + number of protons • # P+= 17 • # e-= 17 • # N0= mass number - #P+= 37 – 17 = 20

42. Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

43. Law of multiple proportions: if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

44. Law of Conservation of Mass

45. Relative Atomic Masses The standard used by scientists to compare units of atomic mass is the carbon-12 atom, which has been arbitrarily assigned a mass of exactly 12 atomic mass units, or 12 amu. One atomic mass unit,or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

46. Average Atomic Masses of Elements Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. Calculating Average Atomic Mass The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

47. Average Atomic Masses of Elements, continued Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 amu, and 30.85% copper-65, which has an atomic mass of 64.927 amu. The average atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.

48. Average Atomic Masses of Elements, continued (0.6915 × 62.929 601 amu) + (0.3085 × 64.927 794 amu) = 63.55 amu The calculated average atomic mass of naturally occurring copper is 63.55 amu.

49. Hotel Head to Toe Activity • How many floors were there in the hotel? • How many double beds were in the s rooms? The p rooms? The d rooms? The f rooms? • How many guests would each room hold? • What was special about how guests had to sleep in the beds? • We will now compare our hotel and its rooms and guests with electrons filling orbitals in atoms

50. Guests = electrons • Floors = Energy levels (1, 2, 3, 4, 5, 6, or 7) • Rooms = Energy sub-levels (s, p, d, and f) • Double beds = Orbitals • The way guests sleep in the bed = Electron spin