2. Objectives Identify important developments in the history of atomic theory. Summarize Dalton’s atomic theory. Describe the size of an atom. Distinguish among protons, electrons, and neutrons in terms of relative mass and change. Describe the structure of an atom, including the location of the protons, electrons, and neutrons with respect to the nucleus. Explain how the atomic number identifies an element. Use the atomic number and mass number of an element to find the number of protons, electrons, and neutrons. Explain how isotopes differ and why the atomic masses of elements are not whole numbers. Calculate the average atomic mass of an element from isotope data. TEKS: 2A, 2B, 2C, 2D, 2E, 3A, 3C, 3E, 4A, 4C,4D, 5A, 6A, 6B, 6C, 8A, 9B, 10A, 11A
3. Early Models of the Atom400 B.C. – Democritus proposed the existence offundamental particles of matter that were indivisible andindestructible - “atomos”.Aristotle thought all matter was continuous; he did notbelieve in atoms.Neither idea was supported by any experimentalevidence – speculation only.
4. Foundations of Atomic Theory The late 1700’s –definitions and basic laws had been discovered and accepted by chemists. Element – substance that cannot be broken down by ordinary chemical means. Chemical Reaction – transformation of substance or substances into one or more new substances.
5. Law of Conservation of Mass – mass cannot be created or destroyed just changed from one form to another. (Antoine Lavosier) Law of Definite Proportions – a chemical compound contains exactly the same elements in the same proportion regardless of sample size. (Joseph Proust from work of Gay-Lussac & Amadeo Avogadro – 1802/1804) Law of Multiple Proportions – If two or more different compounds are composed of the same two elements, then the ratio of the masses of those elements will always exist as a ratio of small whole numbers. (John Dalton - 1808)
6. Dalton’s Atomic Theory All elements are composed of tiny indivisible particles called atoms. Atoms of the same element are identical. The atoms of one element are different from the atoms of another element. Atoms combine in simple whole-number ratios. Atoms are separated, joined or rearranged in chemical reactions. Atoms of one element are never changed into atoms of another element as a result of a chemical reaction.
7. Discovery of Electrons 1897 – J.J. Thomson – “Cathode Ray Tube Experiment” Showed existence of first know sub-atomic particle Determined charge to mass ratio of the electron 1909 – Robert Millikan found the charge of the electron – “Millikan’s Oil Drop Experiment”
8. Cathode Ray Tube High Voltage Gas at very low pressure Metal disk (anode)Metal disk Cathode Ray(cathode) (electrons)
9. Cathode Ray Tube High Voltage Gas at very low Negative plate pressure Metal disk (anode)Metal disk Positive plate Cathode Ray(cathode) (electrons)
10. Rutherford’s Gold Foil Experiment Rutherford, Geiger & Marsden (1912) -showed that most of the atom was empty space, but that atoms had a solid, positive core. Alpha ParticlesLeadshield Radioactive source
11. Discovery of Protons 1919 -J.J. Thomson & James Chadwick– discovered particles traveling opposite of the cathode rays. Determined existence, mass and charge of protons Idea had actually been previously proposed by Goldstein in 1886.
12. Cathode Ray Tube High VoltageGas at verylow pressure Negative plateprotons Metal disk (anode) Metal disk Positive plate Cathode Ray (cathode) (electrons)
13. Neutrons James Chadwick 1932 - confirmed the existence of the neutron. Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton. Walter Bothe had first reasoned the existence of a third subatomic particle in 1930. Bothe’s work was based in part on that of Henry Mosely who showed by X-ray analysis that not all atoms of the same element were identical. (Isotopes – 1907)
14. Counting Particles Atomic Number = number of protons Mass Number = number of protons and neutrons Atomic Mass = average mass of the isotopes (also known as atomic weight)
15. Periodic Table atomic number # of protonsmass number 8-atomic number O round to 16 - mass number ( # protons# of neutrons & neutrons) 15.999 unrounded –mass number (average mass of the isotopes)
16. Masses of Atoms A scale designed for atoms gives their small atomic masses in atomic mass units (amu) An atom of 12C was assigned an exact mass of 12.00 amu Relative masses of all other atoms was determined by comparing each to the mass of 12C An atom twice as heavy has a mass of 24.00 amu. An atom half as heavy is 6.00 amu.
17. Atomic Mass Listed on the periodic table Gives the mass of “average” atom of each element compared to 12C Average atom based on all the isotopes and their abundance %. Atomic mass is not a whole number Na due to isotopes. 22.99
18. Isotopes Isotopes – atoms of the same element with different numbers of neutrons. Oxygen-16 Oxygen-17 Oxygen-18 16 17 18 8 8 8p+ ‗‗‗‗ ‗‗‗‗ ‗‗‗‗e- ‗‗‗‗ ‗‗‗‗ ‗‗‗‗nº ‗‗‗‗ ‗‗‗‗ ‗‗‗‗
19. Calculating Average Atomic Mass Percent(%) abundance of isotopes Mass of each isotope of that element Weighted average = mass isotope1(%) + mass isotope2(%) + … 100 100
20. Atomic Mass of Magnesium Isotopes Mass of Isotope Abundance 24Mg = 24.0 amu 78.70% 25Mg = 25.0 amu 10.13% 26Mg = 26.0 amu 11.17%Atomic mass (average mass) Mg = 24.3 amu Mg 24.3
21. #16 The element copper has naturally occurring isotopeswith mass numbers of 63 and 65. The relative abundanceand atomic masses are 69.2% for mass = 63.0 amu, and30.8% for mass = 65.0 amu. Calculate the averageatomic mass of copper.
22. Finding An Isotopic Mass Naturally occurring boron is 80.20% boron- 11 (atomic mass 11.0 amu) and 19.80% of a different isotope of boron. What must the mass of this isotope be if the average atomic mass of boron is 10.81 amu?
23. Radioactivity Mosely’s X-ray analysis of atoms was an attempt to explain radioactivity. 1896 – Henri Becquerel – Uranium spontaneously emits energy. 1898 – Marie & Pierre Curie – first isolated a radioactive element - Radium
24. Properties of Subatomic ParticlesParticles Symbol Charge Relative Mass MassElectron e- 1- 1/1840 amu 9.11 x 10-28 gProton p+ 1+ 1 amu 1.67 x 10-24 gNeutron nº 0 1 amu 1.67 x 10-24 g
25. “Planetary” Model of the Atom Niels Bohr (1913) – developed the “planetary” model of the atom based upon the following: Rutherford’s Gold Foil Experiment E = mc2 – Albert Einstein (1905) Quantum Theory – Max Planck (1910)
26. Atom 10-13 cmelectrons protons neutrons nucleus 10-8 cm
27. Size of the Atom Aluminum Atom 150 m e-1 mm Outside edge of Al e- e- atom stands e- e- goal post e- nucleus - size e- of a marble e- e- e- Texas Memorial Stadium @ UT