1. Chapter 3 & 24 ATOMIC STRUCTURE

2. Early Models and the Men (and one Woman) Who Discovered Them  Why are the following people important?  Democritus (~450 BC) Proposed that all matter is composed of tiny indivisible particles (atmos) Today’s definition of ATOM – the smallest particle of an element that retains the properties of that element This idea was opposed by Aristotle and others – “What holds everything together?” Democritus could not explain this.  Antoine Lavosier (1770’s) – ***Law of Conservation of Mass*** Matter is neither created nor destroyed in a chemical reaction

3. Early Models and the (Wo)Men  Joseph Proust (1799) – Law of Constant Composition Any given compound always contains the same elements in the same proportions by mass Ex.Water (H 2 0)Oxygen – 88.9%Hydrogen – 11.1% Carbon Dioxide (CO 2 )Carbon – 27.3%Oxygen – 72.7%

4. Early Models and the (Wo)Men  John Dalton – ***Dalton’s Atomic Theory of Matter*** 1. Each element is composed of small particles called atoms 2. Atoms of a given element are identical but different from atoms of other elements 3. Atoms are neither created nor destroyed in any chemical reaction 4.Any molecule of a given compound has the same relative number and kinds of atoms These four postulates are important and you need to know them!!

5. Early Models and the (Wo)Men  Benjamin Franklin (1780s) – static electricity  Michael Faraday (1839) – determined that while atoms are electrically neutral they, somehow, contain particles that have an electric charge.  Investigation of electricity: Negatively charged electrode = CATHODE Positively charged electrode = ANODE Cathode ray tube – an evacuated tube with two electrodes a.streams particles from cathode to anode b.particle turn a paddle therefore confirming they are, indeed, particles c.deflected by a magnet, therefore confirming they are charged

6. Early Models and the (Wo)Men  **J. J. Thomson (~1896), an English physicist, discovered atoms are divisible and that one of the parts carries a negative charge by using a cathode ray tube and a magnet. He called that particle an ELECTRON. He also determined the charge to mass ratio of an electron but was not able to determine the mass directly. He posited the “Plum Pudding” model of the atom.  Robert Millikan (1909), an American physicist, successfully measured the charge of an individual electron and was then able to calculate the mass of an electron using Thomson’s discovery. Charge = 1.60 X 10 -19 coulomb Mass = 9.11 X 10 -28 gram

7. Early Models and the (Wo)Men  Henri Becquerel (1896), a French physicist, accidentally discovered radioactivity by placing a piece of uranium next to an unexposed photographic plate.  **Marie Curie (~1900) isolated two new radioactive elements, radium and polonium.  Ernest Rutherford (early 1900s), a New Zealand scientist, began an extensive study of radiation. His first notable experiment demonstrated the existence of neutrons and of three types of radiation that he called: alpha, beta and gamma. (See description of experiment on pg. 100 of your text!)

8. Rutherford’s Gold Foil Experiemnt  In 1909, Rutherford and his colleagues set out to determine how the atom, which they knew consisted of both positive and negative particles, could be arranged so that each atom was electrically neutral. The Gold Foil Experiment – An radiation source in a lead block with a small opening was used to direct a beam of alpha (positive) radiation toward a piece of gold foil which was surrounded by a circular fluorescent screen. Rutherford believed he could deduce the structure of the atom by observing how the alpha particles were deflected when hitting the gold foil.

9. Rutherford’s Gold Foil Experiemnt The deflection pattern was quite unexpected: - Most of the alpha particles passed straight through the foil - While Rutherford expected only minimal angle scattering, he found particles scattered in all directions - Most notable was within the small fraction of alpha particles, approximately 1 in 8000, were scattered, some were reflected directly back toward the radiation source Rutherford concluded the “Plum Pudding” model was wrong and that - all of the atom’s positive charge was concentrated in a small volume that he called the “nucleus” - the electrons circled the nucleus leaving mostly space inside an atom

10. Modern Atomic Theory  Atoms are composed of three fundamental particles: 1. Protons which have a positive (+) electrical charge, are indicated by the symbol p + and have a mass of ~ 1 amu 2. Neutrons which are neutral and have no (0) charge, are indicated by the symbol n 0 and have a mass of ~ 1 amu 3. Electrons which have a negative (-) electrical charge, are indicated by the symbol e - and have a mass of ~ 0 amu (actually about 1/2000 th of an amu which is insignificant for our purposes)  Protons and neutrons are found in the nucleus at the center of the atom. The nucleus has a positive charge.

11. Modern Atomic Theory (cont.)  Electrons reside in clouds, known as orbitals, that surround the nucleus  As elements not involved in a chemical reaction, atoms are electrically neutral and, as such … the # of p + = the # of e -  Review the chart at the top of pg. 104 for actual charges and mass in grams of the fundamental particles.

12. Modern Atomic Theory (cont.)  The ATOMIC NUMBER of an element is equal to the number of protons in the nucleus.  An element’s identity is determined by the number of protons in the nucleus. For example, an atom with three protons is Lithium but if you add one more proton the atom is now Beryllium.  And since atoms of elements not involved in a chemical reaction are neutral the number of protons equals the number of electrons.

13. Modern Atomic Theory (cont.)  Ion – an atom, or group of atoms, that have gained or lost one or more electrons to obtain a net electrical charge Charge of Ion = number of protons (+) - number of electrons (-) Example:When an atom of magnesium becomes an ion it loses 2 electrons, so …. 12 protons (+12) - 10 electrons (-10) = Mg +2  REMEMBER an element’s identity is determined by the number of protons in the nucleus so a change in charge (# of electrons) does not change the identity of the element!!!

14. Modern Atomic Theory (cont.)  ISOTOPE – atoms that have the same number of protons but a different number of neutrons. Every element has at least two or three isotopes. For example, hydrogen has three, carbon has eight! NameHydrogenHydrogen-2Hydrogen-3 Symbol 1 1 H 2 1 H 3 1 H CommonHydrogenDeuteriumTritium # of p + 111 # of n 0 012 # of e - 111 Mass123 Very common RareVery rare

15. Modern Atomic Theory (cont.)  Isotopes are virtually identical chemically because the chemical properties of an element depend, primarily, on it’s number of electrons and protons.  REMEMBER an element’s identity is determined by the number of protons in its nucleus so the number of neutrons does not effect the identity of the element, just its mass.  MASS NUMBER – the sum of the number of protons and neutrons in an element’s nucleus. Chlorine-37 (Cl-37) has 17 p + and 20 n 0 which yields a mass number of 37 Nitrogen-15 (N-15) has 7 p + and 8 n 0 which yields a mass number of 15

16. Nomenclature With Symbols  To name an isotope using chemical symbols, place the mass number to the upper left of the element’s symbol.  In practice, this method also includes the atomic number (to the lower left of the symbol) and, if it is also an ion, the ionic charge to the upper right of the symbol. 56+2 Fe 26 Element: IronAtomic Number: 26Atomic Mass: 56Charge: +2 Protons (p + ) = 26Neutrons (n 0 ) = 56 – 26 = 30 Electrons (e - ) = 26 – 2 = 24 Note: the “+” charge means there are fewer electrons than protons, in this case, 2 fewer electrons.

17. Nomenclature With Symbols (cont.)  What if I just give you this: p + = 21, n 0 = 24, e - = 18 Answer: 45+3 Sc 21

18. Atomic Mass  1 amu = 1/12 th of the mass of a C-12 atom = 1.66 X 10 -24 g  Atomic Mass is the average mass of all of an element’s isotopes in proportion to their abundance in nature IMPORTANT – the mass shown on the periodic table for each element is NOT the mass of any atom of that element, it is just a weighted average of all the isotopes of that element. Atomic Mass =(%Abundance Isotope#1 / 100)(Mass Isotope#1) + (%Abundance Isotope#2 / 100)(Mass Isotope#2) + … [all isotopes]

19. Atomic Mass (cont.) Example: What is the atomic mass of chlorine? Answer: There are two isotopes of chlorine: Cl-35 and Cl-37 with fractional abundances of 75.53% and 24.47%, respectively. Atomic Mass = (75.53%/100)(35) + (24.47%/100)(37) = 26.4355 + 9.0539 = 35.49 amu NOTE : Typically, the element’s atomic mass shown on the periodic table is close to the mass of the most abundant isotope.

20. Atomic Mass (cont.)  You try one! What is the atomic mass of neon? Neon – 20 has a relative abundance of 90.92% Neon – 21 has a relative abundance of 0.26% Neon – 22 has a relative abundance of 8.82%  Answer: 20.179 amu

21. Changes in the Nucleus  Nuclear Reactions – reactions that occur only in the nucleus of an atom that cause a change in the composition of the nucleus. A change from one element to another may occur.  All three types of radiation – alpha, beta and gamma – are the result of nuclear reactions.  The nucleus is held together by the Strong Nuclear Force (SNF) which is present in each particle in the nucleus. While protons contribute to the SNF in a nucleus, they also repel each other because like charges repel. Neutrons have no charge and as a result contribute more of the SNF in a nucleus. Therefore, neutrons act like a glue to hold a nucleus together.

22. Changes in the Nucleus (cont.)  Pattern of Stability – In the lighter elements, the proton/neutron ratio is generally 1:1. However, as you put more protons in a nucleus you require more neutrons to hold them together. Every element with an atomic number GREATER THAN 83 is unstable and is radioactive.  Nuclei can be unstable for having both too few and too many neutrons. Nuclei with too many neutrons are likely to decay by emitting beta radiation.

23. Radioactive Decay Read AND MEMORIZE Figure 3-29 on Page 114 in your book!!!