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Gilbert Kirss Foster Chapter 2 Atomic Structure Subatomic Particles? • If you cut a piece of graphite from the tip of a pencil into smaller and smaller pieces, how far could you go? • You would eventually end up with atoms (translates to “indivisible” in greek) of pure carbon. • You can not divide a carbon atom into smaller pieces and still have carbon © 2014 W. W. Norton Co., Inc. Atomic Structure • We have established that matter is comprised of atoms. But what are atoms made of? • There are three types of sub-atomic particles that make up the atom are known as: • electrons • protons • neutrons © 2014 W. W. Norton Co., Inc. Setting up the Cathode Ray Exp. • The electron was the first subatomic particle discovered. • In the late 1800’s, J.J Thomson sought to understand the strange “cathode ray” phenomena, which involved the observation of a “strange, flowing energy” through gases. • He developed the cathode ray experiment, which was comprised of: • Glass tube from which most of the air was removed • Two metallic plates, an anode and a cathode, connected to a high voltage power supply. • The cathode is connected to the negative terminal of the power source. The anode has a small hole drilled through its center. © 2014 W. W. Norton Co., Inc. 4 Observation of Cathode Rays • When the connections were made, these mysterious cathode rays flow from the cathode to the anode, and some of these rays escape through the hole in the anode. • The rays are invisible, so a phosphorescent screen lines the back of the tube, which exhibits a glowing spot when struck by the beam. © 2014 W. W. Norton Co., Inc. 5 What Are Cathode Rays? • Thompson soon realized that the cathode rays could be deflected by electric and magnetic fields. The image below shows a cathode ray beam being deflected upwards toward a positive pole. • He also found that the mass of the cathode remained virtually unchanged. What does this mean??? • The beam is not energy, but rather, charged, nearly massless particles, and the particles are negative! © 2014 W. W. Norton Co., Inc. 6 Plum-Pudding Model • Following the discovery of the electron, it became obvious that positive charges, called protons, must also exist since matter is electrically neutral • However, scientist had no idea how these particles were arranged in the atom. The first proposed model was the “plum pudding model”, which described electrons as being spread out in a proton “sea” © 2014 W. W. Norton Co., Inc. 7 Determining the Nuclear Model • Following the discovery of radioactivity, the “gold foil experiment” was designed to test the PP model. • A thin sheet of gold foil was placed in a phosphorescent ring. A radioactive emitter of positive particles was placed in front. • If the PP model was correct, then the positive α-particles would pass through the foil unimpeded. But… • The particles were actually deflected!! How? © 2014 W. W. Norton Co., Inc. 8 Expected Results Expected results from “plum-pudding” model. © 2014 W. W. Norton Co., Inc. 9 Rutherford’s Experiment • The experiment not only disproved the PP model, but also suggested that a very dense, very positive “core” exists at the center of the atom, in which all positive charges are found. • This came to be known as the nucleus. © 2014 W. W. Norton Co., Inc. 10 Neutrons • Rutherford’s model was incomplete. For example, a hydrogen atom has one proton and one electron, but is only ¼th the mass of a helium atom which has two electrons and two protons. • If all of the mass of an atom comes from its sub-atomic particles, how do we explain the unaccounted for mass? • The answer is neutrons, particles that are equal in mass to protons, but with no electrical charge. While scientists knew that neutrons had to exist, they were not officially discovered until 1932. © 2014 W. W. Norton Co., Inc. The Nuclear Atom • Positively charged center of an atom, containing nearly all of the atom’s mass • About 1/10,000 the size of the atom © 2014 W. W. Norton Co., Inc. 12 About the Nucleus • Atomic Mass Units (amu) • Unit used to express the relative masses of atoms and subatomic particles • Equal to 1/12 of a carbon atom © 2014 W. W. Norton Co., Inc. 13 Elemental Symbols 6 Atomic # C Carbon 12.0107 • The number of protons in an atom is called the atomic number. An element is defined by its atomic number. (ex. only carbon has 6 protons) • For a given element, the number of protons DOES NOT CHANGE • In a neutral atom, the number of protons is equal to the number of electrons. © 2014 W. W. Norton Co., Inc. Elemental Symbols 6 C Carbon 12.0107 Mass # • The mass number of an element is the sum of its protons and neutrons. • The mass #’s listed on the periodic table are averages, in units of amu • These averages are used because numerous variations of elements called isotopes exist in nature. © 2014 W. W. Norton Co., Inc. Isotopes • Isotopes are variations of the same element having different numbers of neutrons. Atomic Mass = total number of “nucleons” (protons, neutrons) in the nucleus Atomic Number (Z) = the number of protons A Z X • Isotope symbols are shown below for the two isotopes of nitrogen with their % abundances in nature. The 14N and 15N isotopes have 7 and 8 neutrons, respectively. 𝟏𝟒 𝟕𝑵 © 2014 W. W. Norton Co., Inc. (99.636%) 𝟏𝟓 𝟕𝑵 (0.346%) Group Work • Complete the missing information in the table. 23 ? © 2014 W. W. Norton Co., Inc. 17 Transitional Page Avg. atomic mass is obtained using the % abundance and the isotope mass. 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 = © 2014 W. W. Norton Co., Inc. 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠 𝑥 (% 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒) Group Work • Using the given abundances and isotope masses, calculate the average atomic mass of C. Does it match the reported value? ISOTOPE %A Mass (amu) 𝟏𝟐 𝟔𝑪 𝟏𝟑 𝟔𝑪 𝟏𝟒 𝟔𝑪 98.93 12 1.07 13.003 354 8378 ~0 14.003 2420 • Boron has two isotopes, 10B and 11B. Using the given isotope masses, determine the % abundances of each isotope. Hint: total abundance must equal 100% ISOTOPE © 2014 W. W. Norton Co., Inc. %A Mass (amu) 𝟏𝟎 𝟓𝑩 10.013 𝟏𝟏 𝟓𝑩 11.009 Proton-Neutron Ratio • The nuclei of most naturally occurring isotopes are very stable, despite the massive repulsive forces that exist between the protons in the nucleus. • A strong force of attraction between neutrons and protons known as the nuclear force counteracts this repulsion. • As the number of protons increases, more neutrons are required to stabilize the atom. Stable nuclei up to atomic number 20 have equal numbers of protons and neutrons. • For nuclei with atomic number above 20, the number of neutrons exceeds the protons to create a stable nucleus. © 2014 W. W. Norton Co., Inc. Proton-Neutron Ratio • Radioactive isotopes are unstable (high in energy). This instability is attributed to a neutron/proton ratio that is either too high or too low. • To become stable, they spontaneously release particles or radiation to lower their energy. • This release of energy is called radioactive decay. © 2014 W. W. Norton Co., Inc. Radioactivity • The three most common types of radioactive decay are alpha, beta, and gamma Property α β γ Reason for process Too many protons, too few neutrons (n/p ratio too low) Too few protons, too many neutrons (n/p ratio too high) Too much energy in nucleus Charge 2+ 1- 0 Mass 6.64 x 10-24 g 9.11 x 10-28 g 0 Emitted Radiation Type 2 protons and 2 neutrons ( 42𝐻𝑒) High energy electron. Pure energy (Radiation) Penetrating Power Low. Stopped by paper. Blocked by skin. Moderate. Stopped by aluminum foil. (10α) High. Can penetrate several inches of lead. (10000α) © 2014 W. W. Norton Co., Inc. Radioactive Decay • For example, the 238 92𝑈 isotope undergoes alpha decay to increase its n/p ratio: 238 92𝑈 238 − 92 𝑛 = 1.58 𝑛/𝑝 92 𝑝 © 2014 W. W. Norton Co., Inc. → 234 90𝑇ℎ + 42𝐻𝑒 234 − 90 𝑛 = 1.60 𝑛/𝑝 90 𝑝 Radioactive Decay • The Thorium-234 isotope undergoes beta decay which lowers the n/p ratio: 234 90𝑇ℎ → 234 − 90 𝑛 = 1.60 𝑛/𝑝 90 𝑝 234 91𝑃𝑎 + 0 −1𝑒 234 − 91 𝑛 = 1.57 𝑛/𝑝 91 𝑝 • In beta decay, a neutron is converted to a proton and an electron. This causes the proton count to increase: 1 0𝑛 © 2014 W. W. Norton Co., Inc. → 11𝑝 + −10𝑒 Ions • Thus far, we’ve learned that each element has an exact number of protons. • For example, Hydrogen has only one proton. If you force a second proton onto the atom, you no longer have hydrogen… you now have Helium. • We have also learned that atoms can have variable numbers of neutrons (isotopes). • Next, we will discuss ions. © 2014 W. W. Norton Co., Inc. Ions • Ions are electrically charged atoms, resulting from the gain or loss of electrons. • Positively charged ions are called cations. You form cations when electrons are lost • Negatively charged ions are called anions. You form anions when electrons are gained © 2014 W. W. Norton Co., Inc. Ion Nomenclature • A cation is named by adding the word “ion” to the end of the element name • Anions are named by adding the suffix –ide to the end of an element 𝑳𝒊+ Lithium ion 𝑪𝒍− Chloride 𝑵𝒂+ Sodium ion 𝑺𝟐− Sulfide 𝑴𝒈𝟐+ Magnesium ion 𝑶𝟐− Oxide 𝑨𝒍𝟑+ Aluminum ion 𝑷𝟑− Phosphide © 2014 W. W. Norton Co., Inc. Group Work • Fill in the missing information below ISOTOPE P N E 13 14 10 32 16𝑆 32 216𝑆 ?? ?? 4+ ??𝑃𝑡 © 2014 W. W. Norton Co., Inc. 95