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Atomic Structure, Isotopes, And Ions Lecture 3 Ch.2 Suggested HW: 15, 19, 23, 28, 38, 49, 52, 58 Understanding the Nature of Atoms • 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 Matter • An atom is the smallest identifiable unit of an element • The theory that all matter is composed of atoms grew out of two primary laws 1. Law of constant composition 2. Law of conservation of mass Law of Constant Composition The relative amounts of each element in a given substance are always the same, regardless of how the substance was made. Law of Constant Composition • For a molecule, AB: • mass A + mass B = total mass AB • %A + %B = 100 • %A = 𝑚𝑎𝑠𝑠 𝐴 𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝐴𝐵 𝑥 100% • Example: We analyze 1.630 g of CaS and find that it’s 0.906 g Ca. Find the mass of S? Find the mass% of Ca and S? * 55.6% Ca, 44.4% S * This means that all pure CaS in the universe has the same composition as calculated above, regardless of how it was made or where it was found. Law of Conservation of Mass In a chemical reaction, atoms are not created or destroyed, only rearranged. The total mass of substances present before a reaction is equal to the total mass after. Atomic Structure • We have established that matter is comprised of atoms. But what are atoms made of? • In the 1800’s, physicists conducted numerous experiments which revealed that the atom itself is made up of even smaller, more fundamental particles. • The three types of sub-atomic particles that make up the atom are known as: • electrons • protons • neutrons Discovery of the Electron J.J. Thomson’s Cathode Ray Experiment (late 1800’s) • No Electric Field • With Electric Field Applying voltage to a metal cathode produces a beam of particles. The beam can be deflected by electric fields towards a positive pole. Mass of cathode plate does not change during this process. What does this mean? Plum Pudding Model • Atoms are charge neutral. If electrons reside within an atom, then an equivalent number of positive charges must also exist, appropriately named protons. • How do all these charges coexist? • Thomson proposed the very first theoretical model of the atom, the socalled plum pudding model (PPM) shown to the right. • Electrons reside in a sea of uniform positive charge Protons and The Nucleus • Ernest Rutherford sought to test the PPM using the gold foil experiment (below) • A beam of positively charged α-particles were focused on a very thin sheet of gold • Based on the PM model, this beam would pass right through the gold foil. In actuality, the beam was deflected at odd angles, with some α-particles bouncing directly back!! transmitted through cloud Electron cloud THE ATOM α particles Nucleus scattered, repulsed particles True model of the atom is a dense, positively charged, proton-loaded nucleus surrounded by a sparse electron cloud ! The vast majority of an atom’s mass is contained within the nucleus. 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. Subatomic Particles and Their Relative Masses & Charges Particle Relative Charge Mass (amu) Charges shown in table are relative to the charge of a proton. A proton has an actual charge of 1.602 x 10-19 Coulombs (C), an electron has a charge of -1.602 x 10-19 C. Opposite charges attract! Like charges repel!! 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. 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 (the unit of atomic mass is the amu, or atomic mass unit). • These averages are used because numerous variations of elements called isotopes exist in nature. Isotopes • Isotopes are variations of elements with the same number of protons but different numbers of neutrons. Isotope symbols are shown below for the three isotopes of nitrogen with their % abundances in nature. The 14N and 15N isotopes have 7 and 8 neutrons, respectively. mass number atomic number 𝟏𝟒 𝟕𝑵 (99.636%) 𝟏𝟓 𝟕𝑵 (0.346%) Avg. atomic mass is obtained using the % abundance and the isotope mass. Transitional Page 𝐴𝑣𝑒𝑟𝑎𝑔𝑒 𝑎𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 = 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑚𝑎𝑠𝑠 𝑥 (% 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒) Group Work • For the table below, fill in the numbers of protons, neutrons, and electrons for each isotope of carbon. • Then, using the given abundances and isotope masses, calculate the average atomic mass of C. Does it match the reported value? ISOTOPE P N E %A Mass (amu) 𝟏𝟐 𝟔𝑪 98.93 12 𝟏𝟑 𝟔𝑪 1.07 13.003 354 8378 𝟏𝟒 𝟔𝑪 ~0 14.003 2420 Group Work • Boron has two isotopes, 10B and 11B. Using the given isotope masses, determine the % abundances of each isotope. ISOTOPE %A Mass (amu) 𝟏𝟎 𝟓𝑩 10.013 𝟏𝟏 𝟓𝑩 11.009 Proton-Neutron Ratio and Radioactivity • 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. Proton-Neutron Ratio and Radioactivity • 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. Radioactivity • The three most common types of radioactive decay are alpha, beta, and gamma Property α β γ 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) Radioactivity • For example, the 238 92𝑈 isotope undergoes alpha decay to decrease its n/p ratio: 238 92𝑈 → 234 90𝑇ℎ + 42𝐻𝑒 • The Thorium-234 isotope then undergoes beta decay which lowers the ratio even more: 234 90𝑇ℎ → 234 91𝑃𝑎 + 0 −1𝑒 – In beta decay, a neutron is converted to a proton and an electron. This causes the proton count to increase: 1 0𝑛 → 11𝑝 + −10𝑒 • Gamma (γ) decay usually accompanies α or β decay to release residual excess energy. γ is not shown in equations. Applications of Radiochemistry: Carbon Dating 𝟏𝟒 𝟕𝑵 + 𝟏𝟎𝒏 → 𝟏𝟓 𝟕𝑵 𝟏𝟓 𝟕𝑵 → 𝟏𝟒 𝟔𝑪 + 𝟏𝟏𝒑 𝟏𝟒 𝟔𝑪 → 𝟏𝟒 𝟕𝑵 + −𝟏𝟎𝒆 Beta decay! Half life = 5700 yrs Applications of Radiochemistry: Radiomedicine 131I treatment for thyroid cancer Radioactive tracers can be linked to chemical compounds to allow doctors to monitor physiological processes. These compounds ‘glow’ upon decay via γ-decay • Organ malfunction can be indicated if a radioisotope is taken up either too little or too much. • Can monitor blood flow • Intestinal blockages can be detected by accumulation of tracer. • Tumor detection • 18F, 99Tc 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. 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 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 Magnesium ion 𝑶𝟐− Oxide Aluminum ion 𝑷𝟑− Phosphide 𝑴𝒈𝟐+ 𝑨𝒍𝟑+ Group Work • Fill in the missing information below ISOTOPE P N E 13 14 10 32 16𝑆 32 216𝑆 ?? ?? 4+ ??𝑃𝑡 95