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Unit 2 Atomic Structure & Nuclear Chemistry Section 1 Obj.2 Counting Subatomic Particles The ATOM is defined as the smallest particle of an element that retains the properties of that element. Subatomic Particles & Their Properties Particle Symbol Location Relative Mass Actual mass in Electrical Charge in a.m.u* grams 1 amu +1 1.6710-24 g + p 1 amu 1.6710-24 g 0 .00055 amu 9.1110-28 g -1 n0 e*a.m.u : atomic mass unit 1 amu (“atomic mass unit”) = 1.67 10-24 g **The mass of the proton and neutron are very similar** **The mass of the proton & neutron are both about 2000 times larger than the mass of the electron** What are the 2 regions of the atom as of now? Nucleus: small, dense center containing ___________________ and _____________________. Electron Cloud: region surrounding nucleus that contains the electrons. It’s mostly empty space. Counting Subatomic Particles Atomic Number Every atom has a different number of protons, which is represented by the atomic number. The number of protons determines the identity of the atom Atomic Number = Number of protons Ex: Nitrogen’s atomic number is 7 Calculating Electrons in a Neutral Atom (A.P.E) A neutral atom has the SAME number of protons and electrons. # protons = # electrons if neutral Mass Number The number of protons and neutrons in a nucleus of a specific atom is called the mass number. Calculating Electrons in a Ion 2 Options: [1] Add the number of given protons & neutrons together Mass # = protons + neutrons [2] USE ONLY when no element information has been given: Round atomic mass to a whole number to get an element’s most likely most common mass number. o Nitrogen’s atomic mass = 14.01 amu o Nitrogen’s mass number is ____________ Calculating Neutrons (M.A.N) To calculate the number of neutrons, subtract the atomic number from the mass number Mass number – Atomic number = # Neutrons Nitrogen: 14 – 7 = 7 neutron An ion is an atom with a charge. The number of protons is NOT EQUAL to the number of electrons. Positive charged ions called CATIONS have more protons than electrons & negatively charged ions are called ANIONS have more electrons than protons. Example: [1] How many electrons does Br-1 have? [2] How many electrons does Al+3 have? Overall Charge = Protons – Electrons Element Information: Nuclear Symbol LABEL THE PARTS BELOW A A Z Z X X C C SELF CHECK What are the number of protons, neutrons & electrons in each atom? Hyphen Notation: LABEL THE PARTS! Copper – 65 ___________ __ p __e __ n You Try: (Be Careful: Never use the given atomic mas on the periodic table to calculate mass number unless no other information has been provided. ) Nuclear Symbol Hyphen Notation Atomic # Magnesium-25 Mass # Charge Proton Neutron Electron 126 82 +2 82 Section 2 Obj.2 cont. & Obj. 4 How Do Atoms Differ: Isotopes Isotopes are atoms of the same element that have the SAME NUMBER OF PROTONS but a DIFFERENT NUMBER OF NEUTRONS Most elements contain a mixture of 2 or more isotopes. Each one having its own mass and abundance. Some isotopes are radioactive and unstable. (See nuclear chem. Section) You can identify an isotope by the different mass number of the same element Carbon-12 Carbon-13 12 13 C 6 C 6 Fill in the chart using the diagram in the presentation Atomic Isotope Protons number Neutrons Electrons Mass (a.m.u) Lithium-6 Lithium-7 Lithium-8 You Try ? What the number of protons, neutrons and electrons in each? 19 18 F 9 203 F 9 Hg 80 194 Hg 80 How to calculate Average Atomic Mass? Average atomic mass is the weighted average of the masses of all naturally occurring isotopes. Average Atomic Mass Equation: Average atomic mass = (% abundance of isotope x mass of 1st isotope) + (% abundance of isotope x mass of 2nd isotope) + ……… Cl-35 and Cl-37 Why is the AAM = 35.45 amu ? Example 1: Element x has 2 natural isotopes. Calculate its average atomic mass. 1st isotope has a mass of 10.012 a.m.u. with 19.91% abundance. 80.09% of the 2nd element has a mass of 11.009 a.m.u. Example 2: Calculate the average atomic mass of copper if it has 2 isotopes. 69.11% has a mass of 62.93 a.m.u and the rest has a mass of 64.93 a.m.u. Section 3 Obj.3 Nuclear Decay Reactions Nuclear chemistry is the study of the changes of the NUCLEUS of an atom. Nuclear Reactions involve changes within the nucleus where as chemical reactions involve the loss, gain or sharing of electrons. The Nucleus Contains ________________ and neutrons. They are collectively called NUCLEONS. Radioactivity An unstable nucleus will decay or break down, releasing particles and/or energy in order to become stable. An atom with an unstable nucleus is considered “_____________________”. Transmutation Type of nuclear reaction that will change the number of PROTONS and thus will create a new & different ELEMENT There are several ways radioactive atoms can decay into different atoms! Basic Types of Radioactive Decay Particle Type Symbol Alpha Beta *** or or What Happens? Example Penetrating Power Atomic number decreases by ____ and mass number decreases by ____ (1) LOW: Can be blocked by paper/clothing Atomic number increases by _____ but the mass number stays the same (100) MEDIUM: can penetrate the skin; need to be protected by clothing/thin metals like aluminum ***A neutron becomes a ________________ and a high-speed electron that is discharged from the nucleus** Gamma (100000) HIGH: need to be protected by thick concrete or metal like lead No change in atomic nor mass number; occurs with other types of decay Writing Balanced Nuclear Equations The sums of both the ATOMIC NUMBER & MASS NUMBER on the left side must equal the sums of both the ATOMIC NUMBER & MASS NUMBER. Often problems will have 1 particle missing and you will need to identify it. Alpha Decay of thorium-230 You Try! 1. Beta decay of zircomium-97 Beta Decay of cesium -137 3. Complete this: 2. Alpha decay of americium-241 Making New Elements and Isotopes: Bombarding the Nucleus All the transuranium elements (elements with atomic numbers higher than Uranium) have been made by bombarding the nucleus with neutrons and other atoms in accelerators Example: Section 4 Obj.3 cont. Half Life Radioactive isotopes decay at a characteristic rate measured in HALF LIVES. A half-life time (THL) is the time required for HALF of the amount of radioactive atoms to decay. The time ranges from __________________ to millions of _____________________. HOW TO’s 1. To calculate the number of half-lives, divide the half-life (THL) into the total time (T). Then cut the original amount in half the number of time determined by the # of half lives (HL). T / THL = # HL 2. Algebraic equation to calculate the remaining amount left over after a certain number of half-lives have passed. Amount remaining = (initial amount) (.5) #HL Examples: 1. Suppose you have 20 grams of sodium-24. Its half-life is 15 hours. How much is left over after 60 hours. 2. Uranium-238 has a half-life of 4.46 x 109 years. How long will it take for 7/8th of the sample to decay? You Try! 1. 1.5 grams of a 12.0 g sample are left after 114 s. What is the half-life of radium-222? 2. A sample of 3x107 Radon atoms is trapped in a basement that is sealed. The half-life of Radon is 3.83 days. How many radon atoms are left after 31 days?