Atoms and the Periodic Table

1. Atoms and the Periodic Table 2.1 Atoms First 2.2 Subatomic Particles and Atomic Structure Discovery of the Electron Radioactivity The Proton and the Nuclear Model of the Atom The Neutron 2.3 Atomic Number, Mass Number, and Isotopes 2.4 Average Atomic Mass 2.5 The Periodic Table 2.6 The Mole and Molar Mass The Mole Molar Mass Interconverting Mass, Moles, and Numbers of Atoms 2

2. Dalton’s Atomic Theory 1808 1. Elements are composed of atoms • Atoms are indivisible • Atoms of a given element are identical • Neither created nor destroyed. Are atoms of the same element identical? A = true B = false

3. An atom is the smallest quantity of matter that still retains the properties of matter. An element is a substance that cannot be broken down into two or more simpler substances by any means. • Examples: gold, oxygen, helium Atoms can be divided smaller and smaller until eventually only a single atom remains. Dividing it any smaller would give pieces that are no longer an atom. 2.1

4. Dalton’s Atomic Theory 1808 1. Elements are composed of atoms • Atoms are indivisible • Atoms of a given element are identical • Neither created nor destroyed. 3. Compounds are combinations of atoms in fixed proportions 4. Chemical Reaction = atom rearrangement - Explains Law of conservation of mass.

5. Once a single atom has been obtained, dividing it smaller produces subatomic particles. In the late 1800’s, many scientists used a cathode ray tube, which consists of two metal plates sealed inside a glass tube from which most of the air has been evacuated.

6. J. J. Thomson (1856-1940) noted the rays were repelled by a plate bearing a negative charge, and attracted to a plate bearing a positive charge. By varying the electric field and measuring the degree of deflection of cathode rays, Thomson determined the charge-to-mass ratio of electrons to be -1.76×108C/g. (C is coulomb, the derived SI unit of electric charge.) F = ma (1687) F = q1q2/r2(1785) q1/m = r2a/q2 Why couldn’t he know mass or charge individually?

7. Subatomic Particles and Atomic Structure R. A. Millikan (1868-1953) determined the charge on an electron by examining the motion of tiny oil drops. The charge was determined to be -1.6022×10-19 C. The charge-to-mass ratio of electrons to be -1.76×108 C/g What is the mass of an electron? Apply dimensional analysis _q_ = q x m = q/m q

8. The atom according to ……… Dalton ~ 1808 Thomson ~ 1890 Indivisible sphere raisin pudding?

9. Subatomic Particles and Atomic Structure Alpha (α) rays consist of positively charged particles, called α particles. Beta (β) rays, or β particles, are electrons so they are deflected away from the negatively charged plate. Gamma (γ) rays, like X-rays, have no charge and are unaffected by external electric or magnetic fields.

10. Subatomic Particles and Atomic Structure Ernest Rutherford used α particles to prove the structure of atoms. The majority of particles penetrated the gold foil undeflected. Sometimes, αparticles were deflected at a large angle. Sometimes, α particles bounced back in the direction from which they had come.

11. The subatomic particle discovered by the apparatus shown is …. a) proton b) neutron c) electron d) x-ray Millikan’s oil drop experiment directly determined …. a) the charge-to-mass ratio of the electron b) the charge of the electron c) the mass of the electron d) all of the above

12. Rutherford’s experiment…. a) discovered the proton b) discovered the neutron c) discovered a particles d) led to the development of the nuclear model of the atom.

13. Subatomic Particles and Atomic Structure Rutherford proposed a new model for the atom: Positive charge is concentrated in the nucleus. The nucleus accounts for most of an atom’s mass and is an extremely dense central core within the atom. • A typical atomic radius is about 100 pm • A typical nucleus has a radius of about 5×10–3 pm • 1 pm = 1×10–12 m

14. Subatomic Particles and Atomic Structure Protons are positively charged particles found in the nucleus. Neutrons are electronically neutral particles found in the nucleus. Neutrons are slightly larger than protons. Mass – 2 sig. fig. amu 0.0 1.0 1.0

15. Atomic Number, Mass Number, and Isotopes All atoms can be identified by the number of protons and neutrons they contain. The atomic number (Z)is the number of protons in the nucleus. • Atoms are neutral, so it’s also the number of electrons. • Protons determine the identity of an element. For example, nitrogen’s atomic number is 7, so every nitrogen has 7 protons. The mass number (A)is the total number of protons and neutrons. • Protons and neutrons are collectively referred to as nucleons. Mass number (number of protons + neutrons) 2.3 Elemental symbol Atomic number (number of protons)

16. Atomic Number, Mass Number, and Isotopes Most elements have two or more isotopes, atoms that have the same atomic number (Z) but different mass numbers (A). 1 proton 0 neutrons 1 proton 1 neutron 1 proton 2 neutrons Isotopes of the same element exhibit similar chemical properties, forming the same types of compounds and displaying similar reactivities. Heavy water = D2O or 2H2O

17. Average Atomic Mass Atomic mass unit (amu) by assignment ….. 1 12C atom has a mass of 12 amu The average atomic masson the periodic table represents the average mass of the naturally occurring mixture of isotopes. Average mass (C) = (0.9893)(12.00000 amu) + (0.0107)(13.003355 amu) 2.4 = 12.01 amu

18. Natural Abundance of Isotopes Data from CRC Handbook H 1H 99.985% 1.007825 1.0079 2H 0.015% 2.0140 3H trace 3.01605 He 3He 0.00013% 3.01603 4.0026 4He 99.99987% 4.00260 Li 6Li 7.5% 6.01512 6.941 7Li 92.5% 7.01600 B 10B 20% 10.0129 10.811 11B 80% 11.00931 C 12C 98.89% 12.0000 12.011 13C 1.11% 13.00335 N 14N 99.64% 14.00307 14.007 15N 0.36% 15.00011 O 16O 99.76% 15.99491 15.999 17O 0.04% 18O 0.20% Cl35Cl 75.77% 34.96885 35.453 37Cl 24.23% 36.96993 Verify the AW of B? a) 10.013 b) 10.810 c) 10.764 d) 10.985

19. Natural Abundance of Isotopes Data from CRC Handbook H 1H 99.985% 1.007825 1.0079 2H 0.015% 2.0140 3H trace 3.01605 He 3He 0.00013% 3.01603 4.0026 4He 99.99987% 4.00260 Li 6Li 7.5% 6.01512 6.941 7Li 92.5% 7.01600 B 10B 20% 10.0129 10.811 11B 80% 11.00931 C 12C 98.89% 12.0000 12.011 13C 1.11% 13.00335 N 14N 99.64% 14.00307 14.007 15N 0.36% 15.00011 O 16O 99.76% 15.99491 15.999 17O 0.04% 18O 0.20% Cl35Cl 75.77% 34.96885 35.453 37Cl 24.23% 36.96993 Determine “fraction” natural abundance from AW and ‘2’ isotopes.

20. What is a mole? 1. mole : a pigmented spot, mark, or small permanent protuberance on the human body. 2. mole : any of numerous burrowing insectivores with minute eyes, concealed ears, and soft fur. (one who works in the dark) (a spy)

21. What is a mole? 4. A massive work formed of masonry and large stones or earth laid in the sea as a pier or breakwater. 6. A spicy sauce made with chilies and usually chocolate and served with meat. (molé) 5. The base unit of amount of pure substance in the International System of Units that contains the same number of elementary entities as there are atoms in exactly 12 grams of the isotope carbon-12. 1 mole= the numberof C atoms in 12 g of pure 12C. 1 mole = 6.022 x 1023 of any item (Avogadro’s Number) A mole is like a dozen (only a lot more more)

22. ↑ group/ family ↓ element – Matter comprised of a single type of atom. What are the units You should attach To the AW? ← period →

23. A penny weighs 2.509g. 2.5% of that is Cu. AW of Cu = 63.55 amu NA = 6.022 x 1023 How many atoms of Cu are found in a penny? 0.025 0.1574 g Cu 0.002476 moles Cu 1.5 x 1021 atoms Cu __2.5 g Cu__ • 2.509 g penny • 1 mole Cu • 6.022 x 1023 atoms Cu = 100 g penny 63.55 g Cu 1 mole Cu

24. Formula weight (or atomic weight/molecular weight) is the conversion factor from mass to moles, and vice versa. Avogadro’s # converts from moles to # or particles. Grams to moles to # : X grams • 1 mole • 6.022 x 1023= # atoms/molecules/formula units etc. FW (g) 1 mole How many grams of iron pellets do you need to make 2.53 moles of Iron? How many atoms does this contain? Fe AW = 55.847 g/mol