Unit 1-The Atom and Periodic Table

1. Unit 1-The Atom and Periodic Table https://www.superteachertools.us/jeopardyx/jeopardy-review-game.php?gamefile=327265#.WmPFVLpFzmI

2. Overview • Chemistry is the study of matter • Matter is anything that has mass and volume • Chemistry deals with matter and the changes it can go through • In order to understand modern chemistry, we must go back to the beginning!

3. Early Studies of Atoms • Democritus (Greek 460-370BC) • Was the first person to use the word “atomos” (a-tomos=in-divisible) • Believed that atoms were indivisible & indestructible • But he had no scientific evidence • Aristotle (Greek 384-322BC) • Rejected atoms • Believed in the 4 core elements • Fire, air, water, earth

4. Greek beliefs dominated until the 16- to 1700s • Robert Boyle (below, 1627-1691) believed gold and silver to be sort of like elements, but he was an alchemist so doesn’t really count • Antoine Lavoisier (1743-1794), father of modern chemistry, identified the first 20 or so elements like gold and silver

5. Lavoisier’s work led to the redefinition of an element as a substance whose particles (atoms) cannot be broken down into simpler substances by physical or chemical means and whose particles are all ‘identical’ • He was also guillotined during the French Revolution

6. John Dalton’s Atomic Theory (1766-1844) • 1. All elements are composed of indivisible atoms • 2. All atoms of a given element are identical • 3. Atoms of different elements are different and have different masses • 4. Compounds are formed by the combination of atoms of different elements

7. The “Plum Pudding” Model (1897-1904) • Proposed by JJ Thomson • Used a cathode ray tube to show atoms have a negatively charged part called the electron • Electrons have a charge of -1 and a mass of nearly 0

8. Ernest Rutherford’s Gold Foil Experiment (1909) • Bombarded gold foil with alpha particles • Particles scattered, showing that atoms have a dense, positively charged nucleus • But wouldn’t identify the proton until 1919 • A proton has a charge of +1 and a mass of 1 atomic mass unit

9. Rutherford’s Nuclear Model • Atom is mostly empty space • Atoms have a dense center or “nucleus” • Most of the mass is in the nucleus • The nucleus is positive • Electrons are outside the nucleus and occupy most of the atom’s volume

10. Limitation of Rutherford’s Model • Rutherford’s Nuclear Model still did not explain the atomic masses of the elements • Thought that the nucleus had additional protons and electrons, so the mass was there but their charges cancelled • In 1932, James Chadwick discovered the neutron with the help of Irene Curie’s experiments • Neutrons have a mass of 1 atomic mass unit, but no charge • So now we know all the pieces, but what does an atom look like?

11. Modern Atomic Theory • Combination of: • The Bohr Model (1900s)-”planetary model” • The Wave-Mechanical Model (1900s) • Niels Bohr got a house full of beer as a reward for his Nobel Prize • In Denmark, he would later dissolve 2 of his friend’s Nobel Prizes in acid

12. The Bohr Model-”planetary model” • Atoms have a nucleus with rings of orbiting electrons • The ring is called an orbital, shell, or level • Outermost shell can’t have more than 8 e-s • Outermost shell is called the valence shell • When a valence shell is filled, the atom is stable and in its lowest energy state

13. Limitations of the Bohr Model • 1. Electrons are not only particles • They actually act like particles AND waves at the same time • 2. Electrons do not actually move in rings around the nucleus

14. The Wave-Mechanical Model • (Bohr helped create this after discovering new evidence with a bit of help from Einstein, Planck, and de Broglie) • Energy and matter act as both waves and particles • e- with distinct amounts of energy move in orbitals • Orbital is region where e- most likely to be, not a ring

15. Subatomic Particles • The nucleus of an atom has • Protons and neutrons • Protons have a charge of +1, neutrons are neutral • One proton = 1 atomic mass unit (amu)=1.67x10-24g • Neutrons approximately the same • # of protons determines the element • # of protons=atomic number

16. Atoms are surrounded by electrons (e-) in the electron cloud • Mass of 1/1836 amu • Charge of -1 • Electrons only exist in orbitals, and can never exist in between them • In a neutral atom, # of electrons = number of protons = atomic number

17. Electron Arrangement • electrons may not weigh much, but they determine: • Chemical properties • Type of bond between atoms • Characteristics determined by the number of valence e-s • The outermost shell

18. Location of Electrons • Review: Electrons are found in orbitals • Orbitals = region an electron can most probably be found • Energy levels • Orbitals form a series of energy levels • Orbitals further away from the nucleus have the most energy • e-s can jump or fall energy levels but can’t exist between them • The outermost or highest energy level are the VALENCE electrons

19. Electron Arrangement • Electron configuration • The locations of all electrons are written as an electron configuration • Each number represents the Number of electrons in that “shell” or level

20. Electron Arrangement – Bohr Model • Examples you may want: • Sulfur • Ca

21. Lewis Dot Structures • Chemical symbol surrounded by dots representing valence electrons • Symbol represents the kernel: nucleus + nonvalence electrons • There are four sides around the symbol, so there are four possible pairs of electrons • Electrons will always spread out before pairing • Cl: 2 – 8 – 7

22. Neutral Practice • Ne • B • Mg • K • Na • Si • S • I

23. Lewis Dot Structures for Ions • Must have symbol and valence electrons surrounded by brackets and charge • Mg2+[Mg]2+ • S2-

24. Ions Practice • For the following, use the periodic table to determine its ion’s charge and draw the Lewis Dot Structure (if multiple charges are listed, use the first one on the list) • F • Se • Ge • Al • Mg • Na • Rb

25. Excited Electrons • Ground state • When e-s occupy the lowest available orbitals • Carbon in the ground state is 2-4 • Excited state • When e-s absorb energy (heat, light, electricity) and move to a higher level • Carbon in the excited state could be 2-3-1 or 1-5, etc. • Energy levels get full in the pattern of 2-8-18-32-18-8 • More commonly will be full in the pattern of 2-8-8 or 2-8-18-8

26. Why does a hamburger have less energy than a steak? • It’s in the ground state!

27. Write the ground state electron configuration for each of the following and write one possible excited state configuration • H • He • Li • P • Fe

28. Write the ground state configurations for the following • Fe2+ • Ca2+ • O2-

29. When an electron moves from ground to excited • Gains energy • Electrons can move from the excited to the ground state • Releases energy in the form of light (a photon) • Infrared, UV, visible • Neon lights

30. Bright line spectra • The result of the energy released when an electron falls to ground state • Each element has unique spectra that can be used to identify the element • It’s unique because of each elements unique electron configuration

31. What’s in the mixture?

32. (Stop here)

33. Atomic Masses • Mass of any one nucleus must be an integer • Atomic mass of a nucleus = # protons + # neutrons

34. Atomic number is 3 • The mass is 6.94 amu • That’s not a whole number…

36. Atomic mass on the periodic table is the • Weighted average of all the naturally occurring isotopes of an element • We will learn how to calculate in a minute, but what’s an isotope?

37. Isotopes • Each atom of the same element has the same # of protons • The number of neutrons may vary • When this happens, called an isotope • The atomic mass also changes • So an atom of hydrogen will always have 1 proton, but may have 0, 1, or even 2 neutrons

38. Isotopic Notation • An atom of hydrogen with 1 proton and 2 neutrons • or 3H or hydrogen-3 or H-3 • It all means the same thing • Examples: • Carbon-12: 6 protons, 6 neutrons=12 amu • Carbon-14: 6 protons, 8 neutrons=14 amu • 12C or 14C or ____________ or _____________?

39. Calculating Atomic Mass • Calculate the atomic mass from the following: • Carbon: 98.89% C-12, 1.108% C-13 • Convert % to decimals • 0.9889 and 0.01108 • Multiply each by the mass of its isotope • 0.9889 x 12 amu = 11.87 amu • 0.01108 x 13 amu = 0.1440 amu • Add the masses together • 11.87amu+0.1440amu=12.01amu • Preferably, set it up all at once: (0.9889x12)+(0.01108x13)

40. Calculating Atomic Mass Worksheet

46. Ions • An ion is a charged atom, but we won’t call it an atom • The result of an atom losing or gaining an electron • Example: If Ca loses 2 electrons, it becomes a Ca+2 ion • One electron has a -1 charge, so it loses -2 total • 0-(-2)=+2

47. Examples – (for now) use the first charge given on the periodic table as the charge of the ion Mg: +2 ion, 12 protons, 12 neutrons, 10 electrons +2 ion so it needs 2 more protons than electrons, but you can’t change the protons, so subtract 2 electrons K Ba P S F O Ar

48. Put on your glasses or hold up the spectroscope – look at the “neon” light

49. Turn off the lights/projector and close the blinds

50. COSMOS EPISODE 5: HIDDEN IN THE LIGHT