ATOMS & THE PERIODIC TABLE

1. ATOMS & THE PERIODIC TABLE

2. Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it.

3. Atoms and the Periodic Table

4. Size of Atoms Particle Diameter (meters) atom 10-10 nucleus 10-14 proton 10-15 neutron 10-15 electron 10-18 1 Angstrom = 10-10 m

5. History of the Atomic Model Democritus – 400 BC atoms make up all substances John Dalton – 1766-1844 atom is a solid hard sphere Joseph John Thomson – 1856-1940 discovered the electron in 1897 plum pudding model of atom positive sphere with negative e- embedded

6. Lived from (1766-1844) • All elements are composed of atoms • Atoms of the same element are identical. Each element has unique properties . • Atoms of different elements can be chemically combined in simple whole number ratios to form compounds. • Law of Conservation of Matter/Mass

7. The Electron • Streams of negatively charged particles were found from cathode tubes. • J. J. Thompson is credited with their discovery (1897).

8. The Atom, circa 1900: • “Plum pudding” model, put forward by Thompson. • Positive sphere of matter with negative electrons imbedded in it.

9. Ernest Rutherford – 1871-1937 gold foil experiment most of atom empty space positive nucleus contains most of the mass discovered protons in 1919

10. James Chadwick – 1891-1974 discovered the neutron in 1932. Niels Bohr – 1885-1962 (Danish) electrons move around the nucleus in fixed orbits that have a set amount of energy

11. Electron Cloud Model of Atom 1. came to be used to estimate the positions of electrons in an atom 2. uncertainty principle, which states that it is not possible to obtain precise values of both position and momentum of a particle at the same time 3. probability of finding an electron

12. Protons • The number of protons distinguished 1 atom from another • Most atoms are very stable • It takes a lot of energy to add or remove a proton from an atom • Atomic number = number of protons • The Periodic Table is arranged by number of protons

13. Symbols of Elements Elements are symbolized by one or two letters.

14. Atomic Number Number of protons = The atomic number

15. Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

16. Mass Number • The number of protons plus neutrons in an atom. • Always a whole number. • Written or indicated like this: H C Cu Si K

17. Calculating Number of Neutrons • Subtract: Mass Number - Atomic Number = Neutrons Mass Number - # of protons = Neutrons

18. Notes on Finding Numbers of Protons, Neutrons, Electrons

19. Isotopes • Atoms of the same element with different numbers of neutrons

20. Isotopes • To distinguish between isotopes of an element • Ex: Neon has 3 main isotopes

21. Average Atomic Mass • Atomic mass unit – 1/12 the mass of a C-12 atom • To calculate avg. atomic mass

22. Example – Avg. Atomic Mass

23. Chemical Goals • To be Chemically Stable Unstable atoms are radioactive: their nuclei change or decay by spitting out radiation, in the form of particles or electromagnetic waves. • To be Electronically neutral To have no charge on the atom. To have the same # of protons as electrons.

24. Why does an atom stay together? The strong nuclear force keeps protons and neutrons together in the nucleus in spite of the repulsion of the protons for each other. The strong nuclear force acts only over a very short distance. It doesn’t work outside the nucleus. The strong nuclear force is stronger than the electromagnetic force.

25. Valence Electrons • Electrons in outer energy level. • Can only have 8 or less. This is the Octet Rule. • These electrons are the ones involved in bonding with other atoms and the ones with the most energy. • Looking at the Periodic Table, you can tell the number of valence electrons for the A Families from the Roman Numeral designation

26. Electrons • e- located far from nucleus in a series of energy levels. • Each e- has a certain amount of energy. 3. The further the e- gets from the nucleus the more energy it has. Valence e- have the most energy. Those closest to the nucleus have the least amount of energy.

27. Energy Levels 1. Each energy level can only hold a certain number of e- 2. Electrons always fill low energy orbitals (closest to the nucleus) before filling higher energy ones. 3.The high the energy level occupied by the e-, the easier it is for the e- to escape from the atom. 4. Quantum of energy – amount of energy needed to move an e- from its present energy level to the next one 5. Ground State – the lowest energy level for ane-.

29. Electron Placement on the Periodic Table 2 e- 8 e- 18 e- 32 e-

31. Electron Configuration Chart

32. Aufbau /Orbital Diagram

33. EX: Electron Configuration for Potassium 19 electrons EX: Electron Configuration for Arsenic 33 electrons EX: Electron Configuration for Silver 47 electrons

34. Pauli Exclusion Principle An orbital can hold 0, 1, or 2 electrons and if there are 2 electrons in the orbital they must have opposite spin.

35. Aufbau Principle • Rules for Orbital Filling • Lower-energy orbitals fill first. • An orbital can hold only 2 e- with opposite spins. • The most stable arrangement for e- is one with the maximum number of unpaired e-. This minimized e- to e- repulsion and stabilizes the atom. – Hund’s Rule

36. Hund’s Rule • When filling up sublevels other than s, electrons are placed in individual orbitals before they are paired up.

37. Increasing Energy in Electron Sublevels

38. THE DIAGONAL RULE MUST GO IN THIS ORDER: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d and 7p. These orbitals will account for all the elements now known. This diagonal rule can help account for the octet rule, too.

39. The Energy Flow • The order of increasing energy of the orbitals is then read off by following these arrows, starting at the top of the first line and then proceeding on to the second, third, fourth lines, and so on. This diagram predicts the following order of increasing energy for atomic orbitals. • 1s < 2s < 2p < 3s < 3p <4s < 3d <4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p < 8s ...

40. Bohr Model of Atomsonly represents energy levels, not orbitals Lithium

41. Nobel Gases Neon Krypton Argon

42. Drawing AtomsThe Bohr Model 11 p+ e- 3 p+

44. Lewis StructuresElectron Dot Diagrams • Describes e- arrangement in atoms • Describes bond arrangement in molecules. • Uses dots to represent valence e- around an atom • EX: Li Ne O Si

45. Dimitri Mendeleev • In the late 1860's, Mendeleev began working on his great achievement: the periodic table of the elements. By arranging all of the 63 elements then known by their atomic weights, he managed to organize them into groups possessing similar properties. Where a gap existed in the table, he predicted a new element would one day be found and deduced its properties. And he was right. Three of those elements were found during his lifetime: gallium, scandium, and germanium.

46. Mendeleev’s Periodic Table

47. Moseley’s Periodic Table In 1913 Henry Moseley came up with this Periodic Table. The elements are arranged by increasing atomic number.

48. Groups • A group (also known as a family) is a vertical column in the periodic table of the chemical elements. There are 18 groups in the standard periodic table. • Elements in a group have similar configurations of the outermost electron of their atoms – same number of valence e- • This gives the groups of elements similar physical and chemical characteristics.

49. Periods • With each group across a period, the elements have one more proton and electron and become less metallic. • Rows of elements are called periods. The period number of an element signifies the highest unexcited energy level for an electron in that element.

50. Physical Properties Metals Nonmetals Poor conductors of heat and electricity. Brittle - if a solid. Nonductile. Do not possess metallic luster. Transparent as a thin sheet. Solids, liquids or gases at room temperature. • Good electrical conductors and heat conductors. • Malleable - can be beaten into thin sheets. • Ductile - can be stretched into wire. • Possess metallic luster. • Opaque as thin sheet. • Solid at room temperature (except Hg).