Atoms and the Periodic Table

1. Atoms and the Periodic Table

2. Atomic Structure • ATOM: the smallest particle that has the properties of an element. • From the early Greek concept of the atom to the modern atomic theory, scientists have built on and modified existing models of the atom.

3. ATOM BASICS • Atoms are composed of a positively charged nucleus surrounded by an electron cloud. • Nucleus (99% of atom’s mass): uncharged neutrons and positively charged protons. • Electron cloud: negatively charged electrons in constant motion creating a “cloud” like a fan.

4. DEMOCRITUS • In 400 B.C., this Greek philosopher suggested that the universe was made of indivisible units. • “Atom” – Greek word meaning “unable to be divided”

5. JOHN DALTON In 1808, this English schoolteacher proposed his own atomic theory that became widely accepted. Dalton’s Atomic Theory: • All elements are made of tiny atoms. • Atoms cannot be subdivided. • Atoms of the same element are exactly alike. • Atoms of different elements can join to form molecules.

6. As it turns out, the atom can be divided into subatomic particles. • Thompson and Millikan are given credit for the first discoveries relating to electrons. • Rutherford discovered the positively charged nucleus Millikan Thompson Ernest Rutherford

7. Rutherford Scattering Experiment

8. Rutherford Scattering Experiment

9. Rutherford Scattering Experiment

10. “I remember Geiger coming to me in great excitement and saying, “We have been able to get some of the alpha-particles coming backwards.” It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15 inch shell at a piece of paper and it came back and hit you.” Ernest Rutherford

11. Rutherford Scattering Experiment Rutherford proposes the nucleus, and the “solar system” model.

12. NIELS BOHR • In 1913, this Danish scientist suggested that electrons “orbit” the nucleus. • In Bohr’s model, electrons are placed in different energy levels based on their distance from the nucleus.

13. Electrons are like books in a book case. They can only exist on certain levels.

14. Only certain orbits are allowed

15. Only certain orbits are allowed An electron jumping up absorbs light

16. Only certain orbits are allowed An electron jumping down gives off light

17. Orbits – Energy Levels n = 4 n = 3 n = 2 n = 1

18. Orbits – Energy Levels n = 4 n = 3 n = 2 n = 1

19. Orbits – Energy Levels n = 4 n = 3 n = 2 n = 1 Lyman Series (Ultraviolet)

20. Orbits – Energy Levels n = 4 n = 3 n = 2 Balmer Series (Visible) n = 1

22. Rydberg Equation

23. Bohr Model of the Atom Successful: 1. Explained the existence of line spectra. 2. Explained the frequencies of line spectra.

24. The Bohr Atom was a “Solar System” model.

25. Bohr Model of the Atom Defects: 1. Contradicted known physics. 2. Couldn’t explain various intensities of the line spectra. 3. Couldn’t explain why only certain orbits were allowed. 4. Only worked for hydrogen.

26. Physicists were mystified, but intrigued by Bohr’s theory of the atom. Why are the energies of the hydrogen electron quantized? Why is the electron restricted to orbiting at certain fixed distances?

27. Louis de Broglie For a decade no one, not even Bohr himself had a logical explanation. In 1924 Louis de Broglie provided a solution.

28. In his doctoral dissertation he proposed that matter and radiation have both wave and particle properties. He won the nobel prize in 1929

29. If electrons are waves, what is it that’s waving? Einstein proposes it’s the probability of finding an electron at any given place.

30. MODERN ATOMIC MODEL • By 1925, Bohr’s model of the atom no longer explained all observations. Bohr was correct about energy levels, but wrong about electron movement. • Electrons occupy the lowest energy levels available. • Energy increases as distance from the nucleus increases. • Electrons move in patterns of “wave functions” around the nucleus. • It is impossible to know an electrons velocity and location at any moment in time.

31. This is very difficult stuff! Schrödinger In 1926 Schrodinger proposes an equation that gives the probability of finding an electron at any place in the atom. Don’t worry, this won’t be on the test.

32. ORBITALS • ORBITAL: the regions in an atom where there is a high probability of finding electrons. • s is the lowest energy orbital, and p is slightly higher • d and f are the next two orbitals. They occupy even higher energy levels and take on more complex shapes than s & p s orbital p orbitals

33. VALENCE ELECTRONS Carbon 4 valence electrons • Electrons in the outermost energy level are called VALENCE ELECTRONS. • Valence electrons determine how an atom will act in a chemical reaction. • Atoms with equal numbers of valence electrons have similar properties. • We will learn how to determine the # of valence electrons in an atom later in this unit.

34. What part of the atom is much smaller than the atom, yet contains most of the mass? • the nucleus • the electron cloud Not drawn to scale While its diameter is very small compared to that of the entire atom, 99% of the mass of an atom comes from the protons and neutrons in the nucleus.

35. Which statement is true according to Dalton’s atomic theory? • Atoms of different elements join to form larger atoms. • Atoms can be subdivided into smaller particles. • Atoms of the same element differ in electric charge. • Atoms of the same element are exactly alike. This statement was part of Dalton’s atomic theory. While this statement was holds true if we remove the word “exactly”, we have since found that atoms of an element often differ in number of neutrons (isotopes).

36. According to Bohr’s model of the atom, electrons behave like • planets rotating on their axes. • planets orbiting the sun. • light energy in a vacuum. • waves vibrating on a string. While Bohr’s model of energy levels of electrons around an atom remains supported, his model of electron movement has been revised. Schrödinger and other scientists have since found that electrons travel in wave-like patterns around the nucleus. Schrödinger

37. Dmitri Mendeleev: 1834-1907 1869: created first periodic table of elements. Arranged elements in order of increasing atomic mass.

38. Henry Moseley 1887-1915 One of Rutherford’s students. 1914: Arranged the elements in order of increasing atomic number (responsible for TODAY’S periodic table).

39. Organization of the Periodic Table • PERIODICITY: regular variations (or patterns) of properties with increasing atomic number. Both chemical and physical properties vary in a periodic (repeating) pattern. • PERIOD: horizontal row of elements on P.T. • GROUP (FAMILY): vertical column of elements on P.T.

40. Periodic Key 6 C Carbon 12.011 Atomic number (Z) Element’s symbol Element’s name Atomic mass (A) # of protons = Z # of electrons = # of protons (in a neutral atom) # of neutrons = A-Z

41. ISOTOPES • Isotopes are atoms that have the same # of protons, but a different # of neutrons. • Example: Carbon-12 vs. Carbon-14 12C Mass # = 12; Atomic # = 6 (6P, 6E, 6N) 14C Mass # = 14; Atomic # = 6 (6P, 6E, 8N)

42. IONS • Ionization: the process of adding or removing electrons from an atom or group of atoms. • An ion has a net electric charge. • Cation: ion with a positive charge. Ex: Na+ • Anion: ion with a negative charge. Ex: O2-

43. Valence electrons & electron dot diagrams • Review: The valence electrons are the outermost electrons in an atom. These are the electrons that are involved when there is a chemical reaction. • Looking at a periodic table, you can quickly determine how many valence electrons an atom has by what column it is in.

44. # of Valence electrons 2 1 2 3 4 5 6 7 8

45. ELECTRON DOT DIAGRAMS:visual representations of elements and their valence electrons Standard form: 3 6 R 4 1 7 2 5 8 Order of electron/dot placement element symbol Example: O 6 valence electrons Oxygen

46. Determining # P+, N, and E- from chemical symbols: Mass # • Example 1: atom # protons = 6 #electrons= 6 #neutrons= 14-6 = 8 • Example 2:ion # protons = 7 #electrons= 10 #neutrons= 15-7 = 8 No net charge 14 C 6 Atomic # 15 N3- 7 Net charge of -3

47. The periodic law states that when elements are arranged in order of increasing __________ _________, similarities in their properties occur in a regular pattern. • Atomic mass • Atomic number • Atomic radius Moseley created the modern periodic table when he determined that elements should be placed in order of increasing atomic number (# of protons). It’s a shame that WWI took the life of this brilliant scientist when he was so young.

48. Carbon-12 and carbon-14 are • isomers • isotopes • radioactive elements • different elements Isotopes are atoms of the same element with different numbers of neutrons, and therefore different atomic masses. While carbon-14 is used in radioactive dating, carbon-12 has a more stable nucleus and therefore is not used in this capacity.

49. How many protons, electrons and neutrons are in one atom of oxygen-17? • 17 p, 17 e, 17 n • 17 p, 17 e, 1 n • 8 p, 8 e, 8 n • 8 p, 8 e, 9 n While most oxygen atoms have a mass of 16 g/mol, oxygen-17 is an isotope with a mass of 17 g/mol. The number of protons in an element is the same for every atom of that element. # electrons = # protons if element has no charge.

50. How many protons, electrons and neutrons are in O2-? • 10 p, 8 e, 8 n • 8 p, 8 e, 8 n • 8 p, 10 e, 8 n • 8 p, 6 e, 8 n Every atom or ion of oxygen has 8 protons. The net charge of -2 indicates that the ion has two more negative charges (electrons) than positive charges (protons) All atoms and ions of standard oxygen-16 have 8 neutrons.