Atoms

1. Atoms And what they form…

2. Element Information • Webelements… http://www.webelements.com

3. The Atom • ~400 BCE Democritus: small particle (“atomos” is Greek for indivisible) • 1803 John Dalton: atomic model with several postulates • 1897-now Modern Investigations…

4. Dalton’s Atomic Theory The Postulates • All matter consists of atoms which are indivisible and indestructible. 2. Atoms of one element cannot be converted into atoms of another element. 3. Atoms of an element are identical in mass and other properties and are different from atoms of any other element. 4. Compounds result from the chemical combination of a specific ratio of atoms of different elements. From Silberberg, Principles of Chemistry

5. Atomic Structure • ~1870-1935 saw experimentation that showed that Dalton was right and wrong. • 1874 – Stoney: electricity is made of individual particles with negative charge called electrons • 1879 – Crookes: discovered “cathode rays” have distinct properties like electrons

6. Figure 2.4 Silberberg, Principles of Chemistry Experiments to determine the properties of cathode rays.

7. Atomic Structure • 1896 – Becquerel discovers emissions from materials (radioactivity) • 1898 – Rutherford uncovers properties of the some Becquerel emissions and names them alpha and beta www.lbl.gov/abc/graphics/magnet.gif

8. Atomic Structure • 1897 – Thomson discovers “canal rays” are the same as positive hydrogen atoms • Model of atom is “plum pudding” with + and – particles mixed together http://commons.wikimedia.org/wiki/File:Plum_pudding_atom.svg • 1911 – Rutherford conducts “gold foil experiment”

9. Rutherford’s Gold Foil Experiment

11. Atomic Structure • After gold foil experiment, atomic model changes to one with + charges (protons) in dense center with – charges (electrons) surrounding the center • 1932 – Chadwick: discovers the missing mass in the atom comes from neutral particles named neutrons

12. Location in the Atom Name(Symbol) Relative Absolute(C)* Relative(amu)† Absolute(g) Proton (p+) 1+ +1.60218x10-19 1.00727 1.67262x10-24 Nucleus Neutron (n0) 0 0 1.00866 1.67493x10-24 Nucleus Outside Nucleus Electron (e-) 1- -1.60218x10-19 0.00054858 9.10939x10-28 Properties of the Three Key Subatomic Particles Charge Mass * The coulomb (C) is the SI unit of charge. † The atomic mass unit (amu) equals 1.66054x10-24 g.

13. Atomic Symbols, Isotopes, Numbers A X The Symbol of the Atom or Isotope Z X = Atomic symbol of the element A = mass number; A = Z + N Z = atomic number (the number of protons in the nucleus) N = number of neutrons in the nucleus Isotope = atoms of an element with the same number of protons, but a different number of neutrons Figure 2.8 Silberberg, Principles of Chemistry

14. Figure 2.10 The modern periodic table. Silberberg, Principles of Chemistry

15. The Modern Reassessment of Dalton’s Atomic Theory 1. All matter is composed of atoms that are indivisible and indestructible. The atom is the smallest body thatretains the unique identityof the element. However, it can, under unusual circumstances, be destroyed (converted to energy) and it can be divided into smaller parts. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in nuclear reactions. 3. All atoms of an element have the same number of protons and electrons, which determines the chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in mass number. A sample of the element is treated as though its atoms have anaverage mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios.

16. To understand the electronic structure of the atom we need to review the properties of electromagnetic radiation.

17. Spectra Site • http://jersey.uoregon.edu/vlab/elements/Elements.html Absorption and emission spectra for element arranged on the periodic table

18. Metals and Color Metal ions: • provide the color in fireworks and flares (and when you burn colored newspapers) • cause glass to have different colors (stained glass) • are responsible for the colors of many precipitates (like the purple Co3(PO4)2)

19. The Wave Nature of Light Frequency and Wavelength • c = l n • l = wavelength • n = frequency • C = speed of light Silberberg, Principles of Chemistry

20. λ, ν, and Energy • As λdecreases and ν increases, what happened to the energy of the radiation? where h = Planck’s constant (6.626 × 10-34m2 kg/s)

21. The infinite number of wavelengths of electromagnetic radiation have been classified into groups as shown below. Regions of the electromagnetic spectrum. Silberberg, Principles of Chemistry

22. Light is a wave…right? • Light falling on alkali metals causes electrons to be released from the metal. • The # of electrons depends on the intensity of light. • There are specific wavelengths of light that cause the release of e-. • This is called the photoelectric effect.

23. Light is a wave…right? • Einstein’s interpretation of the photoelectric effect (1905) was that light is quantized in packets of set energy called photons. (He won the Nobel Prize for this.) • This meant that light had characteristics of particles!

24. Electrons are particles…right? • In 1925, de Broglie stated that all particles have a wavelength described by the equation: λ = h/p where p= momentum • Electrons show diffraction pattern like light when passing through a slit. • So light and particles have a dual nature.

25. Back to atomic structure… • Bohr theorized that the emission spectra of atoms described by Rydberg’s equation were caused by the transition of electrons between specific energy levels (orbits). • http://www.upscale.utoronto.ca/GeneralInterest/Harrison/BohrModel/Flash/BohrModel.html

26. Figure 7.10 The Bohr explanation of the three series of spectral lines. Silberberg, Principles of Chemistry

27. Electron locations • When an electron occupies its usual energy level it is in the ground state. • When an electron absorbs a photon and moves to a higher energy level it is in an excited state. • The energy levels are “quantized”. Atoms can only transition between set levels. • Why are the levels set where they are?

28. More on electrons as waves • Since electrons have wave motion Schrödinger applied the classic wave equations to the motion of a hydrogen electron. Certain wavelengths reinforced each other and were allowed. • This generated regions occupied by an electron of set energy termed orbitals.

29. More on electrons as waves • Heisenberg stated that in measuring the electron there is uncertainty so we can only calculate a probable location for the electron. This is called the Heisenberg Uncertainty Principle.

30. Figure 7.16 Electron probability in the ground-state H atom. Silberberg, Principles of Chemistry

31. Figure 7.19 The 2p orbitals. Silberberg, Principles of Chemistry

32. The 3d orbitals. Figure 7.20 Silberberg, Principles of Chemistry

33. Figure 7.21 One of the seven possible 4f orbitals. Silberberg, Principles of Chemistry

34. Atomic Shape • So…when you get all the electrons in their orbitals surrounding the nucleus, the shape is basically spherical. This is why you often see atoms represented as balls.

35. Ions… • Electrons are arranged on levels or “shells”. Atoms are most stable with 8 electrons on their outermost shell (except for the small atoms). This is often referred to as the octet rule. • Number of electrons on the outer shell = Group number for elements in Groups I-VIII A (using the US convention labeling on periodic table)

36. Figure 2.10 The modern periodic table. US Convention IUPAC Convention Silberberg, Principles of Chemistry

37. Ions… • To achieve the octet, atoms with more than 4 electrons on the outer shell will gain enough electrons to reach 8. • Cl Group 7 7 electrons on outer shell Cl will gain 1 e- . • Now #e- = 18 but #p+ = 17 so the chlorine is an ion with the formula Cl1-.

38. Ions… • O Group 6 6 electrons on outer shell O will gain 2 e- . • Now #e- = 10 but #p+ = 8 so the oxygen is an ion with the formula O2-. • P Group 5 5 electrons on outer shell P will gain 3 e- • Now #e- = 18 but #p+ = 15 so the phosphorus is an ion with the formula P3-

39. Ions… • Na Group 1 1 electrons on outer shell Na will lose 1 e- to expose the complete inside shell. • Now #e- = 10 but #p+ = 11 so the sodium is an ion with the formula Na1+. • Mg Group 2 2 electrons on outer shell Mg will lose 2 e- • Now #e- = 10 but #p+ = 12 so the magnesium is an ion with the formula Mg2+

40. -- 1+ 2+ 3+ -- 3- 2- 1- Generally metals form cations and non-metals form anions. CHM 1010 PGCC Barbara Gage

41. Figure 2.11 From Silberberg, Principles of Chemistry The formation of an ionic compound. Transferring electrons from the atoms of one element to those of another results in an ionic compound.

42. Binary Ionic Compounds • Cations(positively charged ions)and anions (negatively charged ions) will associate with each other and form a neutral binary compound to reduce energy. • 1 Na+ 1 Cl-NaCl • 1 Ca2+ 1 O2-CaO • 3 K+ 1 N3- K3N • 2 Al3+ 3 S2- Al2S3

43. Binary Ionic Compounds • The “B” group elements may form cations with more than one charge. • To specify which cation forms a compound, a roman numeral equal to the charge is added to the cation name. Fe2+ = iron (II) Fe3+ = iron (III)

44. Polyatomic Ions • Some atoms form bonds that hold atoms together in a structure that has an overall charge (rather than as a neutral compound. These ions are called polyatomic ions. • CO32- carbonate SO42- sulfate • Compounds with polyatomic ions are named with the cation and anion name. • Na2SO4 – sodium sulfate

45. Common Polyatomic Ions (Silberberg pg 54 – most common) NH4+ ammonium H3O+ hydronium NO2- nitrite NO3- nitrate SO32- sulfite SO42- sulfate S2O32- thiosulfate HSO4- hydrogen sulfate or bisulfate OH- hydroxide CN- cyanide PO43- phosphate HPO42- hydrogen phosphate H2PO4- dihydrogen phosphate CO32- carbonate HCO3- hydrogen carbonate or bicarbonate ClO- hypochlorite ClO2- chlorite ClO3- chlorate ClO4- perchlorate C2H3O2- acetate (or CH3COO- or CH3CO2-) MnO4- permanganate CrO42- chromate Cr2O72- dichromate O22- peroxide

46. Solid ionic compound Molten ionic compound Ionic compound dissolved in water Electrical conductance and ion mobility. Figure 9.9 Silberberg, Principles of Chemistry

47. Covalent Compounds • Non-metals form bonds by sharing electrons rather than transferring them to achieve the octet. • The resulting bond is referred to as a covalent bond. Each pair of shared electrons = 1 bond • Shared electrons move around the nuclei of both atoms in the bond so both atoms have possession of the shared electrons.

48. Covalent Binary Compounds • CO2 carbon dioxide • N2O dinitrogen oxide • P2S5diphosphoruspentasulfide • SiCl4 silicon tetrachloride • For compounds where two vowels occur together when the prefix is added, the vowel from the prefix can be dropped (except for iodine). • N2O5 dinitrogenpentaoxide

49. Electron Distribution in a Covalent Bond • Are electrons shared equally in a covalent bond? • If not, why not? • Distance of electrons from nucleus and number of protons in the nucleus • Electronegativity – attraction of one atom in a bond for the electrons in that bond

50. Figure 9.20 The Pauling electronegativity (EN) scale. Silberberg, Principles of Chemistry