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) 1803John Dalton: atomic model with several postulates 1897-nowModern Investigations…

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

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

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 that retains the unique identity of 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 an average 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/element s/Elements.html http://jersey.uoregon.edu/vlab/element s/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 Co 3 (PO 4 ) 2 )

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

20. λ, ν, and Energy As λ decreases and ν increases, what happened to the energy of the radiation? where h = Planck’s constant (6.626 × 10 -34 m 2 kg/s)

21. Regions of the electromagnetic spectrum. The infinite number of wavelengths of electromagnetic radiation have been classified into groups as shown below. 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/General Interest/Harrison/BohrModel/Flash/Bo hrModel.html http://www.upscale.utoronto.ca/General Interest/Harrison/BohrModel/Flash/Bo hrModel.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. Figure 7.20 The 3d orbitals. 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. Silberberg, Principles of Chemistry US Convention IUPAC Convention

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 Cl 1-.

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 O 2-. 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 P 3-

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 Na 1+. 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 Mg 2+

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

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

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 Ca 2+ 1 O 2- CaO 3 K + 1 N 3- K 3 N 2 Al 3+ 3 S 2- Al 2 S 3

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. Fe 2+ = iron (II) Fe 3+ = 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. CO 3 2- carbonate SO 4 2- sulfate Compounds with polyatomic ions are named with the cation and anion name. Na 2 SO 4 – sodium sulfate

45. Common Polyatomic Ions (Silberberg pg 54 – most common) NH 4 + ammonium H 3 O + hydronium NO 2 - nitrite NO 3 - nitrate SO 3 2- sulfite SO 4 2- sulfate S 2 O 3 2- thiosulfate HSO 4 - hydrogen sulfate or bisulfate OH - hydroxide CN - cyanide PO 4 3- phosphate HPO 4 2- hydrogen phosphate H 2 PO 4 - dihydrogen phosphate CO 3 2- carbonate HCO 3 - hydrogen carbonate or bicarbonate ClO - hypochlorite ClO 2 - chlorite ClO 3 - chlorate ClO 4 - perchlorate C 2 H 3 O 2 - acetate (or CH 3 COO - or CH 3 CO 2 - ) MnO 4 - permanganate CrO 4 2- chromate Cr 2 O 7 2- dichromate O 2 2- peroxide

46. Figure 9.9 Electrical conductance and ion mobility. Solid ionic compound Molten ionic compound Ionic compound dissolved in water 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 CO 2 carbon dioxide N 2 Odinitrogen oxide P 2 S 5 diphosphorus pentasulfide SiCl 4 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). N 2 O 5 dinitrogen pentaoxide

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

51. Polarity When atoms in a bond have different electronegativities, the electron sharing is unequal. As the ΔEN increases, the electron distribution becomes more uneven and the molecule becomes polar.

52. Polarity HCl EN H = 2.1 EN Cl = 3.0 ΔEN = 0.9 The end with the higher EN will be slightly negative and the other will be slightly positive δ+ H – Cl δ- H – Cl

53. Silberberg, Principles of Chemistry Figure 9.22  EN 3.0 2.0 0.0 Boundary ranges for classifying ionic character of chemical bonds.