2. ATOMIC STRUCTURE THE HALF OF KNOWLEDGE IS KNOWING WHERE TO FIND KNOWLEDGE. UNKNOWN Ch2. 1. J.C. Rowe Windsor University School of Medicine

3. CONCEPT MAP Elements, compounds & mixtures Separation methods Solutions Atoms have structure Atomic nucleus Electronic structure Size of atoms Atoms & ions Symbols for atoms & ions

4. Elements, Mixtures and Compounds  Elements:  Are materials that could not be separated into simpler things either physically or chemically  Elements contain only one type of atom  Iron is an example of an element  Mixture:  Composition of two or more substances that can be separated physically  Example of a chemical mixture is the combination of Sand and Salt  Compound:  Combination of two or more elements  The units of a compound have different properties  Compounds have fixed compositions  Example – combination of Iron and Sulphur.

5. Mixtures What is a mixture? How can it be separated? A mixture is a physical blend of two or more substances. Their composition varies (Air). There are two types of mixtures: homogenous and heterogenous

6. Homogenous or Heterogenous? 1.Air 2.Salt water 3.Tea 4.Brass 5.Vinegar 6.Hydrogen peroxide 7.Steel 1.Salad dressing 2.Apple 3.Sand 4.Paint 5.Granite 6.Laundry detergent 7.Cereal

7. Heterogenous Mixtures  Is the type of mixture that is not uniform in composition.  If one portion of the above mixture were sampled, it’s composition would vary.

8. Homogenous Mixtures  Is a type of mixture that has a completely uniform composition throughout.  It’s components are evenly distributed throughout the sample.

9. Solutions Is the special name that scientists give to homogenous mixtures. Solutions may be of gases, liquids or solids. An example: solution of sugar solid in liquid water.

10. What is a Solution ? Ques: A Mixture of 2 or more soluble substances dissolved in a solvent Can you name any solutions ?

11. What is a solution ? Solutions are a special kind of mixture. Solution = solvent (major) + solute (minor). A solvent may be a solid, a liquid or a gas. When the solvent is a liquid the solution is said to be aqueous (water) or non-aqueous (other than water). Solution can be homogeneous or heterogeneous.

12. Solutions Some common types of solutions System Examples System Examples Gas-gas CO 2 and O in N (air) Liquid-gas Water vapor in air Gas-liquid CO 2 in H 2 O (Soda water) Liquid-liquid Acetic acid in H 2 O (vinegar) Solid-liquid NaCl in H 2 O (brine) Solid-solid Cu in Ag (Sterling silver)

13. Separating Mixtures If you were given a mixture of iron nails, salt and water… How would you separate this mixture completely? Based on which physical properties would you base your method on?

14. Separating Mixtures How would you separate the components in tap water? Distillation A liquid is boiled to produce vapor that is then condensed again to a liquid

15. Separating Mixtures

16. Conditions for solubility  Temperature  Solids usually are more soluble in liquids as the temperature increases;  Gases are less soluble in warm water than cold water;  Gases dissolve more in liquids at higher pressure.  Ionic materials dissolve in water  Covalent materials dissolve in covalent liquids

17. Separating mixtures  Filtration  Evaporation  Distillation/fractional distillation  Gravity separation  chromatography

18. The nuclear atom  Rutherford showed in 1911 that atoms:  Have a very tiny nucleus at the center  Have all their positive charge in the nucleus  Have almost all their mass in the nucleus  Have only negative electrical charges making up the rest of their volume

19. Model of an Atom

21. The Rutherford Experiment

24. Properties of the 3 sub-atomic particles Name of particle symbolRelative mass chargelocation Electrone0Always outside the nucleus Neutronn10Always in the nucleus Protonp1+1Always in the nucleus

25. Protons/Mass number/Isotope  The atomic number of an element is the number of protons in an atom of that element.  The mass number of an element is the sum of the number of protons & the number of neutrons in one atom of that element.  Isotopes of an element are atoms of that element containing the same number of protons but different numbers of neutrons.

26. Standard Atomic Notation

27. Radioactivity  If there are the “wrong” number of neutrons in a nucleus, the atom changes to get the number right. The atom fires out one or more fragments from its nucleus. This type of change is called radioactivity.  The changing radioactive atoms are said to decay.

28. Electron shells  The electrons are not moving at random in the space around the nucleus.  They are arranged in layers, one outside the other. These layers of electrons are called electron shells.  A shell can only hold up to a maximum number of electrons. (2n^2)  1rst shell (2electrons); 2 nd shell (8electrons)  3 rd shell (18electrons); 4 th shell (32 electrons)

30. Atoms & ions  Atoms are electrically neutral.  Removing an electron from an atom leaves the atom with a net positive charge.  Adding an extra electron to an atom gives it a net negative charge.  Charged atoms or groups of atoms are called ions.

31. Atoms vs. ions

32. Symbols for some common atoms & ions Atom or ion symbolProton # # of electrons Net charge Chlorine atomCl17 0 Chloride ionCl - 1718 Iron atomFe26 0 Iron ionFe 2+ 2624+2 Nitrate ionNO 3 - ** Oxide ionO 2- 810-2 Sodium atomNa11 0 Sodium ionNa + 1110+1 Sulphate ionSO 4 2- **-2

33. An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 2.5

34. HISTORY OF THE ATOM 460 BC Democritus develops the idea of atoms he pounded up materials in his pestle and mortar until he had reduced them to smaller and smaller particles which he called ATOMA (greek for indivisible)

35. Dalton’s Atomic Theory (1808) 1. Elements are composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements. 2. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same. 3.Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2.1

36. 8 X 2 Y 16 X8 Y + 2.1

37. 2

38. HISTORY OF THE ATOM 1913 Niels Bohr studied under Rutherford at the Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons.

39. Bohr’s Atom electrons in orbits nucleus

40. HELIUM ATOM + N N + - - proton electron neutron Shell What do these particles consist of?

41. ATOMIC STRUCTURE Particle proton neutron electron Charge + ve charge -ve charge No charge 1 1 nil Mass

42. ATOMIC STRUCTURE He the number of protons and neutrons in an atom4 Atomic mass number the number of protons in an atom2 Atomic number number of electrons = number of protons

43. atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m Rutherford’s Model of the Atom 2.2

44. Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4  + 9 Be 1 n + 12 C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10 -24 g 2.2

45. Subatomic Particles mass p = mass n = 1840 x mass e - 2.2

46. Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in the nucleus X A Z H 1 1 H (D) 2 1 H (T) 3 1 U 235 92 U 238 92 Mass Number Atomic Number Element Symbol 2.3

48. Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal 2.4

49. ATOMIC STRUCTURE Electrons are arranged in Energy Levels or Shells around the nucleus of an atom. first shella maximum of 2 electrons second shella maximum of 8 electrons third shella maximum of 8 electrons

50. ATOMIC STRUCTURE There are two ways to represent the atomic structure of an element or compound; 1.Electronic Configuration 2.Dot & Cross Diagrams

51. ELECTRONIC CONFIGURATION With electronic configuration elements are represented numerically by the number of electrons in their shells and number of shells. For example; N Nitrogen 14 2 in 1 st shell 5 in 2 nd shell configuration = 2, 5 2 + 5 = 7 7

52. ELECTRONIC CONFIGURATION Write the electronic configuration for the following elements; Ca O ClSi Na 20 40 11 23 8 17 16 35 14 28 B 11 5 a)b)c) d)e)f) 2,8,8,2 2,8,1 2,8,72,8,42,3 2,6

53. DOT & CROSS DIAGRAMS With Dot & Cross diagrams elements and compounds are represented by Dots or Crosses to show electrons, and circles to show the shells. For example; Nitrogen N XX X X XX X N 7 14

54. DOT & CROSS DIAGRAMS Draw the Dot & Cross diagrams for the following elements; OCl 817 16 35 a)b) O X X X X X X X X Cl X X X XX X X X X X X X X X X X X X

55. SUMMARY 1. The Atomic Number of an atom = number of protons in the nucleus. 2. The Atomic Mass of an atom = number of Protons + Neutrons in the nucleus. 3. The number of Protons = Number of Electrons. 4. Electrons orbit the nucleus in shells. 5. Each shell can only carry a set number of electrons.

56. 2.6

57. A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance. An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2OH2O H2OH2O molecularempirical C 6 H 12 O 6 CH 2 O O3O3 O N2H4N2H4 NH 2 2.6

58. ionic compounds consist of a cation and an anion e.g. NaCl is really composed of Na + & Cl _ the formula is always the same as the empirical formula the sum of the charges on the cation and anion in each formula unit must equal zero… Is this always true? The ionic compound NaCl 2.6

59. The term molecule is not normally used to apply to ionic substances.. Can you guess why? The?

60. Formula of Ionic Compounds Al 2 O 3 2.6 2 x +3 = +63 x -2 = -6 Al 3+ O 2- CaBr 2 1 x +2 = +22 x -1 = -2 Ca 2+ Br - Na 2 CO 3 1 x +2 = +21 x -2 = -2 Na + CO 3 2-

61. Some Polyatomic Ions 2.7

62. Chemical Nomenclature Ionic Compounds –often a metal + nonmetal –anion (nonmetal), add “ide” to element name BaCl 2 barium chloride K2OK2O potassium oxide Mg(OH) 2 magnesium hydroxide KNO 3 potassium nitrate 2.7

63. Transition metal ionic compounds –indicate charge on metal with Roman numerals FeCl 2 2 Cl - -2 so Fe is +2 iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3 iron(III) chloride Cr 2 S 3 3 S -2 -6 so Cr is +3 (6/2)chromium(III) sulfide 2.7

64. Molecular compounds 1.nonmetals or nonmetals + metalloids 2.common names H 2 O, NH 3, CH 4, C 60 3.element further left in periodic table is 1 st 4.element closest to bottom of group is 1 st 5.if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom 6.last element ends in “ide”… this only applies for members of Group 5, 6, & 7 2.7

65. HI hydrogen iodide NF 3 nitrogen trifluoride SO 2 sulfur dioxide N 2 Cl 4 dinitrogen tetrachloride NO 2 nitrogen dioxide N2ON2Odinitrogen monoxide Molecular Compounds 2.7 TOXIC ! Laughing Gas

66. A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2H2 H2OH2ONH 3 CH 4 A diatomic molecule contains only two atoms H 2, N 2, O 2, Br 2, HCl, CO A polyatomic molecule contains more than two atoms O 3, H 2 O, NH 3, CH 4 2.5

67. Quality is never an accident. It is always the result of high intention, sincere effort and skillful execution. It represents the wise choice of many alternatives. unknown