1. Atomic Theory

2. Defining the Atom The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) –He believed that atoms were indivisible and indestructible –His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy

3. Dalton’s Atomic Theory (experiment based!) 3)Atoms of different elements combine in simple whole-number ratios to form chemical compounds 4)In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element. 1)All elements are composed of tiny indivisible particles called atoms 2)Atoms of the same element are identical. Atoms of any one element are different from those of any other element. John Dalton (1766 – 1844)

4. Dalton’s Atomic Theory (1808) 1. Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. 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

5. 2

6. 8 X 2 Y 16 X8 Y + 2.1

7. Structure of the Nuclear Atom One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles: –Electrons, protons, and neutrons are examples of these fundamental particles –There are many other types of particles, but we will study these three Dalton Wasn’t Exactly Correct…..

8. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

9. J.J. Thomson, measured mass/charge of e - (1906 Nobel Prize in Physics) 2.2

10. Cathode Ray Tube 2.2

11. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. J. J. Thomson

12. Modern Cathode Ray Tubes  Cathode ray tubes pass electricity (electrons) through a gas that is contained at a very low pressure. Television Computer Monitor

13. Mass of the Electron 1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge The oil drop apparatus Mass of the electron is 9.11 x 10 -28 g

14. e - charge = -1.60 x 10 -19 C Thomson’s charge/mass of e - = -1.76 x 10 8 C/g e - mass = 9.10 x 10 -28 g Measured mass of e - (1923 Nobel Prize in Physics) 2.2

15. (Uranium compound) 2.2

17. 1.atoms positive charge is concentrated in the nucleus 2.proton (p) has opposite (+) charge of electron (-) 3.mass of p is 1840 x mass of e - (1.67 x 10 -24 g)  particle velocity ~ 1.4 x 10 7 m/s (~5% speed of light) (1908 Nobel Prize in Chemistry) 2.2

18. Rutherford’s Findings a) The nucleus is small b) The nucleus is dense c) The nucleus is positively charged  Most of the particles passed right through  A few particles were deflected  VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions:

19. 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 “If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.”

20. The Rutherford Atomic Model Based on his experimental evidence: –The atom is mostly empty space –All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus” –The nucleus is composed of protons and neutrons (they make the nucleus!) –The electrons distributed around the nucleus, and occupy most of the volume –His model was called a “nuclear model”

21. Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 · In 1932 James proved the existence of neutral particles in an atom · James said that the neutrons were just about the same weight as protons · He discovered this by using alpha rays, which are charged, and therefore repelled by considerable electrical forces present in the nuclei of heavier atoms · Chadwick led the way to the starting of penetrating and splitting the nuclei of atoms. · Also led the way to the fission of uranium 235, which eventually created the atomic bomb

22. HOMEWORK 1. Describe JJ Thompson’s CRT (Cathode Ray Tube) experiment & how it showed that atoms contain particles he called “electrons.” 2. Describe JJ’s model of the atom. 3. Explain Rutherford’s scattering experiment and what it helped to prove. Also, how did it disprove Thompson’s model? 4.Describe Rutherford’s atomic model. 5.What led to Chadwick’s discovery.

23. mass p + = mass n o = 1840 x mass e - 2.2 Mass (amu) 0 1 1

24. 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 their nuclei 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

26. 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons Do You Understand Isotopes? 2.3 How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ?

27. Ions: Atoms with different number of protons (p + ) and electrons (e - ) 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

28. 13 protons, 10 (13 – 3) electrons 34 protons, 36 (34 + 2) electrons Do You Understand Ions? 2.5 How many protons and electrons are in ?Al 27 13 3+ How many protons and electrons are in ? Se 78 34 2-

29. Practice: 2.5 1)Write the elemental symbol if 10 p +, 11n o, 10e - 21 Ne 2) Write the elemental symbol if 20p +, 20n o, 18e - 40 Ca 2+

30. Period Group Alkali Metal Noble Gas Halogen Alkaline Earth Metal 2.4

31. Chemistry In Action Natural abundance of elements in Earth’s crust Natural abundance of elements in human body 2.4

32. 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

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. A monatomic ion contains only one atom A polyatomic ion contains more than one atom 2.5 Na +, Cl -, Ca 2+, O 2-, Al 3+, N 3- OH -, CN -, NH 4 +, NO 3 -

35. 2.5

36. 2.6

37. 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

38. ionic compounds consist of a combination of cations and an anions the formula is always the same as the empirical formula the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6

39. 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-

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41. 2.7

42. Polyatomic Ions +1-2-3 Ammonium NH 4 Nitrate NO 3 Carbonate CO 3 Phosphate PO 4 Acetate C 2 H 3 O 2 Sulfate SO 4 Phosphite PO 3 Cyanide CN Sulfite SO 3 Hydroxide OH Dichromate Cr 2 O 7 Permanganate MnO 4 Chlorate ClO 3 Bicarbonate HCO 3 Nitrite NO 2

43. WRITING FORMULAE (DIATOMIC & POLYATOMIC) Polyatomic 1) Potassium Nitrate 2) Sodium Sulfate 3) Potassium Dichromate 4) Ammonium Phosphate 5) Copper I Carbonate 6) Iron (III) Cyanide 7) Silver Sulfite 8) Tin (II) Nitrite 9) Calcium Hydroxide 10) Boron Acetate Mixed 11. Sodium Sulfide 12. Sodium Bicarbonate 13. Potassium Oxide 14. Potassium Permanganate 15. Magnesium Chloride 16. Calcium Chlorate 17. Iron (II) Phosphate 18. Copper Nitride 19. Ammonium Phosphide 20. Aluminum Acetate

44. 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

45. 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

46. Hydrates: – ionic compounds with Water Molecules contained in their crystalline structure. CuCl 2 · 2 H 2 O Copper Chloride Dihydrate Fe(SO 3 ) · 5 H 2 O Iron (II) Sulfate Pentahydrate Cu(NO 3 ) 2 · 4 H 2 O Copper Nitrate Tetrahydrate 2.7

47. Hydrates: – USE PREFIXES TO DENOTE THE # OF WATER MOLECULES 1 = 2 = 3 = 4 = 5 = 6 = 7 = 8 = 9 = 10 = mono di tri tetra penta hexa hepta octa nona deca

48. Molecular compounds nonmetals or nonmetals + metalloids common names H 2 O, NH 3, CH 4, C 60 element further left in periodic table is 1 st element closest to bottom of group is 1 st if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom last element ends in ide 2.7

49. HIhydrogen 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

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51. An acid can be defined as a substance that yields hydrogen ions (H + ) when dissolved in water. HCl Pure substance, hydrogen chloride Dissolved in water (H + Cl - ), hydrochloric acid An oxoacid is an acid that contains hydrogen, oxygen, and another element. HNO 3 nitric acid H 2 CO 3 carbonic acid H 2 SO 4 sulfuric acid 2.7 HNO 3

52. 2.7

55. A base can be defined as a substance that yields hydroxide ions (OH - ) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH) 2 barium hydroxide 2.7