1 Chapter 5 Atoms, Molecules, and Ions

2 History n Greeks n Democritus and Leucippus - atomos n Aristotle- 4 elements. n Alchemy-experimentation n Robert Boyle- experimental definition of element. n Lavoisier- Father of modern chemistry. n He wrote the book.

3 Conservation of Mass n In any chemical reaction the total mass of the reactants will equal the total mass of the products g Reactants g Products n 2H 2 + O 2 2H 2 O 4.04g g = 36.04g 4.04g g = 36.04gLaws

Law of Definite Proportion- n The law stating that a pure substance, e.g. H 2 O, will always have the same percent by weight, e.g. 11.2% H and 88.8% O. n elements will combine in specific ratios by mass. 4

Law of Definite Proportions n H 2 0 H 2 x 1.01g = 2.02g O 1 x 16.00g = 16.00g O 1 x 16.00g = 16.00g 18.02g 18.02g 2.02g H / g H 2 O = 0.11 x 100 = 11% 16.00g H / g H 2 O = 0.88 x 100 = 88 % What are the proportions in a CO 2 molecule? 5

6 Law of Multiple Proportions In chemistry, the law of multiple proportions is sometimes called Dalton's Law after its discoverer, the English chemist John Dalton. The law is: If two elements form more than one compound between them, then the ratios of the masses of the second element which combine with a fixed mass of the first element will be ratios of small whole numbers.

7 For example carbon oxide: CO and CO grams of carbon may react with 133 grams of oxygen to produce carbon monoxide (CO) 100 grams of carbon may react with 266 grams of oxygen to produce carbon dioxide (CO 2 ) The ratio of the masses of oxygen that can react with 100 grams of carbon is 266:133 ≈ 2:1, a ratio of small whole numbers.

8 What?! n Compare Water H 2 O and Hydrogen Peroxide H 2 O 2 n H 2 O - H 2 x 1.01g = 2.02 g O 1 x 16.00g = g O 1 x 16.00g = g g Oxygen / 2.02 g Hydrogen = g Oxygen / 2.02 g Hydrogen = 8gO/1gH 8gO/1gH n Water has 8 g of oxygen per 1 g of hydrogen.

Law of Multiple Proportions n H 2 O 2 – H 2 x 1.01g = 2.02g O 2 x 16.00g = 32.00g O 2 x 16.00g = 32.00g g Oxygen / 2.02g Hydrogen = 16gO/1gH 16gO/1gH n Hydrogen peroxide has 16 g of oxygen per g of hydrogen. n Compare the 2 samples n 16g H2O2 / 8g H2 = 2 / 1 ratio n Small whole number ratios. 9

10 Dalton’s Atomic Theory 1) Elements are made up of atoms 2) Atoms of each element are identical. Atoms of different elements are different. 3) Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. 4) Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

11 Experiments to determine what an atom was n J. J. Thomson- used Cathode ray tubes

12 Thomson’s Experiment Voltage source +-

13 Thomson’s Experiment Voltage source +-

14 n Passing an electric current makes a beam appear to move from the negative to the positive end. Thomson’s Experiment Voltage source +-

15 Voltage source Thomson’s Experiment n By adding an electric field

16 Voltage source Thomson’s Experiment n By adding an electric field, he found that the moving pieces were negative + -

17 Thomsom’s Model n Found the electron. n Couldn’t find positive (for a while). n Said the atom was like plum pudding. n A bunch of positive stuff, with the electrons able to be removed.

18 Rutherford’s Experiment n Used uranium to produce alpha particles. n Aimed alpha particles at gold foil by drilling hole in lead block. n Since the mass is evenly distributed in gold atoms alpha particles should go straight through. n Used gold foil because it could be made atoms thin.

19 Lead block Uranium Gold Foil Florescent Screen

20 What he expected

21 Because

22 Because, he thought the mass was evenly distributed in the atom.

23 What he got

24 How he explained it + n Atom is mostly empty n Small dense, positive piece at center. n Alpha particles are deflected by it if they get close enough.

25 +

26 Modern View n The atom is mostly empty space. n Two regions n Nucleus- protons and neutrons. n Electron cloud- region where you might find an electron.

27 Sub-atomic Particles n Z - atomic number = number of protons determines type of atom. n A - mass number = number of protons + neutrons. n Number of protons = number of electrons if neutral.

28 Symbols X A Z Na Mass # Atomic# (p + and n 0 ) (p + )

29 Periodic Table Click Link

30 Metals n Conductors n Lose electrons n Malleable and ductile

31 Nonmetals n Brittle n Gain electrons n Covalent bonds

32 Semi-metals or Metalloids

33 Alkali Metals

34 Alkaline Earth Metals

35 Halogens

36 Transition metals

37 Noble Gases

38 Inner Transition Metals

40 Ions n Atoms or groups of atoms with a charge. n Cations- positive ions - metals losing electrons(s). n Anions- negative ions - nometals gaining electron(s). n Ionic bonding- held together by the opposite charges. n Ionic solids are called salts.

41 Polyatomic Ions n Groups of atoms that have a charge. n No, you don’t have to memorize them. n List on page 97 or on the back of your periodic table. periodic table.

42 Naming compounds n Two types n Ionic - metal and non metal or polyatomics. n If the metal is in the 1 st or 2 nd column (representative element). NO (roman numeral) in the name n If the metal is a transition metal a (roman numeral) may be needed (check your periodic table) n Covalent- we will just learn the rules for 2 non- metals.

43 Ionic compounds n If the cation is monoatomic- Name the metal (cation) just write the name. Mg 2+ Magnesium n If the cation is polyatomic- name it. NH 4 + Ammonium n If the anion is monoatomic- name it but change the ending to –ide. Cl 1- Chloride n If the anion is polyatomic- just name it PO 4 3- Phosphate n Practice.

44 Ionic Compounds n Have to know what ions they form n off table, polyatomic, or figure it out n Two types –Binary ionic – 2 elements (cation + anion) –Ternary ionic – 3 elements –(cation + Polyatomic ion) n CaSK 2 S n AlPO 4 K 2 SO 4 n FeSCoI 3

45 Ionic Compounds n Fe 2 (C 2 O 4 ) n MgO n MnO n KMnO 4 n NH 4 NO 3 n Hg 2 Cl 2 n Cr 2 O 3

46 Ionic Compounds n KClO 4 n NaClO 3 n YBrO 2 n Cr(ClO) 6

47 Writing Formulas n Two sets of rules, ionic and covalent n To decide which to use, decide what the first word or element is. n If it is a metal or polyatomic use ionic. n If it is a non-metal use covalent.

48 Ionic Formulas n Charges must add up to zero. n Get charges from table, name of metal ion, or memorized from the list. n Use parenthesis to indicate multiple polyatomics.

49 Ionic Formulas n Sodium nitride n sodium- Na is always +1 n nitride - ide tells you it comes from the table n nitride is N -3

50 Ionic Formulas n Sodium nitride n sodium- Na is always +1 n Nitride - ide tells you it comes from the table n nitride is N -3 n Doesn’t add up to zero. Na +1 N -3

51 Ionic Formulas n Sodium nitride n sodium- Na is always +1 n nitride - ide tells you it comes from the table n nitride is N -3 n Doesn’t add up to zero n Need 3 Na Na +1 N -3 Na 3 N

52 Ionic Compounds n Sodium sulfite n calcium iodide n Lead (II) oxide n Lead (IV) oxide n Mercury (I) sulfide n Barium chromate n Aluminum hydrogen sulfate n Cerium (IV) nitrite

53 Chemical Bonds n The forces that hold atoms together. n Covalent bonding - sharing electrons. n Makes molecules. n Chemical formula- the number and type of atoms in a molecule. n C 2 H carbon atoms, 6 hydrogen atoms, n Structural formula shows the connections, but not necessarily the shape.

54 H H HH H HCC n There are also other models that attempt to show three dimensional shape. n Ball and stick.

55 Binary Molecular Compounds n Two words, with prefixes. n Prefixes tell you how many atoms. n 1-mono, 2-di, 3-tri, 4-tetr(a), 5-pent(a), 6-hex(a), 7-hept(a), 8 oct(a),9-non(a), 10-dec(a) Name : P 4 Cl 7 n First element whole name with the appropriate prefix, except monoP 4 = tetraphosphorus n Second element, -ide ending with appropriate prefix Cl 7 =heptachloride n Practice

56 n CO 2 n CO n CCl 4 nN2O4nN2O4nN2O4nN2O4 n XeF 6 nN4O4nN4O4nN4O4nN4O4 n P 2 O 10 Naming Covalent Compounds

57 Binary Molecular (Covalent) Compounds n The name tells you how to write the formula n Sulfur dioxide n diflourine monoxide n nitrogen trichloride n diphosphorus pentoxide

58 More Names and formulas

59 Acids n Substances that produce hydrogen (H + ) ions when dissolved in water. n All acids begin with H. n Two types of acids: n Oxyacids – oxygen in the formula n Non-oxyacids – no oxygen in the formula

60 Naming Acids Hydrogen _______ide becomes hydro____ic acid Hydrogen_______ate becomes _________ic acid Hydrogen_______ite becomes _______ous acid Hydrogen always #1 on periodic table

61 Using naming acids table Name : HCl from the periodic table ionic name is hydrogen chloride so use: Hydrogen _______ide becomes hydro______ic acidchlor Transfer word root from one side to the other

62 Using naming acids table Write the formula for sulphuric acid: use Hydrogen_______ate becomes _________ic acidsulphursulph Ionic name is hydrogen sulphate H+H+ SO 4 2- So acid formula is H 2 SO 4

63 Naming acids n If the formula has oxygen in it(oxyacid) n write the name of the anion, but change –ate to -ic acid –ite to -ous acid n Watch out for sulfuric and sulfurous n H 2 CrO 4 n HMnO 4 n HNO 2

64 Naming acids n If the acid doesn’t have oxygen(nonoxyacid) n add the prefix hydro- n change the suffix -ide to -ic acid n HCl nH2SnH2SnH2SnH2S n HCN

65 Formulas for acids n Backwards from names. n If it has hydro- in the name it has no oxygen n Anion ends in -ide n No hydro, anion ends in -ate or -ite n Write anion and add enough H to balance the charges.

66 Formulas for acids n hydrofluoric acid n chromic acid n carbonic acid n hydrophosphoric acid n fluorous acid n perchloric acid n phosphorous acid

67 Hydrates n Some salts trap water crystals when they form crystals. n These are hydrates. n Both the name and the formula needs to indicate how many water molecules are trapped. n In the name we add the word hydrate with a prefix that tells us how many water molecules.

68 Hydrates n In the formula you put a dot and then write the number of molecules. Calcium chloride dihydrate = CaCl 2  2  Calcium chloride dihydrate = CaCl 2  2  Chromium (III) nitrate hexahydrate = Cr(NO 3 ) 3  6H 2 O Chromium (III) nitrate hexahydrate = Cr(NO 3 ) 3  6H 2 O Simulation of a Hydrate Hydrate Nomenclature Practice

69 compounds binary ternary molecular ionicacids molecules hydrates formula unit All nonmetals Metal or hydrogen and nonmetals or Polyatomics trapped water