Elements and Compounds

Transcription

1 Elements and Compounds Composition of Matter Elements 1

2 Element Identification 2

3 Source of Light Elements Solar Spectra 3

4 Detecting Elements Molecular Spectrum of CO 4

5 Chemical and Physical Properties Elements An Element is the basic form of matter it cannot be broken down into smaller parts by chemical methods Each element has a unique identity Compounds A Compound is a unique substance that is formed from two or more elements and has different properties from all other substances 5

6 Pure Substances Pure substances have unchanging, definite compositions and properties. Table Salt: crystalline white substance that is commonly used to modify food taste NaCl : 1 sodium atom + 1 chlorine atom. Two different elements, unique properties. Properties of Substances A physical property is one that can be observed without performing a change to the substance: Color Density Hardness etc. 6

7 Names of the Elements Latin names from historical context produce unusual symbols for some elements: first letter is capitalized, second is lower case Ferrum: Fe (Iron) Aurum: Au (Gold) Cobalt: Co 7

8 Atomic Theory John Dalton used Greek ideas with experimentation to reveal the atomic nature of the elements The atom is defined as the smallest particle of an element that retains the properties of the bulk element Daltons atomic theory allows the prediction of composition Daltons Atomic Theory Matter is composed of atoms All atoms of each element are identical and possess the same properties Chemical compounds are composed of atoms of different elements in whole number ratio Reactions produce re-arranged atoms 8

9 Law of Constant Composition Any compound is made up of elements in the same proportion by mass. A compound is identified by a specific formula Compounds and Molecules Two or more elements bonded to each other are molecules or molecular compounds Elements bonded together are polyatomic i.e. Cl 2 Molecules may be formed by both types: compounds or polyatomic elements 9

10 Chemical Properties Chemical properties are related to the identity of the element or the compound The ability of a substance to undergo a chemical change: reaction Chemical change is a transformation of a substance into a new substance Mixtures of elements and compounds 10

11 Mixtures Heterogeneous: a non-uniform mixture containing two or more phases Sample of mixture produces different amounts of each substance in container Homogeneous Mixtures One physical state with uniform properties May be a Solution or an Alloy Solution: homogeneous mixture with one liquid phase Alloy: homogeneous mixture of solids with one phase (solid solution) 11

12 Electrons 1897 J.J. Thompson proves existence of the electron Is common to all atoms Has a charge of -1 Used Plum Pudding analogy : electrons and a positive medium mixed as a solid particle Rutherford Experiment Used positively charged helium ions to bombard thin gold foil Showed a dense positively charged center- not a conglomeration of positive and negative charges Calculated charge to mass ratio =-1.76 x 10 8 coulombs/gram From an electron beam 12

13 Electron Properties R.A. Millikan (Univ. Of Chicago) determined charge magnitude of electrons which led to the determination of the mass q= -1.6 x Coulombs Mass = 9.11 x grams (from charge to mass ratio) 13

14 The Nuclear Atom Counting The Particles All atoms of each element are not quite the same Mass of atoms of an element is variable, a result of a new particle, the neutron The nuclear mass is the sum of protons and neutrons: isotopes 14

15 Calculating Isotopic Values Atomic Number = element identity = Protons in nucleus neutrons=mass number protons 90 Sr has 38 protons: = 52 neutrons Still has 38 electrons! atomic number Every atom with an atomic number of 1 is a hydrogen atom. 1 proton in the nucleus 15

16 Every atom with an atomic number of 6 is a carbon atom. 6 protons in the nucleus atomic number atomic number Every atom with an atomic number of 92 is a uranium atom. 92 protons in the nucleus 16

17 Isotopes of the Elements Atoms of the same element can have different masses. They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. The difference in the number of neutrons accounts for the difference in mass. These are isotopes element of the same 17

18 Isotopes of the Same Element Have Equal numbers of protons Different numbers of neutrons Isotopic Notation 18

19 Isotopic Notation 6 protons + 6 neutrons 12 6C 6 protons Isotopic Notation 6 protons + 8 neutrons 14 6C 6 protons 19

20 Isotopic Notation 8 protons + 8 neutrons 16 O 8 8 protons Isotopic Notation 8 protons + 9 neutrons 17 O 8 8 protons 20

21 Isotopic Notation 8 protons + 10 neutrons 18 O 8 8 protons Hydrogen has three isotopes 1 proton 0 neutrons 1 proton 1 neutron 1 proton 2 neutrons 21

22 Examples of Isotopes Element Protons Electrons Neutrons Symbol Hydrogen H Hydrogen H Hydrogen H Uranium U Uranium U Atomic Mass 22

23 The mass of a single atom is too small to measure on a balance. Using a mass spectrometer, the mass of the hydrogen atom was determined. A Modern Mass Spectrometer 5.8 Positive ions formed from sample. Electrical field at slits accelerates positive ions. From the intensity and positions A mass of the lines Deflection the mass of spectrogram spectrogram, positive the different ions is recorded. isotopes and occurs their relative amounts can magnetic be determined. field. 23

24 5.9 A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given. Using a mass spectrometer, the mass of one hydrogen atom was determined to be x g 24

25 This number is very small. small small small small small small small small small small small small small small small small small small The mass of a hydrogen atom is very small. Numbers To overcome of this this size problem are too small a system for of practical relative atomic use. masses using atomic mass units was devised to express the masses of elements using simple numbers x g 25

26 The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon A mass of exactly 12 atomic mass units (amu) was assigned to

27 1 1 amu is defined as exactly equal to the mass of a carbon-12 atom 12 1 amu = x g 12 6 Average atomic mass amu. 27

28 Average atomic mass amu. Average atomic mass amu. 28

29 Average Relative Atomic Mass Most elements occur as mixtures of isotopes. Isotopes of the same element have different masses. The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly amu 29

30 To calculate the atomic mass multiply the atomic mass of each isotope by its percent abundance and add the results. Isotope Isotopic mass (amu) Abundance (%) Average atomic mass (amu) 63 Cu Cu ( amu) = amu ( amu) = amu amu Relationship Between Mass Number and Atomic Number 30

31 The mass number minus the atomic number equals the number of neutrons in the nucleus. mass number atomic number Ag mass number - atomic number = = number of neutrons 62 31

32 Atomic Mass Units Isotopic mass is determined by comparison with a standard: AMU = 1/12 the mass of a 12 C atom Boron isotopes: 10 B is times the mass of 12 C: Therefore has AMU 11 B has AMU Table Values for Atomic Mass Since elements exist in nature with isotopic masses, the reported value is the weighted average of all of these isotopes based on their percent abundance 32

33 Calculating atomic masses Nature contains 19.9% 10 B and 80.1% 11 B. What is the weighted average? 10 B: x AMU = 1.99 AMU 11 B: x AMU = 8.82 AMU AMU Binary Compounds Metal-Nonmetal Single Oxidation Multiple oxidation Polyatomic ions Al 2 O 3 CaCl 2 Fe 2 O 3 ; FeCl 2 Mg 3 (PO 4 ) 2 33

34 Practice: Nomenclature NaCl Sodium Chloride HCl (blue Stuff) Hydrogen Chloride H 2 O (blue Stuff) Dihydrogen oxide FeS Iron(II) oxide Al 2 O 3 Aluminum oxide SiO 2 (blue stuff) Silicon dioxide Binary Compounds NaCl Sodium chloride CaBr 2 Calcium bromide NH 4 Cl Ammonium chloride N 2 O Dinitrogen oxide FeS Iron(II) sulfide Mn 2 O 3 Manganese(III) oxide Sn(NO 2 ) 4 Tin(IV) nitrite Fe(OH) 3 Iron(III) hydroxide 34

35 Non-metal Binary Compounds Use Greek Prefixes (table 4-3) Write first element which is closest to metal Hydrogen is written second First element written as english name Second element(s) end with ide Never use mono- for first element; only remaining element Oxides of Nitrogen N 2 O Dinitrogen monoxide NO Nitrogen monoxide N 2 O 3 Dinitrogen trioxide N 2 O 5 Dinitrogen pentoxide 35

36 Binary Acids Hydrogen plus halogen (group VII) or other anions Name of compound ends with ic acid HCL hydrogen choride Hydrochloric acid HBr- hydrogen bromide Hydrobromic acid Oxy Acids Formed by the combination of hydrogen and most polyatomic ions Name after the root of the anion plus the following endings: -ic acid if the polyatomic ion ended with ate -ous if the anion ended with ite Mineral acids use common names 36

37 Oxyacids C 2 H 3 O - 2 acetate ion CO 2-3 carbonate ion NO - 3 nitrate ion PO 3-4 phosphate ion ClO - hypochlorite ion ClO 2- chlorite ion ClO 3- chlorate ion ClO - 4 perchlorate ion HC 2 H 3 O 2 acetic acid H 2 CO 3 carbonic acid HNO 3 nitric acid H 3 PO 4 phosphoric a. HClO hypochlorous a. HClO 2 chlorous acid HClO 3 chloric acid HClO 4 perchloric acid 37