History German J. W. Dobereiner Grouped elements into triads

Transcription

1 The Periodic Table

2 History 1829 German J. W. Dobereiner Grouped elements into triads One of these triads included chlorine, bromine, and iodine; another consisted of calcium, strontium, and barium. In each of these triads, the atomic weight of the intermediate element is approximately the average of the atomic weights of the other two elements. The density of that element is approximately the average of the densities of the other two elements.

3 History The problem with this arrangement was that Dobereiner s model became outdated as new elements were identified. A good model is able to incorporate newly understood information. Dobereiner s Triad Model was not useful, since several newly discovered elements did not fit into it. Not all elements had triads

4 History Russian scientist Dmitri Mendeleev taught chemistry in terms of properties Mid 1800 atomic masses of elements were known Wrote down the elements in order of increasing atomic mass Found a pattern of repeating properties

5 Mendeleev s Table Grouped elements in columns by similar properties in order of increasing atomic mass Found some inconsistencies - felt that the properties were more important than the mass, so switched order. Found some gaps Must be undiscovered elements Predicted their properties before they were found

6 The Modern Table Elements are still grouped by properties Similar properties are in the same column Order is in increasing atomic number Added a column of elements Mendeleev didn t know about. The noble gases weren t found because they didn t react with anything.

7 Horizontal rows are called periods There are 7 periods

8 Vertical columns are called groups. Elements are placed in columns by similar properties. Also called families

9 1A 2A The elements in the A groups are called the representative elements 3A 4A 5A 6A7A 8A 0

10 Metals

11 Metals Luster shiny. Ductile drawn into wires. Malleable hammered into sheets. Conductors of heat and electricity.

12 Transition metals The Group B elements

13 Dull Brittle Nonconductors - insulators Non-metals

14 Metalloids or Semimetals Properties of both Semiconductors

15 These are called the inner transition elements and they belong here

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17 Group 1A are the alkali metals Group 2A are the alkaline earth metals

18 Group 7A is called the Halogens Group 8A are the noble gases

19 Group Characteristics Alkali metals Group 1: very reactive metals which do not occur freely in nature. 1 electron in outer shell Alkaline Earth Metals Group 2: next reactive metals, found in earths crust but not in elemental form. 2 electrons in outer shell Transition Elements Group 3-12: metals with varying reactivities. Greater density than Group 1 or 2 elements. 1-2 electrons in outer shell Lanthanides and Actinides These elements are also transition elements but have been taken out to prevent the periodic table being so wide.

20 Boron Group Group 13: reactive, contains metal and metalloid. 3 electrons in outer shell Carbon Group Group 14: contains metalloids, metals and non metals. 4 electrons in outer shell Nitrogen Group Group 15: contains metalloids, metals and non metals. 5 electrons in outer shell Oxygen Group Group 16: contains contains metalloids, metals and non metals. Reactive 6 electrons in outer shell Halogens Group 17: non-metals, very reactive. 7 electrons in outer shell Nobel gas Group 18: non-metals, non reactive. 8 electrons in outer shell

21 Periodic Table Part 2 Periodic trends Identifying the patterns

22 What we will look for Periodic trends- How properties vary as you go across a period Group trends How properties vary as you go down a group Why? Explain why they vary

23 What we will investigate Atomic size how big the atoms are Ionization energy How much energy to remove an electron Electronegativity The attraction for the electron in a compound Ionic size How big ions are

24 Ionization Energy The amount of energy required to completely remove an electron from a gaseous atom. Removing one electron makes a +1 ion The energy required is called the first ionization energy

25 What determines IE The greater the nuclear charge the greater IE. Increased atomic radius decreases IE

26 Group trends As you go down a group first IE decreases because of Bigger atoms so outer electron less attracted even though there are more +charges in the nucleus

27 Periodic trends All the atoms in the same period Same approximate size Increasing nuclear charge So IE generally increases from left to right.

28 Discussion Question What is Atomic Radius? The radius of an atom The distance an atom travels The width of an atom The distance around an atom

29 Atomic Size First problem where do you start measuring The electron cloud doesn t have a definite edge. They get around this by measuring more than 1 atom at a time

30 Atomic Size } Radius Atomic Radius = half the distance between two nuclei of molecule

31 Trends in Atomic Size Influenced by two factors Energy Level (Electron Shell) Higher energy level (shell) is further away Charge on nucleus More charge pulls electrons in closer

32 Group trends As we go down a group each atom has another energy level (electron shell) So the atoms get bigger H Li Na K Rb

33 Periodic Trends As you go across a period the radius gets smaller. Same energy level (electron shell) More nuclear charge Pulls outermost electrons closer Na Mg Al Si P S Cl Ar

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35 What is Ionic Radius? the radius of an atom the radius of the most common ion of an atom the size of an atom the width of an ion

36 Ionic Size Cations are positive ions Cations form by losing electrons Metals form cations Cations are smaller than the atom they come from

37 Ionic size Anions are negative ions Anions form by gaining electrons Nonmetals form anions Anions are bigger than the atom they come from

38 Group trends Adding energy level (electron shell) Ions get bigger as you go down H 1+ Li 1+ Na 1+ K 1+ Rb 1+ Cs 1+

39 Periodic Trends Across the period nuclear charge increases so they get smaller. Energy level (electron shell) changes between anions and cations Li 1+ B 3+ N 3- O 2- F 1- Be 2+ C 4+

40 Electronegativity The tendency for an atom to attract electrons to itself when it is chemically combined with another element. How greedy Big electronegativity means it pulls the electron toward it.

41 Group Trend The further down a group Bigger atoms (outer electrons further from nucleus) More electrons per atom Less attraction for electrons Lower electronegativity.

42 Periodic Trend Metals - left end Low nuclear charge Low attraction for extra electrons Low electronegativity Right end - nonmetals High nuclear charge Large attraction for extra electrons High electronegativity Not noble gases- no compounds

43 Ionization energy, electronegativity INCREASE

44 Atomic size increases, Ionic size increases

45 Energy Levels & Shielding Nuclear Charge

46 How to answer why questions Trend Periodic and Group Reason Nuclear charge Energy level and distance from nucleus Result What happens to which electron