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Chapter 6 The Periodic Table Development/History of the Modern Periodic Table Using the Periodic Table An Introduction to the Elements Periodic Trends.

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Presentation on theme: "Chapter 6 The Periodic Table Development/History of the Modern Periodic Table Using the Periodic Table An Introduction to the Elements Periodic Trends."— Presentation transcript:

1 Chapter 6 The Periodic Table Development/History of the Modern Periodic Table Using the Periodic Table An Introduction to the Elements Periodic Trends

2 Periodic Table  Why Periodic????  The properties of the elements repeat in in a periodic way.  Invaluable tool for chemistry  Used for organization

3 History of the Periodic Table  Timeline  Trace the development of the Periodic Table by making a timeline  Aristotle  Newlands  Dobereiner  Meyer  Mendeleev  Moseley

4 The Basics  Elements are arranged by atomic number  Typical box contains:  Name of the element  Symbol  Atomic number  Atomic mass

5 Periods Horizontal Rows Numbered 1-7

6 Groups: Vertical Columns Numbered 1-18

7 Interactive Periodic Tables  http://periodic.lanl.gov/index.shtml http://periodic.lanl.gov/index.shtml  www.webelements.com www.webelements.com  www.chemicool.com www.chemicool.com  http://education.jlab.org/itselemental/ele 016.html http://education.jlab.org/itselemental/ele 016.html

8 The Families of Elements http://www.privatehand.com/flash/elements.html

9 Classification of the Elements Metals Nonmetals Metalloids

10 Metals Occupy the left side of the periodic table Have luster, shiny Solids at room temp except Hg Ductile: ability to be drawn into wires Malleable: ability to be hammered into sheets Excellent conductors of heat and electricity Tend to form positive ions

11 NonMetals  Occupy the right side of the Periodic Table  Generally gases or brittle solids  Dull-looking  Brittle  Poor conductors of heat and electricity  Bromine is the only liquid at room temp  Tend to form negative ions

12 Metalloids Characteristics of metals and nonmetals

13 Classification of the Elements  Families of elements share the same ending electron configuration  therefore they share similar chemical characteristics  Valence Electrons: electrons in the highest principal energy level  Determine Chemical reactivity  Elements in a group share the same number of valence electrons

14 The s, p, d and f blocks

15 Number of Valence Electrons Elements on the left Metals 3 or less valence electrons tend to lose valence electrons form positive ions Elements on the right Nonmetals 4 or more valence electrons tend to gain electrons become negative ions

16 Most Common Ions

17 Families of elements  Elements of the same family (group) share structural and chemical (behavioral) characteristics  Alkali Metals  Alkaline Earth Metals  Transition Elements  Halogens  Nobel Gases

18 Group 1: Alkali Metals  Soft, highly reactive metals  Usually stored under oil or kerosene to prevent their interaction with air and water

19 Properties of Alkali Metals  React vigorously with water  Oxidize readily in air  Good conductors of electricity

20 Alkali Metals  Have one valence electron  Will lose this electron very easily  when electron is lost the metal gains a stable non- reactive noble gas configuration Comparison of the Reactivity of the Alkali Metals http://www.youtube.com/watch?v=uixxJtJPVXk http://www.youtube.com/watch?v=uixxJtJPVXk

21 Group 2: Alkaline Earth Metals  Harder, denser, stronger, and have higher melting points than alkali metals  All are reactive not as reactive as group 1

22 Alkaline Earth Metals  Must lose two electrons to gain a stable configuration

23 Groups 3-12:Transition Metals  Not as reactive as Groups 1 and 2  Huge variety but all shiny  Multi valent…form multiple ions  d-block elements  Also include: Inner Transition Elements (Rare Earth Elements)  Elements 58-71 Lanthanides  Elements 90-103 Actinides

24 Group 17: Halogens  Most reactive non-metals  Combine easily with metals; especially the alkali metals

25 Halogens  7 valence electrons, one short of a stable octet.  Will gain one electron to become stable  -1 ions Reaction of chlorine (a halogen) with sodium (an alkali metal) https://www.youtube.com/watch?v=1xT4OFS03jE

26 Element Dating

27 Hydrogen  Most common element in the universe  Chemical family by itself because it behaves so differently  Reacts with most other elements  Rarely found in a free state in nature  1 valence electron

28 The Hindenberg  Filled with H  Very reactive with oxygen gas  He used in blimps today  much less reactive than H

29 Group 18: Noble Gases  Very low reactivity  Filled valence shells: s and p levels in the highest principal energy levels are full  Very stable electron configuration  Many uses: signs, weather balloons and the airships (Blimps)

30 The Octet Rule  Atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons.  Elements on the left (metals) tend to lose valence electrons and form positive ions  Elements on the right (nonmetals) tend to gain electrons to become negative ions

31

32 Periodic Trends  Properties of Elements tend to occur in a predictable way  Known as a trend, as you move across a period or down a group  Knowing element trends allows us to make predictions about an element’s behavior

33 Periodic Properties  Properties  Atomic Radius  Ionic Radius  Electronegativity  Ionization Energy  Questions we will answer:  Definition  How does the property vary across the table?  Why?  How does it vary down a group?  Why?

34 For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms. Atomic Radius

35  The atomic radius is a measure of the size of an atom.  The larger the radius, the larger is the atom.

36 Trends in Atomic Radius  There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus.  Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.

37 Trends in Atomic Radius  The atomic radius decreases as you move across a period  Why?  Increased nuclear charge pulls the electrons in tighter  Added electrons are in the same principal energy levels

38 Group Trends in Atomic Radius  Atomic Radius increases as you move down a group  Why? The increasing number of electrons are in higher energy levels and instead of pulling the electrons closer to the nucleus we see the …

39 Atomic Radius

40 Atomic radius generally increases as you move down a group. The outermost orbital size increases down a group, making the atom larger.

41 Shielding Effect  More inner electrons shield the outer electron from the nucleus and reduce their attraction to the nucleus therefore the overall atomic radius is larger

42 Ionic Radius  Atoms can gain or lose electrons to form ions  Ion: an atom with a charge  Recall that atoms are neutral in charge,  If an electron is lost, then the overall charge is positive  If an electron is gained the atom becomes negative

43 Positive Ion (Cation) Formation  When atoms lose electrons  Radius always becomes smaller Because…  The loss of a valence electron can leave an empty outer orbital resulting in a small radius.  Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.

44 Negative Ion (Anion) Formation  When atoms gain electrons  Radius always increasesWhy?  More electrons mean more electrostatic repulsion resulting in increased diameter.

45 Period Trend for Ionic Radius  As you move left to right across a period  the ionic radius gets smaller for the positive ions  The ionic radius for the negative ions also decreases

46 Group Trend for Ionic Radius  Both positive and negative ions increase in size moving down a group.

47 Ionic Radius

48 Ionization Energy  the amount of energy need to remove an electron from a specific atom or ion in its ground state in the gas phase  High Ionization Energy: atom is holding onto electrons very strongly  Low Ionization Energy: atom is holding electrons less tightly

49  For any element (A) the process of removing an electron can be represented as follows:  A + energy -----> A + + e-  What is the periodic trend in ionization energy? Why?

50 Trends for Ionization Energy  Generally increases as you move across a period  because increased nuclear charge causes an increased hold on the electrons  Ionization Energy decreases as you move down a group  due to increasing atomic size

51

52 Successive Ionization Energies  There is an ionization energy for each electron that is removed from an atom  After the valence electrons are removed Ionization Energies Jump Dramatically  First Ionization Energy: removes 1 electron  Second Ionization Energy: removes a second electron  Third Ionization Energy: removes a third electron

53 Comparing Successive Ionization Energies

54 Trends in Ionization Energy

55 Electronegativity  The ability of an an atom to attract electrons to itself when it is combined with another atom  Expressed in terms of a relative scale: fluorine is assigned a value of 4 and all other elements are calculated relative to this.  The units of electronegativity are arbitrary units called Paulings.  Noble gases have no values because of few chemical compounds

56 Electronegativity  Greater the electronegativity  the higher an atom’s ability to pull an electron to itself when it is bonded to another atom  What are the periodic trends in electronegativity?  Why?

57 Trends in Electronegativity  Electronegativity Increases as you move across a period  Electronegativity decreases you move down a group Where are the elements with highest electronegativity? Where are the elements with lowest electronegativity?

58 Electronegativity

59

60 Summary of Trends

61 Another Summary


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