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Chapter 6 The Periodic Table

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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

8 The Families of Elements

9 Classification of the Elements
Metals Nonmetals Metalloids

10 Occupy the left side of the periodic table Have luster, shiny
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 Occupy the right side of the Periodic Table
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 right Nonmetals 4 or more valence electrons tend to gain electrons become negative ions Elements on the left Metals 3 or less valence electrons tend to lose valence electrons form positive 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

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 Lanthanides Elements 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)

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


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

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

35 Atomic Radius 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 increases Why? 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


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


60 Summary of Trends

61 Another Summary

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