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History Structure Trends
The Periodic Table History Structure Trends
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Mendeleev Arranged elements in order of atomic mass with 8 long columns and several short columns Arrangement reflected properties of the elements Predicted existence of several elements to fill gaps in his table These elements were later discovered and had the properties predicted
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Mendeleev’s Periodic Law
Properties of the elements are a periodic function of their atomic masses.
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Moseley Performed experiments to determine an accurate mass for several elements which seemed out of place on the table Noticed a pattern in the number of protons Reorganized elements in order of atomic number rather than mass
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Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers.
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Periods 1 2 7 rows 3 4 5 6 7
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Groups or Families 1 18 2 13 14 15 16 17 3 4 5 6 7 8 9 10 11 12
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This Arrangement Reflects:
Properties Increasing atomic number Electron configuration
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Li – 1s22s1 Na – 1s22s22p63s1 K – 1s22s22p63s23p64s1
Electron configuration and properties by group, ex. Group 1 Li Li – 1s22s1 Na – 1s22s22p63s1 K – 1s22s22p63s23p64s1 Na K Same Group (family) All have 1 e- in the outer level All are similar in color and hardness All react vigorously with water
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Electron configuration by period, ex. Period 2
C - 1s22s22p2 N - 1s22s22p3 O - 1s22s22p4 Same period All have a full 1st energy level with e- in s and p orbitals of the second level Number of e- is increasing by 1
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Blocks Electron configuration for members of block end by filling designated orbital within period s p d f
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General Properties of Metals
1 to 3 e- in outer energy level of most Lose e- to form + ions (cations) Shiny (lustrous) Hard Malleable and ductile Good conductors of heat and electricity Solid at room temp (298K), except Hg
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General Properties of Nonmetals:
5 to 8 electrons in the outer level Gain e- to form – ions (anions) Brittle solids or gases Poor conductors of heat Poor conductors of electricity
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Classification of Elements:
Some elements have properties similar to both metals and nonmetals. These are found bordering the stair-step dividing line. Exception: Al is a metal These elements are called metalloids
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Metals, Metalloids, and Nonmetals
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Some of the Families Have Special Names.
Metals Non-metals 17 – halogens 18 – noble gases 1 – alkali metals 2 – alkaline earth metals 3-12 – transition metals Elements 58 – 71 lanthanides Elements 90 – actinides Rare Earth Metals
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Trends: Atomic Radius Radius of an atom without regard to surrounding atoms (size) Radius depends on: the number of energy levels the strength of the nucleus r
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As n increases, size of the e- cloud increases.
Each period represents a higher principal quantum number (n) As n increases, size of the e- cloud increases. F Atomic radius increases down a family. Cl Br
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Across a period, nuclear charge increases by 1 for each element.
A stronger nucleus acts like a stronger magnet which attracts the e- cloud. C N O Atomic radius decreases across a period.
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General Trend for Atomic Radius:
Increases
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Use position in the periodic table to determine which is larger?
Ag or Au? Ni or Cu? La or U? H or He? Na or Rb? Cl or I? Al or Si? K or Ca?
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Ionic Radius Ions are charged atoms formed when: Atoms lose e-
(+ Ion) = Cation Atoms gain e- (- Ion) = Anion
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Cations are smaller than their respective neutral atoms.
Metals usually lose all valence electrons. So then, the atom has one less energy level. - e- + The nucleus pulls tighter on the remaining electrons.
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Anions are larger than their respective neutral atoms.
Nonmetals usually gain electrons to complete the valence shell. + e- - Electrons repel each other and spread out more. - - more e- than p+
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Which is larger? Ca or Ca+2 F or F–1 K or K+1 O or O-2 F-, Ne or Na+
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First Ionization Energy
Energy needed to remove the most loosely held electron from a gaseous atom. Factors that affect ionization energy: Radius Nuclear charge Shielding effect Stability of sublevels
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Radius: The greater the distance between the nucleus and the valence electrons, the easier it is to lose an electron.
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Nuclear charge: # of p+ Within a period, the higher the nuclear charge, the higher the ionization energy.
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Shielding Effect: Other e- block the pull of the nucleus on the outer e-. Electrons repel each other.
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General Trend for Ionization Energy:
Increases
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Electron Affinity: Attraction of an atom for an additional electron.
Factors affecting EA: Size Nuclear charge Shielding effect Stability of sublevels
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General Trend for Electron Affinity:
Increases
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Electronegativity = ability of atom to become negative ion
Ability to “steal” electrons Increases across a period and up a row
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Valence electrons – see group #
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