The Periodic Table

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The Periodic Table
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Dimitri Mendeleev

• Russian scientist developed the first published
  table in 1869.
• Arranged elements in order of atomic mass
• Elements with similar properties were placed in
  columns
• left spaces for undiscovered elements

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Periodic Law the pattern of the table

• Mendeleevs table was called a periodic table
  because..
• Original Periodic Law states
• The physical and chemical properties of elements
  were periodic functions of their atomic masses.
• In other words . . .

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In other words...

• If the elements were arranged in order of
  increasing atomic mass then elements with similar
  properties would show up PERIODICALLY.
• These similar elements were organized into groups
  or families

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Henri Moseley (1914)

• Used X-rays to reveal atomic numbers of several
  elements.
• Suggested that the elements should be arranged in
  order of atomic number instead of atomic mass.

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Modern Periodic Law(Moseley)

• The chemical and physical properties of the
  elements are the periodic functions of their
  atomic number.

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In other words . . .

• when elements are arranged according to
  increasing number of protons (Atomic Number) the
  same chemical properties show up repeatedly

of p of e- in an ATOM
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Reading the Periodic Table

• Families or Groups
• Numbered 1 to 18
• Elements in groups (families) have similar
  characteristics

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• Periods

.
When you write a sentence the period goes at the
end
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Classifications of Elements

• Metals
• Non-metals
• Metalloids (semi-metals)

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Metal characteristics

• Solids at STP (except mercury)
• good conductors of electricity and heat
• Mobile electrons
• shiny luster
• ductile - drawn (made) into wire
• malleable - hammered into sheets
• when combined in compound have a positive
  oxidation state (, cation)

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• Elements to the left of the steps are metals
• except H, Ge Sb

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• Metallic character increases R to L
• The elements with the strongest metallic
  characteristics are found in Group 1
• Shiniest, conduct electricity the best . . .

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Non-metal characteristics

• Solid, liquid or gas at STP
• used as insulators
• solids are dull
• brittle
• when combined in a compound have negative
  oxidation state (-, anion)

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Non-metals
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• Non-metallic character increases L to R
• The elements with the strongest non-metal
  characteristics are found in Group 18

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Inert non-metals

• Filled valence shells

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Metalloids

• a.k.a. Semimetals
• both metal and non-metal properties

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Metalloids

• sit on the steps
• 2 hide underneath

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Elements are found as (s), (l), (g)

• Solid
• particles have vibratory motion and are tightly
  packed
• Almost ALL of the elements
• Liquid
• particles can move throughout substance
• particles are farther apart than in solids
• conforms to shape of container
• ONLY Hg and Br

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• Gas
• molecules in constant, random, straight line
  motion
• molecules fill container
• large space between volumeless molecules
• conforms to the shape of the container
• H, N, O, F, Cl and Group 18

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Allotropes

• Different forms of an element found naturally in
  the same state
• Different molecular structures
• Carbon
• diamond, coal, graphite
• Oxygen
• O2 O3

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Diatomic elements

• Elements found bonded to itself
• uncombined
• N2 O2 F2 Cl2 Br2 I2 and H2

H
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Monoatomic elements

• Lonely drifters
• no friends

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Electron configuration

• Periods correspond to the orbital which starts to
  be filled is the valence shell.
• Blocks refer to the sublevels being filled.
• Sublevels s, p, d, f

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e-config
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Valence electrons

• The number of valence electrons correspond to
  group number
• 1, 2 , 13 - 18

• The shell theyre in is the same as the period

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Group / Family Characteristics

• Each group has characteristic properties that are
  directly related to electron configuration
  especially the number of valence electrons

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Group 1

• Alkali Metals
• all elements except hydrogen
• most reactive metal group
• SO reactive they are never found alone in nature
• always bonded to another element
• form 1 ions

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Group 2

• Alkaline Earth Metals
• second most reactive metals group
• SO reactive they dont occur alone in nature
• always bonded to another element
• form 2 ions

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d- block elements

• Transition Metals
• COLORful

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• Transition metals not as reactive as other metals
• some found in free state
• Au, Ag, Pt
• multiple oxidation states
• Iron (IV) oxide FeO2
• Iron (II) oxide FeO

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Group 17

• Halogens
• Most reactive nonmetal group
• SO reactive that they dont occur by themselves
  in nature
• at least bonded to themselves (diatomic)
• form -1 ions

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Group 18

• Noble gases
• Inert gases
• dont like to combine with other elements
• Valence shells filled

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Ionization Energy
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Ionization Energy (first)

• Definition
• energy needed to remove the most loosely held e-

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Whats holding the e- in place?

• Opposites attract
• There is a force of attraction between protons
  () and electrons (-)
• The more protons in the nucleus the stronger the
  positive charge is and the e- are held tighter

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Ionization energy

• Hint energy needed to make an ion by losing
  electrons

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MOST elements want 8 e- (octet rule)

• Elements with only a few valence electrons will
  tend to have lower ionization E
• In other words
• It doesnt take a lot of E to remove their e-

Metals
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MOST elements want 8 e- (octet rule)

• Elements with almost 8 valence electrons wont
  give them up so easy
• It takes a lot of E to remove their e-
• Non-metals have high Ionization E

Who has the highest IEs?
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Graph Ionization Energies

• Label X axis with Atomic 1-20
• Label y axis with 1st IE (KJ/mol atoms)
• AFTER you plot the points
• Label data points with element symbols
• Connect data points of same period only
• Separate your graph into Periods

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He
Period 2
Period 3
Period 1
Period 4
Ne
Ar
First Ionization Energy (KJ / mol)
H
Ca
Li
Na
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• What is the periodic trend for ionization energy?

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Ionization E trend?

• As move across a periodIE increases. WHY?

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• As you go across a period
• a large proton and a tiny electron are being
  added.
• more p hold the e- tighter

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• As move UP a group IE increases. WHY?
• Period 1
• Period 2
• Period 3
• Period 4
• Lets figure it out . . .

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Potassium 19 p

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Sodium 11 p

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Lithium 3 p

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because . . .

• As you go up a group
• the valence shell gets closer to the nucleus and
  can hold on to the e- tighter.

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SHIELDING

• Reason why IE decreases as you go down a group
• the distance increases between the p and valence
  e-
• AND
• e- in outer shells repel each other

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Electronegativity

• Ability of an atom to attract electrons of other
  atoms.
• In other words. . .
• Atoms with high e-neg are bullies that steal
  electrons

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Electronegativity Graph
Fluorine
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Electronegativity

• Determine the periodic trend

EXCEPT for Noble gases. WHY?
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• Atoms with
• high electronegativities
• also have
• high ionization E

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Atomic Radius

• Radius
• distance from the center of a circle (Nucleus)
  to the outermost edge (valence shell)

R
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Atomic Radius Periodicity
DECREASES
DECREASES
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Why?????

• Why does atomic radius DECREASE as you move up a
  group?

• Losing layers of e-

• Why does atomic radius DECREASE as you move
  across a period?

• Increasing the of p holds the e- in tighter

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• What happens to atomic radius as you create a
  ion?

• radius decreases

• What happens to atomic radius as you create a -
  ion?

• radius increases

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• Compare the ionic radus of Mg2 and the atomic
  radius of Ne.

The radius of the Mg2 ion is smaller than the
atom of Ne, because the Mg2 ion has more p (12)
than the Ne atom (10).
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Compare . . .

• Fluoride ion Fluorine atom
• Sodium ion Sodium atom
• Fluoride ion Neon atom

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General Formulas of compounds

• You can look at the groups on the periodic table
  and determine how they will combine with elements
  of different groups.

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Write these formulas

• lithium fluoride
• sodium fluoride
• potassium fluoride
• lithium chloride
• sodium chloride
• potassium chloride

• What do you notice?
• Write the general formula for Group 1 and Group
  17 elements

AB
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Write these formulas

• beryllium fluoride
• magnesium fluoride
• calcium fluoride
• beryllium chloride
• magnesium chloride
• calcium chloride

• What do you notice?
• Write the general formula for Group 2 and Group
  17 elements

AB2
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Write these formulas

• lithium oxide
• sodium oxide
• potassium oxide
• lithium sulfide
• sodium sulfide
• potassium sulfide

• What do you notice?
• Write the general formula for Group 1 and Group
  16 elements

A2B
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Write these formulas

• beryllium oxide
• magnesium oxide
• calcium oxide
• beryllium sulfide
• magnesium sulfide
• calcium sulfide

• What do you notice?
• Write the general formula for Group 2 and Group
  16 elements

AB
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Formula Writing Naming Review

• Lead IV oxide
• Phosphorus pentoxide
• SO3
• Oxygen
• Argon
• Aluminum oxide
• NiO

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Ionization Energy
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Ionization Energy (first)

• Definition
• energy needed to remove the most loosely held e-

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Second Ionization energy

• Energy needed to remove a second electron.
• Which element has a higher second IE?
• K or Ca
• Why?

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Ionization energy

• Hint energy needed to make an ion by losing
  electrons

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ALL elements want 8 e- (octet rule)

• Elements with only a few valence electrons will
  tend to have lower ionization E
• In other words
• It doesnt take a lot of E to remove their e-

Metals
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ALL elements want 8 e- (octet rule)

• Elements with almost 8 valence electrons wont
  give them up so easy
• It takes a lot of E to remove their e-
• Non-metals have high Ionization E

Who has the highest IEs?
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Graph Ionization Energies

• Label X axis with Atomic
• Label y axis with 1st IE (KJ/mol atoms)
• AFTER you plot the points
• Label data points with element symbols
• Connect data points of same period only
• Separate your graph into Periods

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He
Period 2
Period 3
Period 1
Period 4
Ne
Ar
First Ionization Energy (KJ / mol)
H
Ca
Li
Na
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• What is the periodic trend for ionization energy?

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IE trend?
Harder to remove

• As move across a periodIE increases. WHY?

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Whats holding the e- in place?

• Across a period
• The more protons in the nucleus the stronger the
  nuclear charge is and the e- are held tighter
• Takes more energy to remove them

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• As move UP a group IE increases. WHY?
• Period 1
• Period 2
• Period 3
• Period 4
• Lets figure it out . . .

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Potassium 19 p

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Sodium 11 p

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Lithium 3 p

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because . . .

• As you go up a group
• the valence shell is closer to the nucleus and
  can hold on to the e- tighter.

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SHIELDING

• Reason why IE decreases as you go down a group
• The valence electrons are farther away from the
  nucleus
• AND
• the kernel electrons are pushing it away
• MORE REACTIVE

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Electronegativity

• Ability of an atom to attract electrons of other
  atoms.
• In other words. . .
• Atoms with high e-neg are bullies that steal
  electrons
• MORE REACTIVE

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Electronegativity Graph
Fluorine
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Electronegativity

• Determine the periodic trend

EXCEPT for Noble gases. WHY?
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• Atoms with
• high electronegativities
• also have
• high ionization E

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Atomic Radius

• Radius
• distance from the center of a circle (Nucleus)
  to the outermost edge (valence shell)

R
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Atomic Radius Periodicity
DECREASES
DECREASES
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Why?????

• Why does atomic radius DECREASE as you move up a
  group?

• Losing layers of e-

• Why does atomic radius DECREASE as you move
  across a period?

• Increasing the of p holds the e- in tighter

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• What happens to atomic radius as you create a
  ion?

• radius decreases

• What happens to atomic radius as you create a -
  ion?

• radius increases

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• Compare the ionic radius of Mg2 and the atomic
  radius of Ne.

The radius of the Mg2 ion is smaller than the
atom of Ne, because the Mg2 ion has more p (12)
than the Ne atom (10).
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Compare . . .

• Fluoride ion Fluorine atom
• Sodium ion Sodium atom
• Fluoride ion Neon atom

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General Formulas of compounds

• You can look at the groups on the periodic table
  and determine how they will combine with elements
  of different groups.

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Write these formulas

• lithium fluoride
• sodium fluoride
• potassium fluoride
• lithium chloride
• sodium chloride
• potassium chloride

• What do you notice?
• Write the general formula for Group 1 and Group
  17 elements

AB
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Write these formulas

• beryllium fluoride
• magnesium fluoride
• calcium fluoride
• beryllium chloride
• magnesium chloride
• calcium chloride

• What do you notice?
• Write the general formula for Group 2 and Group
  17 elements

AB2
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Write these formulas

• lithium oxide
• sodium oxide
• potassium oxide
• lithium sulfide
• sodium sulfide
• potassium sulfide

• What do you notice?
• Write the general formula for Group 1 and Group
  16 elements

A2B
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Write these formulas

• beryllium oxide
• magnesium oxide
• calcium oxide
• beryllium sulfide
• magnesium sulfide
• calcium sulfide

• What do you notice?
• Write the general formula for Group 2 and Group
  16 elements

AB
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Formula Writing Naming Review

• Lead IV oxide
• Phosphorus pentoxide
• SO3
• Oxygen
• Argon
• Aluminum oxide
• NiO

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Different Forms of the Periodic Table

• Changes were made to Mendeleevs table to look
  like the modern Periodic table we use today.
• This is not the only form of the periodic table
  that exists, however it is the most widely
  accepted.

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Stowes Physicists p.t.
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Benfey p.t.
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Zmaczynski p.t.
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Alexander Arrangement p.t.
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??????????? p.t.
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THE END