Chapter 6 The Periodic Table p. 154

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Chapter 6The Periodic Tablep. 154

• The Elements by Tom Lehrer

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Organizing the Elements

• Chemists used elements properties to sort into
  groups.
• 1829 - J. W. Dobereiner
• triads groups of 3 w/ similar properties
• One element in triad
• had properties intermediate
• of other 2 elements
• Cl, Br, and I look different.
• similar chemically

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Mendeleevs Periodic Table

• 1800s, about 70 elements known
• 1869 - Dmitri Mendeleev Russian chemist
  teacher
• Arranged elements by atomic mass

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Mendeleevs Periodic Table

• Blank spaces
• undiscovered elements
• Predicted properties
• predictions very accurate
• Problems w/ order
• Te to I atomic mass decreases
• I belongs w/ Br Cl
• Mendeleev broke rule put Te before I

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A better arrangement

• 1913, Henry Moseley
• British physicist
• Determined atomic s
• Modern PT arranged by atomic

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The Elements by Tom Lehrer
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Periodic Law

• Elements arranged by increasing atomic ,
  periodic repetition of properties present
• Horizontal rows periods
• 7 periods
• Vertical column group (family)
• Similar properties
• IUPAC labels (1-18)
• U.S. system ( letteri.e. IA, IIA)

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Areas of periodic table

• 3 classes of elements
• 1) Metals electrical conductors, lustrous,
  ductile, malleable

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• 2) Nonmetals generally brittle/non-lustrous,
  poor conductors of heat and electricity
• Some gases (O, N, Cl)
• some brittle solids (B, S)
• fuming red liquid (Br)

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• 3) Metalloids border the line-2 sides
• Properties are intermediate between metals and
  nonmetals

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Section 6.2Classifying the Elements p. 161
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Groups of elements - family names

• Group IA alkali metals
• Forms base (or alkali) when reacting w/ H2O
  (not just dissolved!)
• Group 2A alkaline earth metals
• Also form bases with H2O dont dissolve well,
  hence earth metals
• Group 7A halogens
• Greek hals (salt) genesis (to be born)

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Electron Configurations in Groups

• sorted based on e- configs
• Noble gases
• Representative elements
• Transition metals
• Inner transition metals

Lets now take a closer look at these.
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Electron Configurations in Groups

• Noble gases in Group 8A (also called Group 18)
• very stable dont react
• e- configuration
• full outer s p sublevels

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Electron Configurations in Groups

• Representative Elements Groups 1A - 7A
• Properties vary
• Represent all elements
• s p sublevels of highest PEL NOT filled
• Group valence e-s

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Electron Configurations in Groups

• Transition metals in B columns
• outer s sublevel full
• Start filling d sublevel
• Transition btwn metals nonmetals

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Electron Configurations in Groups

• Inner Transition Metals below PT, 2 horizontal
  rows
• outer s sublevel full
• Start filling f sublevel

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• Elements 1A-7A groups called representative
  elements
• outer s or p filling

8A
1A
2A
3A
4A
5A
6A
7A
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• The group B called transition elements

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• Group 1A called alkali metals (but NOT H)
• Group 2A called alkaline earth metals

H
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• Group 8A are noble gases
• Group 7A called halogens

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Periodic table rap

• Lets take a quick break

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• 1s1
• 1s22s1
• 1s22s22p63s1
• 1s22s22p63s23p64s1
• 1s22s22p63s23p64s23d104p65s1
• 1s22s22p63s23p64s23d104p65s24d10 5p66s1
• 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67
  s1

Do you notice any similarity in these
configurations of the alkali metals?
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He

• 1s2
• 1s22s22p6
• 1s22s22p63s23p6
• 1s22s22p63s23p64s23d104p6
• 1s22s22p63s23p64s23d104p65s24d105p6
• 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6

Do you notice any similarity in the
configurations of the noble gases?
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Ne
10
Ar
18
Kr
36
Xe
54
Rn
86
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Elements in the s - blocks
s1
s2
He

• Alkali metals end in s1
• Alkaline earth metals end in s2
• should include He, but
• properties of noble gases
• full outer EL
• group 8A

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Transition Metals - d block
Note the change in configuration.
s1 d5
s1 d10
d1
d2
d3
d5
d6
d7
d8
d10
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The P-block
p1
p2
p6
p3
p4
p5
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F - block

• Called inner transition elements

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1 2 3 4 5 6 7
Period Number

• Each period energy level for s p orbitals.

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• d orbitals fill up in levels 1 less than period
  
• first d is 3d found in period 4.

1 2 3 4 5 6 7
3d
4d
5d
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1 2 3 4 5 6 7
4f 5f

• f orbitals start filling at 4f.2 less than
  period

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Demo p. 165
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Trends in Atomic Size
Section 6.3 Periodic Trends p. 170

• Atomic Radius - half distance btwn 2 nuclei of
  identical atoms
• Increases top-bottom
• Decreases L-R
• picometers
• 10-12 m 1 trillionth

Radius
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ALL PT Trends

• Influenced by 3 factors
• 1. Energy Level
• Higher energy levels further from nucleus
• 2. Charge on nucleus ( protons)
• More charge pulls e-s in closer
• 3. Shielding effect

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1. Atomic Size - Group trends

• Going down a group, atoms gain another PEL
  (floor)
• atoms get..

b
H
i
Li
g
Na
g
K
e
r
Rb
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1. Atomic Size - Period Trends

• L to R across period
• More p in nucleus
• More e-s occupy same energy level
• stronger nuclear charge
• Pulls e- cloud closer to nucleus
• atoms get.

Here is an animation to explain the trend.
Si
Ar
Al
P
S
Cl
Mg
Na
m
a
S
l
l
e
r
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Trends of Atomic Radius
increases
increases
decreases
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Ions p. 172

• Some compounds composed of ions
• Ion - atom (or group of atoms) w/ or - charge
• formed when e- transferred btwn atoms
• Cation (loses e-s ion)
• Anion (gains e-s - ion)

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Cation Formation
Effective nuclear charge on remaining e-s
increases.
Na atom 1 valence e-
11p
Remaining e- pulled closer to nucleus. Ionic
size decreases.
Valence e- lost in ion formation
Result a smaller sodium ion, Na
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Anion Formation
A chloride ion is produced. It is larger than
the original atom.
Chlorine atom with 7 valence e-
17p
One e- is added to the outer shell (from Na for
example).
Effective nuclear charge is reduced and the e-
cloud expands.
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2. Trends in Ionization Energy p.173

• Ionization energy - energy required to completely
  remove e- (from gaseous atom)
• energy required to remove only 1st e-called first
  ionization energy.

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Ionization Energy

• second IE is E required to remove 2nd e-
• Always greater than first IE
• third IE greater than 1st or 2nd IE
• IE helps predict what ions elements form
• Li 1
• Mg 2
• Al 3

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Table 6.1, p. 173
Symbol First Second Third
11810 14840 3569 4619 4577
5301 6045 6276
5247 7297 1757 2430 2352 2857 3391 3375 3963
1312 2731 520 900 800 1086 1402 1314 1681 2080
HHeLiBeBCNO F Ne
Why did these values increase so much?
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Cation Formation
11p
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Anion Formation
17p
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What factors determine IE?

• greater nuclear charge greater IE
• Greater distance from nucleus decreases IE
• Filled half-filled orbitals have lower energy
• Easier to achieve (lower IE)
• Shielding effect

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Shielding Effect

• e-s in outer PEL look thru other PELs to
  see nucleus
• Stays same thru blocks
• Greater influence on IE than nuclear charge

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Shielding Trends
increases
remains constant
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Ionization Energy - Group trends p. 174

• going down group
• first IE decreases b/c...
• e- further from p attraction
• more shielding

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Ionization Energy - Period trends p. 174

• Same period atoms have same
• energy levels
• shielding (within a block slight decrease
  btwn s and p)
• Increasing nuclear charge
• IE generally increases left - right
• Exceptionsfull 1/2 full orbitals

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He

• He greater IE than H.
• Both w/ same shielding (e- in 1st level)
• He - greater nuclear charge

H
First Ionization energy
Atomic number
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He

• Li lower IE than H
• more shielding
• further away
• These outweigh greater nuclear charge

H
First Ionization energy
Li
Atomic number
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He

• Be higher IE than Li
• shielding (period)
• greater nuclear charge

H
First Ionization energy
Be
Li
Atomic number
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• B has lower IE than Be
• greater nuclear charge
• shielding has greater influence on IE
• Slight decrease (p e-)
• p e- removed
• s orbital ½ filled

He
H
First Ionization energy
Be
B
Li
Atomic number
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He
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
N
C
H
First Ionization energy
Be
B
Li
Atomic number
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He

• Oxygen breaks the pattern, b/c removing e- leaves
  it w/ 1/2 filled p orbital

N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
F
N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
Ne

• Ne has a lower IE than He
• Both full but
• Ne more shielding
• b/c greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
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He
Ne

• Na has a lower IE than Li
• Both are s1
• Na - more shielding
• Greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
Na
Atomic number
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Trends in Ionization Energy (IE)
decreases
decreases
increases
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Trends in Ionic Size Cations

• Cations lose e-s
• metals
• Cations smaller than atom they came from
• lose e-s
• lose entire energy level.
• Cations of representative elements have noble gas
  config before them

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Trends in Ionic size Anions

• Anions gain e-s
• nonmetals
• Anions bigger than atom they came from
• same energy level
• greater area nuclear charge needs to cover

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Configuration of Ions

• Ions always have noble gas configurations (full
  outer level)
• Na atom is 1s22s22p63s1
• Forms 1 Na ion 1s22s22p6
• Same as Ne

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Configuration of Ions

• Non-metals form ions by gaining e-s to achieve
  noble gas configuration
• configuration of noble gas after them

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Ion Group trends
Li1
Na1

• Each step down a group adds energy level
• Ions - bigger going down
• more energy levels

K1
Rb1
Cs1
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Ion Period Trends

• Across period
• nuclear charge increases
• Ions get smaller
• energy level changes btwn anions cations

N3-
O2-
F1-
B3
Li1
Be2
C4
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3. Trends in Electronegativity

• Electronegativity (EN)- tendency for atom to
  attract e-s when atom in cmpd
• Sharing e-, but how equally?
• Element w/ big EN pulls e- towards itself
  strongly!

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Electronegativity Group Trend

• Further down group, farther e- away from nucleus
• plus more e-s atom has
• more willing to share
• Low EN

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Electronegativity Period Trend

• Metals let e-s go easily
• low EN
• Nonmetals want more e-s
• take e-s from others
• High EN

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Trends in Electronegativity
decreases
decreases
increases
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• Chemistry Song "Elemental Funkiness" - Mark
  Rosengarten

The Elements Tom Lehrer