Chapter 6 The Periodic Table

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Chapter 6The Periodic Table

• The Elements by Tom Lehrer

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Organizing the Elements

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

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Mendeleevs Periodic Table

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

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Mendeleev

• blanks for undiscovered elements
• When discovered, his predictions accurate
• Problems w/ order
• Co to Ni
• Ar to K
• Te to I

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A better arrangement

• 1913, Henry Moseley British physicist, arranged
  elements according to increasing atomic number

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The Elements by Tom Lehrer
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Periodic Law

• When elements arranged in order of increasing
  atomic , periodic repetition of phys chem
  props
• Horizontal rows periods
• 7 periods
• Vertical column group (or family)
• Similar phys chem prop.
• IDed by letter (IA, IIA)

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Areas of periodic table

• 3 classes of elements
• 1) Metals electrical conductors, have luster,
  ductile, malleable
• 2) Nonmetals generally brittle and
  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

• OBJECTIVES
• Describe the information in a periodic table.

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Section 6.2Classifying the Elements

• OBJECTIVES
• Classify elements based on electron configuration.

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Section 6.2Classifying the Elements

• OBJECTIVES
• Distinguish representative elements and
  transition metals.

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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
• salt-forming

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Electron Configurations in Groups

• Elements sorted based on e- configurations
• 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 w/ outer s p sublevels full

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Electron Configurations in Groups

• Representative Elements Groups 1A - 7A
• wide range of properties
• Representative of all elements
• s p sublevels of highest energy level NOT
  filled
• Group equals of e- in highest energy level

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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 main body of PT, in
  2 horizontal rows
• outer s sublevel full
• Start filling f sublevel
• Once called rare-earth elements
• not true b/c some abundant

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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?
2
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
• He has properties of noble gases
• has a full outer level of e-s
• 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 row (or period) is 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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Section 6.3Periodic Trends

• OBJECTIVES
• Describe trends among the elements for atomic
  size.

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Section 6.3Periodic Trends

• OBJECTIVES
• Explain how ions form.

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Section 6.3Periodic Trends

• OBJECTIVES
• Describe periodic trends for first ionization
  energy, ionic size, and electronegativity.

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Trends in Atomic Size

Radius

• Measure Atomic Radius - half distance btwn 2
  nuclei of diatomic molecule (i.e. O2)
• Units of picometers (10-12 m 1 trillionth)

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ALL Periodic Table Trends

• Influenced by 3 factors
• 1. Energy Level
• Higher energy levels further away from nucleus.
• 2. Charge on nucleus ( protons)
• More charge pulls electrons in closer. ( and
  attract each other)
• 3. Shielding effect

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What do they influence?

• Energy levels Shielding have effect on GROUP (
  ? )
• Nuclear charge has effect on PERIOD ( ? )

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1. Atomic Size - Group trends
H

• Going down a group, each atom has another energy
  level (floor)
• atoms get bigger

Li
Na
K
Rb
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1. Atomic Size - Period Trends

• left to right across period
• size gets smaller
• e-s occupy same energy level
• more nuclear charge
• Outer e-s pulled closer

Here is an animation to explain the trend.
Na
Mg
Al
Si
P
S
Cl
Ar
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K
Period 2
Na
Li
Atomic Radius (pm)
Kr
Ar
Ne
H
Atomic Number
10
3
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Trends of Atomic Radius
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Ions

• Some compounds composed of ions
• ion is atom (or group of atoms) w/ or - charge
• Atoms are neutral because the number of protons
  electrons
• - ions formed when e- transferred (lost or
  gained) btwn atoms

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Ions

• Metals LOSE electrons, from outer energy level
• Sodium loses 1 e-
• more p (11) than e- (10)
• charge particle formedcation
• Na called sodium ion

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Ions

• Nonmetals GAIN one or more electrons
• Cl gains 1 e-
• p (17) e- (18), so charge of -1
• Cl1- called chloride ion
• anions

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2. Trends in Ionization Energy

• 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 ionization energy is E required to remove
  2nd e-
• Always greater than first IE.
• third greater than 1st or 2nd IE.

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Table 6.1, p. 173
Symbol First Second Third
5247 7297 1757 2430 2352 2857 3391 3375 3963
1312 2731 520 900 800 1086 1402 1314 1681 2080
HHeLiBeBCNO F Ne
11810 14840 3569 4619 4577
5301 6045 6276
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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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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

• e-s in outer energy level looks through all
  other energy levels to see nucleus

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Ionization Energy - Group trends

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

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Ionization Energy - Period trends

• Atoms in same period
• same energy level
• Same shielding
• Increasing nuclear charge
• So IE generally increases left - right
• Exceptionsfull 1/2 full orbitals

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He

• He has greater IE than H.
• Both have 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
• same shielding
• greater nuclear charge

H
First Ionization energy
Be
Li
Atomic number
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He

• B has lower IE than Be
• same shielding
• greater nuclear charge
• By removing an electron we make s orbital
  half-filled

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, because removing an
  electron leaves it with a 1/2 filled p orbital

N
O
C
H
First Ionization energy
Be
B
Li
Atomic number
63
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 are full,
• Ne has more shielding
• 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 has more shielding
• Greater distance

F
N
O
C
H
First Ionization energy
Be
B
Li
Na
Atomic number
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First Ionization energy
Atomic number
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Trends in Ionization Energy (IE)
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Driving Forces

• Full Energy Levels require high E to remove e-
• Noble Gases full orbitals
• Atoms want noble gas configuration

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2nd Ionization Energy

• For elements w/ filled or ½ filled orbital by
  removing 2 e-, 2nd IE lower than expected.
• True for s2
• Alkaline earth metals form 2 ions.

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3rd IE

• Using the same logic s2p1 atoms have an low 3rd
  IE.
• Atoms in the aluminum family form 3 ions.
• 2nd IE and 3rd IE are always higher than 1st IE!!!

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Trends in Ionic Size Cations

• Cations form by losing electrons.
• metals
• Cations are smaller than the atom they came from
  
• they lose electrons
• they lose an entire energy level.
• Cations of representative elements have noble gas
  configuration before them.

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Trends in Ionic size Anions

• Anions gain electrons
• Anions bigger than the atom they came from
• same energy level
• greater area the nuclear charge needs to cover
• Nonmetals

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Configuration of Ions

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

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Configuration of Ions

• Non-metals form ions by gaining electrons to
  achieve noble gas configuration.
• They end up with the configuration of the noble
  gas after them.

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Ion Group trends
Li1
Na1

• Each step down a group is adding an energy level
• Ions get bigger going down, b/c of extra energy
  level

K1
Rb1
Cs1
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Ion Period Trends

• Across period
• nuclear charge increases
• Ions get smaller.
• energy level changes between anions and cations.

N3-
O2-
F1-
B3
Li1
Be2
C4
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Size of Isoelectronic ions

• Iso- means the same
• Isoelectronic ions have the same of electrons
• Al3 Mg2 Na1 Ne F1- O2- and N3-
• all have 10 electrons
• all have the same configuration 1s22s22p6
  (which is the noble gas neon)

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Size of Isoelectronic ions?

• Positive ions that have more protons would be
  smaller (more protons would pull the same of
  electrons in closer)

N3-
O2-
F1-
Ne
Na1
Al3
7
10
9
8
11
13
12
Mg2
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3. Trends in Electronegativity

• Electronegativity is tendency for atom to
  attract e-s when atom in a compound
• Sharing e-, but how equally do they share it?
• Element with big electronegativity means it pulls
  e- towards itself strongly!

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Electronegativity Group Trend

• Further down a group, farther e- is away from
  nucleus, plus the more e-s an atom has
• more willing to share
• Low electronegativity

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Electronegativity Period Trend

• Metals let e-s go easily
• low electronegativity
• Nonmetals want more electrons
• take them away from others
• High electronegativity.

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