# Chapter 6 “The Periodic Table”

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Chapter 6 “The Periodic Table”
The Elements by Tom Lehrer

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

Mendeleev’s Periodic Table
mid-1800s, about 70 elements known Dmitri Mendeleev – Russian chemist & teacher Arranged elements by increasing atomic mass

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

A better arrangement 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number

The Elements by Tom Lehrer

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. ID’ed by # & letter (IA, IIA)

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)

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

Section 6.2 Classifying the Elements
OBJECTIVES: Describe the information in a periodic table.

Section 6.2 Classifying the Elements
OBJECTIVES: Classify elements based on electron configuration.

Section 6.2 Classifying the Elements
OBJECTIVES: Distinguish representative elements and transition metals.

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; don’t dissolve well, hence “earth metals” Group 7A – halogens “salt-forming”

Electron Configurations in Groups
Elements sorted based on e- configurations: Noble gases Representative elements Transition metals Inner transition metals Let’s now take a closer look at these.

Electron Configurations in Groups
Noble gases in Group 8A (also called Group 18) very stable = don’t react e- configuration w/ outer s & p sublevels full

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

Electron Configurations in Groups
Transition metals in “B” columns outer s sublevel full Start filling “d” sublevel “Transition” btwn metals & nonmetals

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

Elements 1A-7A groups called representative elements
outer s or p filling 1A 8A 2A 3A 4A 5A 6A 7A

The group B called transition elements
These are called the inner transition elements, and they belong here

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

Group 8A are noble gases Group 7A called halogens

Let’s take a quick break……
Periodic table rap

1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Do you notice any similarity in these configurations of the alkali metals?

1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 He Do you notice any similarity in the configurations of the noble gases? 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

Elements in the s - blocks
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.

Transition Metals - d block
Note the change in configuration. s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10

The P-block p1 p2 p3 p4 p6 p5

F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
Called “inner transition elements” f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13

Each row (or period) is energy level for s & p orbitals.
1 2 3 4 5 6 7 Period Number Each row (or period) is energy level for s & p orbitals.

3d 4d 5d “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

f orbitals start filling at 4f….2 less than period #
1 2 3 4 5 6 7 4f 5f f orbitals start filling at 4f….2 less than period #

Demo p. 165

Section 6.3 Periodic Trends
OBJECTIVES: Describe trends among the elements for atomic size.

Section 6.3 Periodic Trends
OBJECTIVES: Explain how ions form.

Section 6.3 Periodic Trends
OBJECTIVES: Describe periodic trends for first ionization energy, ionic size, and electronegativity.

} 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)

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

What do they influence? Energy levels & Shielding have effect on GROUP (  ) Nuclear charge has effect on PERIOD (  )

#1. Atomic Size - Group trends
H Going down a group, each atom has another energy level (floor) atoms get bigger Li Na K Rb

#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

Rb K Period 2 Na Li Atomic Radius (pm) Kr Ar Ne H 3 10 Atomic Number

ion is atom (or group of atoms) w/ + or - charge
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

Metals LOSE electrons, from outer energy level
Ions Metals LOSE electrons, from outer energy level Sodium loses 1 e- more p+ (11) than e- (10) + charge particle formed…“cation” Na+ called “sodium ion”

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

#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.

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

Symbol First Second Third
Table 6.1, p. 173 Symbol First Second Third HHeLiBeBCNO F Ne

Symbol First Second Third
HHeLiBeBCNO F Ne Why did these values increase so much?

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

Shielding e-’s in outer energy level “looks through” all other energy levels to see nucleus

Ionization Energy - Group trends
going down group first IE decreases b/c... e- further away from nucleus attraction more shielding

Ionization Energy - Period trends
Atoms in same period: same energy level Same shielding Increasing nuclear charge So IE generally increases left - right Exceptions…full & 1/2 full orbitals

Both have same shielding (e- in 1st level)
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

These outweigh greater nuclear charge
Li lower IE than H more shielding further away These outweigh greater nuclear charge H First Ionization energy Li Atomic number

greater nuclear charge
He Be higher IE than Li same shielding greater nuclear charge First Ionization energy H Be Li Atomic number

greater nuclear charge
He B has lower IE than Be same shielding greater nuclear charge By removing an electron we make s orbital half-filled First Ionization energy H Be B Li Atomic number

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

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

He Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital N First Ionization energy H C O Be B Li Atomic number

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

Ne has a lower IE than He Both are full, Ne has more shielding
Greater distance F N First Ionization energy H C O Be B Li Atomic number

Na has a lower IE than Li Both are s1 Na has more shielding
He Ne Na has a lower IE than Li Both are s1 Na has more shielding Greater distance F N First Ionization energy H C O Be B Li Na Atomic number

First Ionization energy
Atomic number

Trends in Ionization Energy (IE)

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

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.

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!!!

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.

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

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

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.

Ion Group trends Each step down a group is adding an energy level
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+

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+

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)

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+ 10 9 8 7 13 12 11 Mg2+

#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!

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

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

Trends in Electronegativity

Chemistry Song "Elemental Funkiness" - Mark Rosengarten