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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 propertiesOne element in each triad had properties intermediate of the other two elementsCl, 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
Mendeleevblanks for undiscovered
elementsWhen discovered, his predictions
accurate
Problems w/ orderCo to NiAr to KTe to I
A better arrangement1913, Henry Moseley – British
physicist, arranged elements according to increasing atomic number
Periodic Law When elements arranged in order of
increasing atomic #, periodic repetition of phys & chem props
Horizontal rows = periods7 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 6.2Classifying the ElementsOBJECTIVES:
Describe the information in a periodic table.
Section 6.2Classifying the ElementsOBJECTIVES:
Classify elements based on electron configuration.
Section 6.2Classifying the ElementsOBJECTIVES:
Distinguish representative elements and transition metals.
Groups of elements - family names
Group IA – alkali metalsForms “base” (or alkali) when
reacting w/ H2O (not just dissolved!)
Group 2A – alkaline earth metalsAlso 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:
1) Noble gases2) Representative elements
3) Transition metals
4) Inner transition metals
Let’s now take a closer look at these.
Electron Configurations in Groups
1) Noble gases in Group 8A (also called Group 18)
very stable = don’t react e- configuration w/ outer s & p
sublevels fullfull
Electron Configurations in Groups
2) 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
3) Transition metals in “B” columns outer s sublevel full Start filling “d” sublevel “Transition” btwn metals &
nonmetals
Electron Configurations in Groups
4) 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
1A
2A 3A 4A 5A 6A7A
8A Elements 1A-7A groups called
representative elements
outer s or p filling
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
Periodic table rap
Let’s take a quick break……
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10
5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1
H1
Li3
Na11
K19
Rb37
Cs55
Fr87
Do you notice any similarity in these configurations of the alkali metals?
He2
Ne10
Ar18
Kr36
Xe54
Rn86
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6
Do you notice any similarity in the configurations of the noble gases?
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.
s2s1
He
Transition Metals - d block
d1 d2 d3s1
d5 d5 d6 d7 d8s1
d10 d10
Note the change in configuration.
The P-blockp1 p2 p3 p4 p5 p6
F - block
Called “inner transition elements”
f1 f5f2 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
“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
4d4d5d5d
f orbitals start filling at 4f….2 less than period #
1
2
3
4
5
6
7 4f4f
5f5f
Demo p. 165
Section 6.3Periodic Trends
OBJECTIVES:
Describe trends among the elements for atomic size.
Section 6.3Periodic Trends
OBJECTIVES:
Explain how ions form.
Section 6.3Periodic Trends
OBJECTIVES:
Describe periodic trends for first ionization energy, ionic size, and electronegativity.
Trends in Atomic Size
Measure Atomic Radius - half distance btwn 2 nuclei of diatomic molecule (i.e. O2)
Units of picometers (10-12 m… 1 trillionth)
}Radius
ALLALL Periodic Table Trends Influenced by 3 factors:
1. Energy LevelHigher 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 Going down a
group, each atom has another energy level (floor)
atoms get
bbiiggggeerr
HLi
Na
K
Rb
#1. Atomic Size - Period Trends left to right across period:
size gets smaller
e-’s occupy same energy levelmore nuclear chargeOuter e-’s pulled closer
Na Mg Al Si P S Cl Ar
Here is an animation to explain the trend.
Atomic Number
Ato
mic
Rad
ius
(pm
)
H
Li
Ne
Ar
10
Na
K
Kr
Rb
3
Period 2
Trends of Atomic Radius
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
IonsMetals LOSE electrons, from outer
energy level Sodium loses 1 e-
more p+ (11) than e- (10)+ charge particle formed…“cation”
Na+ called “sodium ion”
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 ThirdHHeLiBeBCNO F Ne
1312 2731 520 900 800 1086 1402 1314 1681 2080
5247 7297 1757 2430 2352 2857 3391 3375 3963
11810 14840 3569 4619 4577 5301 6045 6276
Table 6.1, p. 173
Symbol First Second ThirdHHeLiBeBCNO F Ne
1312 2731 520 900 800 1086 1402 1314 1681 2080
5247 7297 1757 2430 2352 2857 3391 3375 3963
11810 14840 3569 4619 4577 5301 6045 6276
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
energyEasier to achieve (lower IE)
Shielding effect
Shieldinge-’s in outer
energy level “looks through” all other energy levels to see nucleus
Ionization Energy - Group trendsgoing 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 levelSame shielding Increasing nuclear charge So IE generally increases left - right
Exceptions…full & 1/2 full orbitals
Firs
t Ion
izat
ion
ener
gy
Atomic number
He He has greater IE than H.
Both have same shielding (e- in 1st level) He = greater nuclear
charge
H
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li lower IE than H more shielding further away These outweigh
greater nuclear charge
Li
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Be higher IE than Li
same shielding greater nuclear
charge
Li
Be
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He B has lower IE
than Be same shielding greater nuclear
charge By removing an
electron we make s orbital half-filled
Li
Be
B
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne Ne has a lower
IE than He Both are full, Ne has more
shielding Greater
distance
Firs
t Ion
izat
ion
ener
gy
Atomic number
H
He
Li
Be
B
C
N
O
F
Ne Na has a lower
IE than Li Both are s1
Na has more shielding
Greater distance
Na
Firs
t Ion
izat
ion
ener
gy
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 levelgreater 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
Ions get bigger going down, b/c of extra energy level
Li1+
Na1+
K1+
Rb1+
Cs1+
Ion Period Trends Across period
nuclear charge increases Ions get smaller.
energy level changes between anions and cations.
Li1+
Be2+
B3+
C4+
N3-O2- F1-
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)
Al3+
Mg2+
Na1+ Ne F1- O2- N3-
13 12 11 10 9 8 7
#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 TrendFurther down a group,
farther e- is away from nucleus, plus the more e-’s an atom has
more willing to shareLow electronegativity
Electronegativity Period Trend Metals let e-’s go easily
low electronegativity
Nonmetals want more electronstake them away from othersHigh electronegativity.
Trends in Electronegativity
0
Chemistry Song "Elemental Funkiness" - Mark Rosengarten
