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I II III
Ch. 6 - The Periodic Table & Periodic Law
I. Development of the Modern Periodic Table(p. 174 - 181)
A. Mendeleev
Dmitri Mendeleev (1869, Russian) Organized elements
by increasing atomic mass
Elements with similar properties were grouped together
There were some discrepancies
A. Mendeleev
Deduced elements existed, but were undiscovered elements, their properties could be predicted
B. Moseley
Henry Moseley (1913, British)
Organized elements by increasing atomic number
Resolved discrepancies in Mendeleev’s arrangement
This is the way the periodic table is arranged today!
C. Modern Periodic Table
1
2
3
4 5
6
7
Group (Family)Period
1. Groups/Families
Vertical columns of periodic tableEach group contains elements with similar
chemical & physical properties (same amount of valence electrons in each column)
2 numbering systems exist: Groups # I through VIII with ea. # followed by A or B
• A groups are Main Group Elements (s&p electrons)• B groups are Transition Elements (d electrons)
Numbered 1 to 18 from left to right
2. Periods
Horizontal rows of periodic table
Periods are numbered top to bottom from 1 to 7
Elements in same period have similarities in energy levels, but not properties
Main Group ElementsTransition MetalsInner Transition Metals
3. Blocks
3. Blocks
1
2
3
4
5
6
7
Lanthanides - part of period 6
Actinides - part of period 7
Overall Configuration
I II III
II. Classification of theElements(pages 182-186)
Ch. 6 - The Periodic Table
A. Metallic Character
1
2
3
4
5
6
7
MetalsNonmetalsMetalloids
1. Metals
Good conductors of heat and electricityFound in Groups 1 & 2, middle of table in
3-12 and some on right side of tableHave luster, are ductile and malleableMetallic properties increase as you go
from left to right across a periodForm bases in water
a. Alkali Metals
Group 1(IA)1 Valence electronVery reactive, form metal oxides
(ex: Li2O)
Electron configuration ns1
Lowest melting pointsForm 1+ ion: Cations
Examples: Li, Na, K
b. Alkaline Earth Metals
Group 2 (IIA)2 valence electronsReactive (not as reactive as alkali metals) form
metal oxides (ex: MgO)Electron Configuration
ns2
Form 2+ ionsCations
Examples: Be, Mg, Ca, etc
c. Transition Metals
Groups 3 – 12 (IB – VIIIB) Reactive (not as reactive as Groups 1 or 2), can
be free elements Highest melting pointsElectron Configuration
ns2(n-1)dx where x is column in d-blockForm variable valence state ionsAlways form Cations
Examples: Co, Fe, Pt, etc
3. Metalloids
Sometimes called semiconductorsForm the “stairstep” between metals and
nonmetalsHave properties of both metals and
nonmetalsExamples: B, Si, Sb, Te, As, Ge, Po, At
2. Nonmetals
Not good conductorsUsually brittle solids or gases (1 liquid Br)Found on right side of periodic table –
AND hydrogenHydrogen is it’s own group, reacts rapidly
with oxygen & other elements (has 1 valence electron)
Form acids in water
Nonmetal Groups/Families
Boron Group: IIIA typically 3 valence electrons, also mix of metalloids and metals
Carbon Group: IVA typically 4 valence electrons, also has metal and metalloids
Nitrogen Group: VA typically 5 valence electrons, also has metals & metalloids
Oxygen Group: VIA typically 6 valence electrons, also contains metalloids
a. Halogens
Group 17 (VIIA)Very reactiveElectron configuration
ns2np5
Form 1- ions – 1 electron short of noble gas configuration
Typically form salts (NaCl)Anions
Examples: F, Cl, Br, etc
b. Noble Gases
Group 18 (VIIIA)Unreactive, inert, “noble”, stableElectron configuration
ns2np6 full energy level Have an octet or 8 valence e-
Have a 0 charge, no ionsHelium is stable with 1s2, a duetExamples: He, Ne, Ar, Kr, etc
B. Chemical ReactivityMetals Period - reactivity decreases as you go from left to right across a period.
Group - reactivity increases as you go down a groupReact to form bases when combined with water Non-metals Period - reactivity increases as you go from the left to the right across a period.
Group - reactivity decreases as you go down the group. React to form acids when combined with water
C. Valence Electrons
Valence Electrons e- in the outermost s & p energy levels Stable octet: filled s & p orbitals (8e-) in one
energy level
1A
2A 3A 4A 5A 6A 7A
8A
C. Valence ElectronsYou can use the Periodic Table to determine
the number of valence electronsEach group has the same number of valence
electrons Group #A = # of valence e- (except He)
1A
2A 3A 4A 5A 6A 7A
8A
I II III
III. Periodic Trends(p. 187-194)
Ch. 6 - The Periodic Table
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
chemical and physical properties appear
at regular intervals.
0
50
100
150
200
250
0 5 10 15 20
Ato
mic
Ra
diu
s (
pm
)
Atomic Number
Atomic Radius size of atom
© 1998 LOGALIonization Energy
Energy required to remove an e- from a neutral atom
© 1998 LOGAL
Electronegativity
Properties of Atoms
Shielding Effect
There is a Nuclear charge experienced by the outer (valence) electron(s) in a multi-electron atom is due to the difference between the charge on the nucleus and the charge of the core electrons (inner electron shells). As atoms add more protons the nuclear charge increases Atoms are also adding more e- which are attracted to the p+
Results in the reduction of attractive force between the positive nucleus and the outermost electrons due to “shielding effect” of the inner electron shells(core electrons).
Periodic Trend,
1. Shielding effect increases down a group.
2. Shielding effect remains constant across a period.
Atomic Radius = ½ the distance between two identical bonded atoms
1. Atomic Radius
1
2
3
4 5
6
7
Atomic Radius Increases to the LEFT and DOWN
1. Atomic Radius
Why larger going down?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus and the valence e-
Why smaller to the right?
Increased nuclear charge(total charge of protons in nucleus) without additional shielding pulls e- in tighter
1. Atomic Radius
The minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. The ease with which an atom loses an e-.First Ionization Energy (IE1) = Energy required to remove one e- from a neutral atom.
Na(g) + IE1 (energy) → Na+(g) + e- ; +∆H (positive)
Second Ionization Energy (IE2) = energy needed to remove a second electron, and so forth
Na+(g) + IE2 (energy) → Na2+ (g) + e- ; +∆H (positive)
2. Ionization Energy
1
2
3
4 5
6
7
First Ionization Energy Increases UP and to the RIGHT
2. Ionization Energy
Why does it increase up a group?
The closer the e- are to the nucleus the more difficult it is to remove them
Decreased shielding effect increases the positive nuclear charge
Why does it increase across a period?
Atomic radius decreases
Positive nuclear charge increases pulling e- closer to the nucleus
2. Ionization Energy
Successive Ionization Energies
Mg 1st I.E. 736 kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ
Large jump in I.E. occurs when a CORE e- is removed.
The greater the IE the more difficult it is to remove an electrons
2. Ionization Energy
Al 1st I.E. 577 kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ
Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
2. Ionization Energy
Electron Affinity
Electron Affinity
The greater the attraction between a given atom and an added e-, the more negative the atom’s EA. Halogens’ ns2p5 have the most negative EA.
Noble Gases have EA > 0; as do Be, Mg, & N because e- have to enter previously unoccupied, higher energy orbitals, an unfavorable energy state.
Periodic Trend
1. Electron affinity slightly increases up a group.
2. Electron affinity generally tends to increase across a period.
Electron Affinity
Electron affinity increases up a groupdecreases the atomic radius taking the electrons
closer to the nucleus’ positive attraction. decreasing shielding effect increases the effective
positive nuclear charge (+) as additional shells are added and e- are held on tighter.
Electron affinity increases across a period atomic radius decreases effective positive nuclear charge increases steadily
and the e- are drawn closer to the nucleus making it easier to add e- to unfilled sublevels.
3. Electronegativity
The measure of the ability of an atom in a chemical compound to attract electrons
Given a value between 0 and 4, 4 being the highestTendency for an atom to attract e- closer to itself when forming
a chemical bond with another atom.
1
2
3
4 5
6
7
Why increase as you move right?
More valence electrons, need less to fill outer shell
Increased nuclear charge
Why increase as you move up?
Smaller electron cloud, more pull by + nucleus
3. Electronegativity
Ionic Radius
The size atoms become when losing or gaining electrons.
Positive Ions – Metal - Atoms that lose e- and form positive ions become smaller.
The lost e- is a valence e- and the atom may lose a shell.The repulsion between the remaining e- is lessened and allows the effective positive nuclear charge to pull the remaining e- closer.
Negative Ions – Nonmetal - Atoms that gain e- and form negative ions become larger.
The repulsion between the added e- and existing e- is increased and the effective positive nuclear charge cannot hold onto the e- tightly.
Periodic Trend
1. Ionic Radius increases down a group.
2. Ionic radius tends to gradually decrease across a period for the positive ions, then beginning in group VA or VIA the much larger negative ions also gradually decreases
Which atom has the larger radius?
Be or
or Br
Examples
Ba
Ca
Which atom has the higher 1st I.E.?
or Bi
Ba or
Examples
N
Ne
Which element has the higher electronegativity?
Cl or
or Ca
Examples
F
Be