Ch. 6 - The Periodic Table & Periodic Law
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Dmitri Mendeleev (1869, Russian)
 Organized elements by increasing ____________ _______________
 Elements with similar _______________________ were grouped together
 There were some discrepancies
 Deduced elements existed, but were ______________________ elements, their properties could be predicted
Henry Moseley (1913, British)
 Organized elements by increasing __________ ______________
 Resolved discrepancies in Mendeleev’s arrangement
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 This is the way the periodic table is arranged today!
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Modern Periodic Table
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Groups/Families
o ________________ columns of periodic table
o Each group contains elements with similar _______________ &
______________________ properties (same amount of _________________ electrons in each column)
o 2 numbering systems exist:
o Groups # I through VIII with ea. # followed by ____ or ____
 A groups are ________________ Group Elements (s&p electrons)
 B groups are _____________________ Elements (d electrons)
o Numbered 1 to 18 from _________ to _____________
Periods
 ________________________ rows of periodic table
 Periods are numbered top to bottom from _____ to ______
 Elements in same period have similarities in _______________ levels, but _______ properties
Blocks
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Main Group Elements
Transition Metals
Inner Transition Metals
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2
3
4
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6
7
Metallic Character
Metals
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1
2
3
4
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7
Good ________________ of heat and electricity
Found in Groups ___ & ___, ___________ of table in 3-12 and some on
_________ side of table
 Have luster, are ductile and _______________
 Metallic properties __________________ as you go from left to right
across a period
 Alkali Metals
o Group 1(_____)
o ___ Valence electron
o Very _______________, form metal oxides (ex: Li2O)
o Electron configuration _________________
o Lowest __________________ points
o Form 1+ ions: ________________________
Examples: Li, Na, K
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Alkaline Earth Metals
Group _______(IIA)
____ valence electrons
_____________ (not as reactive as alkali metals) form
metal oxides (ex: MgO)
Transition Metals
Groups 3 – 12 (IB – VIIIB)
Reactive (not as reactive as Groups 1 or 2), can be
__________ elements
 Highest _________________ points
 Electron Configuration
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Metalloids
 Sometimes called _________________________
 Form the “stairstep” between metals & nonmetals
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Electron Configuration _______
Form 2+ ions
Cations
Examples:__________
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ns2(n-1)dx
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Form ______________valence state ions
_____________ form Cations
Examples: ________________________
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Have properties of both ___________ and nonmetals
Examples:___________________________
where x is column in ______________
Nonmetals
 __________good conductors
 Usually brittle _______________ or _____________ (1 liquid Br)
 Found on right side of periodic table – AND ____________________
 Hydrogen is it’s own _______________, reacts rapidly with oxygen & other elements (has 1 valence electron)
Nonmetal Groups/Families
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o Boron Group: IIIA typically ___valence electrons, also mix of metalloids and ____________
o Carbon Group: IVA typically ___ valence electrons, also has metal and metalloids
o Nitrogen Group: VA typically ___valence electrons, also has metals & metalloids
o Oxygen Group: VIA typically ___valence electrons, also contains _______________
Halogens
Group 17 (VIIA)
 Typically form _____________ (NaCl)
Very _________________
 _______________
 Examples: F, Cl, Br, et
Electron configuration____________
Form ____ions – 1 e- short of noble gas configuration
Noble Gases
 Group 18 (VIIIA)
 ____________________, inert, “noble”, stable
 Electron configuration
o __________ full energy level
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Have an
or 8 valence eHave a 0 charge, _____ ions
Helium is stable with 1s2, a duet
Examples: He, Ne, Ar, Kr, etc
Ch. 6 Periodic Trends
Periodic Law: When elements are arranged in order of _____________ atomic __________, elements
with similar ___________ and physical properties appear at regular ________________.
Properties of Atoms:
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Shielding Effect
-There is a Nuclear charge experienced by the ________ (valence) electron(s) in a multi-electron atom is
due to the difference between the charge on the ___________ and the charge of the __________
electrons (inner electron shells).
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As atoms add more ____________ the nuclear charge __________________
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Atoms are also adding more e- which are _______________ to the p+
-Results in the reduction of attractive force between the ____________ nucleus and the outermost
electrons due to “______________________” of the inner electron shells(core electrons).
Periodic Trend,
1. Shielding effect increases _____________ a group.
2. Shielding effect remains _________________ across a period.
Atomic Radius: = ½ the ____________ between two identical bonded ______________
Periodic Trend: Increases to the LEFT and DOWN
Why larger going down?
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Higher energy levels have larger _____________
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Shielding - core e- block the attraction between
the nucleus and the ____________ e-
Why larger to the left?
Decreased nuclear _____________ charge(total charge of protons in nucleus) without additional shielding
pulls e- ________________
Ionization Energy:
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The______________ energy required to remove an electron from the ground state of an isolated
gaseous atom or ion.
The ease with which an atom _____________ an e-.
______________________________ = Energy required to remove one e- from a neutral atom.
o Na(g) + IE1 (energy) → Na+(g) + e- ; +∆H (positive)
Second Ionization Energy (IE2) = energy needed to remove a __________ electron, and so forth
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Na+(g) + IE2 (energy) → Na2+ (g) + e- ; +∆H (positive)
First Ionization energy: Increases UP and to the RIGHT
Why does it increase up a group?
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1
2
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4
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The _____________the e- are to the nucleus the more
___________________ it is to remove them
Decreased shielding effect ____________ the positive nuclear charge
Why does it increase across a period?
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Atomic radius ______________________
Positive nuclear charge increases, pulling e- _______________ to the nucleus
Electron Affinity:
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Most atoms can attract e- to form _____________charged ions
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The _________ that occurs when an e- is added to a gaseous atom or ion.
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The ease with which an atom _______ an e-.
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For most atoms, the energy _____________ when an e- is added. (in kJ/mol)
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Cl(g) + e- → Cl—(g) + EA (kJ/mol) ; -∆H (negative)
Periodic Trend
1. Electron affinity slightly _____________ up a group.
2. Electron affinity generally tends to ______________ across a __________________.
Why Electron affinity increases up a group
_____________ the atomic radius taking the electrons closer to the nucleus’ positive attraction.
decreasing shielding effect _________________ the positive nuclear charge (+) as additional shells
are added and e- are held on tighter.
Electron affinity increases across a period
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atomic radius _______________
effective positive nuclear charge increases steadily and the e- are drawn closer to the nucleus
making it easier to add e- .
ElectroNegativity:
 The measure of the ability of an atom in a chemical compound to attract _________________
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Given a value between 0 and 4, 4 being the ______________
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Tendency for an atom to attract e- closer to itself __________________________________
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Why increase as you move right?
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_____________________________________
Increased _____________ _______________
Why increase as you move up?
_____________ electron cloud, more pull by + nucleus
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Examples:
Which atom has the higher 1st I.E.?
Which atom has the larger radius?
-N
-Be
or
Ba
-Ca
or
Br
or
-Ba
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Which element has the higher
electronegativity?
-Cl or F
Chemical Reactivity
Bi
-Be or Ca
Ne
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Metals
Period - reactivity _____________________ as you go from left to right across a period.
Group - reactivity ______________________ as you go down a group
o React to form ___________ when combined with water.
o Non-metals
o Period - reactivity increases as you go from the ___________ to the _____________ across a period.
Group - reactivity ________________________ as you go down the group.
o React to form _____________ when combined with water.
Valence Electrons
 e- in the outermost
energy levels
 ______________ ________________: filled s & p orbitals (8e-) in one energy level
 You can use the Periodic Table to determine the number of valence electrons
 Each group has the _____________ number of valence electrons
 Group #A = # of valence e- (except He)