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The Periodic Table Flashcards

1
Q

Periodic Law

A

chemical/physical properties of elements are dependent in a periodic fashion based off of atomic numbers

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2
Q

The seven periods represent

A

seven rows based on quantum numbers n=1-7 for s and p blocked elements

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3
Q

Groups/Families

A

columns

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4
Q

Elements with the same electronic config in valence shell share…

A

similar chemical properties

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5
Q

Why valence electrons form chemical bonds…

A

because they are less tightly held by nucleus and hold high potential energy

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6
Q

A elements

A

representative elements

IA - VIIA

orbitals in s or p shells

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7
Q

B Elements

A

non representative elements

transition elements, lanthanide and actinide series

s & d, s & f subshells

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8
Q

Metals are found in the _

A

left and middle side of the periodic table

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9
Q

Metals contain the following:

A

active metals , transition metals , lanthanide and actinide

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10
Q

Metal Properties

A 
Lustrous
High mp and density
Maleable
Ductile 
Good Conductor
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11
Q

Atomic Metal Properties

A 
Low effective nuclear charge
low electronegativity 
large atomic radius 
small ionic radius 
low ionization energy
low electron affinity

all together make metals easy to give up electrons!!!!

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12
Q

Oxidation States

A

charges formed when forming bonds with other atoms

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13
Q

Metals are good conductors because..

A

they have free moving electrons

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14
Q

Active Metal Subshell

A

s

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15
Q

Transition Metal Subshells

A

s and d

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16
Q

Lanthanide and Actinide Subshells

A

s and f

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17
Q

Nonmetals are found predominantly on the

A

upper right side of periodic table

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18
Q

Nonmetal Properties

A

Brittle as solids
Little to no metallic luster
Poor Conductors

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19
Q

Nonmetal Atomic Properties

A

small atomic radii (opposite of metals)
large ionic radii (opposite of metals)

Hard to give up electrons!!!

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20
Q

Metalloid Properties

A

electronegativity/ionization between that of metals (loq) and nonmetals (high)

others (density, mp, etc) vary widely

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21
Q

Lanthanides and Actinides are located…

A

separated bottom 2 rows respectively

22
Q

All the Nonmetals:

A

Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium, Astatine

23
Q

Moving left to right on periodic table…

A

electrons and protons are added one at a time

effective nuclear charge (Zeff) increases

atomic radii decreases

24
Q

As the POSITIVITY of the nucleus increases…

A

the electrons around the nucleus experience a STRONGER ELECTROSTATIC PULL toward center of atom

electron cloud moves closer ot binds more tightly to nuclues

25
Effective Nuclear Charge (Zeff)
electrostatic attraction net positive charge experienced by outermost electron increases left to right!!! constant up and down!
26
Electrons inbetween nucleus and outermost electron...
can cancel some nuclear charge
27
Moving down on the periodic table....
principle quantum number increases by 1 each time (number of shells increases) reduction in electrostatic attraction reduced positivity of nucleus valence electrons held less tightly
28
Going down the table, increased shielding caused by inner shell electrons causes...
cancellation of increased positivity of nucleus hence Zeff of a group is constant!!!
29
Atomic radius
1/2 the distance between the centers of two atoms briefly in contact with one another decreases left to right as electrons are pulled more inward increases going down the periodic table
30
2 things to generalize
1) metals lose electrons and become more positive and non metals gain electrons to become more negtive 2) metalloids act either way with a tendency to act more like one dependent on what side of the line they fall on
31
Nonmetals close to the metalloid possess...
larger ionic radii than their counterparts in VIIIA
32
Metals close to the metalloids possess...
smaller radii than other metals
33
Ionization Energy (IE)
energy required to remove electron from a gas endothermic process; increases from left to right, bottom to top
34
Why are some metals 'active'?
they contain low ionization energies that they are not found in neutral state
35
First/Second Ionization Energies
Energy required to remove the first and second electrons respectively
36
Electron Affinity
opposite of IE energy dissipated (exothermic) when a gas gains an electron increases left to right, bottom to top and is the opposite sign of IE
37
Electronegativity
measure of attractive force an atom will exert on an electron in chemical bond
38
Electronegativity is ___ related to ionization energy
directly! lower , lower or higher, higher
39
Pauling Electronegativity Scale
0.7 - 4.0 (least electronegative to most electronegative)
40
Left to Right Trends Periodic Table
``` ionization energy increases electron affinity increases electronegtavity increases effective nuclear charge (Zeff) increases atomic radii decreases ```
41
Bottom Up Trends
``` ionization energy Increases electron affinity Increases electronegtativity increases effective nuclear charge (Zeff) constant atomic radii decreases ```
42
Chemistry Groups
``` Alkali Metals Alkaline Earth Metals Chalcogens Halogens Noble Gases Transition Metals ```
43
Alkali metals (Group 1)
lower densities than other metals one valence electron lowest Zeff , electron affinity/negativity and ionization energy highest atomic radius active metals
44
Alkaline Earth Metals (Group 2)
like all metals but higher Zeff and smaller atomic radii active metals
45
Chalcogens (Group 16)
nonmetals and metalloids six valence electrons small atomic radii and large ionic radii since close to metalloids biological importance
46
Halogens (17)
highly reactive nonmetals 7 valence electrons physical properties variable high electron affinity/negativity reactive towards alkali and alkaline earth metals
47
Noble Gases (18)
inert with no chemical reactivity high ionization energies, but no tendency to lose or gain electrons and no electronegativity low boiling points and gases
48
Transition Metals (Groups 3-12)
low electron affinities, ionization energies and electronegativities hard, high melting and boiling points, malleable, conductos loose electrons in d orbitals different possible charged forms/oxidation states to form many ionic compounds
49
Color is seen because...
it is reflected
50
Complementary color
the color that is complementary to the color of the wavelength absorbed; this is what is seen ex: blue gets absorbed so we see yellow

Chemistry (9 decks)

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