Part I - Introduction to the Periodic Table

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Unit 3.2 Part I - Introduction to the Periodic
Table
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• History of the Periodic Table

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The Search for a Periodic Table

• Until 1800s no clear system
• Elements grouped by similar properties or atomic
  mass

• In 1829, J.W. Döbereiner classified some elements
  into groups of three, which he called triads.

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Döbereiners Triads

• The elements in a triad had similar chemical
  properties, and their physical properties varied
  in an orderly way according to their atomic
  masses.

Element Atomic mass (g) Density (g/mL) Melting point (C) Boiling point (C)
Chlorine 35.5 0.00321 -101 -34
Bromine 79.9 3.12 -7 59
Iodine 127 4.93 114 185
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Döbereiners Triads
Element Atomic mass (g) Density (g/mL) Melting point (C) Boiling point (C)
Chlorine 35.5 0.00321 -101 -34
Bromine 79.9 3.12 -7 59
Iodine 127 4.93 114 185

• Density increases with increasing atomic mass.

• The concept of triads suggested that the
  properties of an element are related to its
  atomic mass.

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Which of the Dobereiner triads shown are still
listed in the same column of the modern periodic
table?
Triad 1 Triad 2 Triad 3
Li Mn S
Na Cr Se
K Fe Te

• Triad 1 and triad 3

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Mendeleevs Periodic Table

• The Russian chemist, Dmitri Mendeleev, developed
  a periodic table of elements.

• organized the elements according to increasing
  atomic mass.

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Mendeleevs Periodic Table

• Mendeleev later developed an improved version of
  his table with the elements arranged in
  horizontal rows.

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Mendeleevs Periodic Table

• Patterns of changing properties repeated for the
  elements across the horizontal rows.

• Elements in vertical columns have similar
  properties.

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Mendeleevs Periodic Table

• properties of the elements repeat in an orderly
  way from row to row of the table.

• This repeated pattern is an example of
  periodicity in the properties of elements.

• Periodicity is the tendency to recur at regular
  intervals.

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Mendeleevs Periodic Table

• In order to group elements with similar
  properties in the same columns, Mendeleev had to
  leave some blank spaces in his table.

• He suggested that these spaces represented
  undiscovered elements.

Mendeleev correctly predicted the properties
of several undiscovered elements. Why is this
important?
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Mendeleevs Periodic Table
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What are two factors that contributed to the
acceptance of Mendeleevs periodic law?

• Grouping of elements with similar chemical
  properties
• Ability to predict properties of undiscovered
  elements

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The Modern Periodic Table

• the basis for ordering the elements in the table
  is the atomic number, not atomic mass.

• The atomic number of an element is equal to the
  number of protons in the nucleus.

• Each row (except the first) begins with a metal
  and ends with a noble gas.

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The Modern Periodic Table

• In between, the properties of the elements change
  in an orderly progression from left to right.

• This regular cycle illustrates periodicity in the
  properties of the elements.

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The Modern Periodic Table

• periodic law - physical and chemical properties
  of the elements repeat in a regular pattern when
  they are arranged in order of increasing atomic
  number

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Use the periodic table to separate these 12
elements into 6 pairs fo elements having similair
properties.Ca, K, Ga, P, Si, Rb, B, Sr, Sn, Cl,
Bi, Br
Ca K Ga P Si Cl
Sr Rb B Bi Sn Br
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• Layout of the Periodic Table

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Layout of the periodic table

• A group, also called a family, consists of the
  elements in a vertical column.

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Groups are numbered 1 18 ORIA VIIIA for
main group elements andIB VIIIB for transition
elements
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As you move left to right across a period the
number of valence electrons increases by one
1 valence e-
3 valence e-
2 valence e-
4 valence e-
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Elements in the same group have same number of
valence electrons and similar properties
1 valence e-
3 valence e-
2 valence e-
4 valence e-
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• A period consists of the elements in a horizontal
  row

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Periods are numbered 1-7 and each new row begins
a new energy level
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• The elements in the middle are called transition
  elements

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The others are main group elements
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• Lithium is the first element in Group 1 and in
  Period 2. Check this location on the periodic
  table.

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• 4 groups have commonly used names alkali metals
  in Group 1 (IA)

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• the alkaline earth metals in Group 2 (IIA)

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• the halogens in Group 17 (VIIA) -from the Greek
  words for salt former , compounds that halogens
  form with metals are salt-like.

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the noble gases in Group 18 (VIIIA) full outer
shell (8 valence electrons), generally unreactive
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• In the periodic table, two series of elements are
  placed below the main body of the table.

The elements in these two series are known as the
inner transition elements.
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• The first series of inner transition elements is
  called the lanthanides because they follow
  element number 57, lanthanum.

• Because of their natural abundance on Earth is
  less than 0.01 percent, the lanthanides are
  sometimes called the rare earth elements.

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• The second series of inner transition elements
  are the actinides

• All of the actinides are radioactive, and all
  beyond uranium (92) are man made (synthetic).

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• Elemental Funkiness
• By Mark Rosengarten UF

• http//www.youtube.com/watch?v1PSzSTilu_s

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• Classification of elements

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• Elements are classified as metals, metalloids, or
  nonmetals on the basis of their physical and
  chemical properties.

• The majority of the elements are metals (solids).
  They occupy the entire left side and center of
  the periodic table.

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Physical States and Classes of the Elements

• Nonmetals occupy the upper-right-hand corner.
  green, yellow, orange

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Physical States and Classes of the Elements

• Metalloids are located along the staircase
  boundary between metals and nonmetals. - purple

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Metals

• Metals are elements that have luster, conduct
  heat and electricity, and usually bend without
  breaking.

• All metals except mercury are solids at room
  temperature in fact, most have extremely high
  melting points.

Click box to view movie clip.
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Metals

• With the exception of tin, lead, and bismuth,
  metals have one, two, or three valence electrons.

• The periodic table shows that most of the metals
  (coded blue) are not main group elements.

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Nonmetals

• Most nonmetals dont conduct electricity, are
  much poorer conductors of heat than metals, and
  are brittle when solid.

• Their melting points tend to be lower than those
  of metals.

• Many are gases at room temperature

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• With the exception of carbon, nonmetals have
  five, six, seven, or eight valence electrons.

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Properties of Metals and Nonmetals
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Metalloids

• Metalloids have some chemical and physical
  properties of metals and other properties of
  nonmetals. - purple

• In the periodic table, the metalloids lie along
  the border between metals and nonmetals.

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some metalloids are semiconductors

• A semiconductor is an element that does not
  conduct electricity as well as a metal, but does
  conduct slightly better than a nonmetal.

• Some metalloids such as silicon, germanium (Ge),
  and arsenic (As) are semiconductors.

• Silicons semiconducting properties made the
  computer revolution possible.

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Part II - Periodic Trends
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Periodic Properties of the Elements

• The electron structure of an atom determines many
  of its chemical and physical properties.

• Understanding the relationship between electron
  configuration and position in the periodic table
  enables you to predict the properties of the
  elements and the outcome of many chemical
  reactions.

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• Atomic Radius
• size of atom

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Atomic Size

• size of an atom INCREASES in any group as you go
  DOWN the column because the valence electrons are
  in energy levels farther from the nucleus.

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• shielding effect electrons in energy levels
  closer to the nucleus shield the valance
  electrons from the positive pull of the nucleus

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• The shielding effect
• Increases down a group because electrons are
  being added to higher energy levels
• There is no shielding effect as you go across a
  period because electrons are being added to the
  same principal energy level

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• size of an atom DECREASES in any period as you
  go to the RIGHT in any row because there is an
  increased nuclear () charge pulling e- in
  tighter.

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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 without additional
  shielding pulls e- in tighter

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Atomic Radii of Main Group Elements
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Examples

• Which atom has the larger radius?

• Be or Ba
• Ca or Br

Ba Ca
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For each of the following pairs, predict which
atom is larger.
a. Mg, Sr
Sr
b. Sr, Sn
Sr
c. Ge, Sn
Sn
d. Ge, Br
Ge
e. Cr, W
W
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• Octet Rule
• reactivity of atoms is based on achieving a
  complete octet of valence electrons(8/8)
• Everybody wants to be like a noble gas!

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Atoms achieve noble gas configuration by gaining
or losing their valence electronsAn ion is an
atom or group of atoms that has a charge because
of the loss or gain of electrons.
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cation - An ion that has LOST an e- and now has a
positive () chargeanion an ion that has
GAINED an e- and now has a negative (-) charge
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Common Ion Charges aka oxidation number
1
0
1-
3
3-
2-
2
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• GROUP VALENCE WHEN FORMING IONS
• OUT OF 8
• Group IA 1 loses 1
• Group IIA 2 loses 2
• Group IIIA 3 loses 3
• Group IVA 4 can lose or gain
• Group VA 5 gains 3
• Group VIA 6 gains 2
• Group VIIA 7 gains 1
• Group VIIIA 8 does not form ions

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Ionic Size

• positive ions (cations) acquire the
  configuration of the noble gas in the preceding
  period.

• the outermost electrons of the ion are in a lower
  energy level than the valence electrons of the
  neutral atom.

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• The electrons that are not lost by the atom
  experience a greater attraction to the nucleus
  and pull together in a tighter bundle with a
  smaller radius.

• all cations ions have smaller radii than their
  corresponding atoms.

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• anions acquire the electron configuration of the
  noble gas at the end of its period.

• But the nuclear charge doesnt increase with the
  number of electrons.

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• In the case of fluorine, a nuclear charge of 9
  must hold ten electrons in the F ion all the
  electrons are held less tightly

• the radius of the anion is larger than the
  neutral atom.

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Ionic Radius

• Ionic Radius

• Cations ()
• lose e-
• smaller

• Anions ()
• gain e-
• larger

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Examples

• Which particle has the larger radius?

• S or S2-
• Al or Al3

S2- Al
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For each of the following pairs, predict which
atom or ion is larger
a. Mg, Mg2
Mg
b. S, S2
S2
c. Ca2, Ba2
Ba2
d. Cl, I
I
e. Na, Al3
Na
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• ionization energy - the energy needed to REMOVE
  an electron from an atom, in kJ/mol

• First Ionization Energy
• Energy required to remove the 1st e- from a
  neutral atom.

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• group trends
• (first) ionization energy decreases from top to
  bottom along a group
• reason outermost electron is farther and farther
  from the nucleus in larger atoms, so it is more
  easily removed

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• periodic trends
• (first) ionization energy increases from left to
  right in a period
• reason ?nuclear charge() increases more
  attraction between electrons and protons

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• Successive Ionization Energies

• Large jump in I.E. occurs when a CORE e- is
  removed.

• Mg 1st I.E. 736 kJ
• 2nd I.E. 1,445 kJ
• Core e- 3rd I.E. 7,730 kJ

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• Successive Ionization Energies

• Large jump in I.E. occurs when a CORE e- is
  removed.

• 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

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Examples

• Which atom has the higher 1st I.E.?

• N or Bi
• Ba or Ne

N Ne
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For each of the following pairs, predict which
atom has the higher first ionization energy.
a. Mg, Na
Mg
O
b. S, O
c. Ca, Ba
Ca
Cl
d. Cl, I
e. Na, Al
Al
f. Se, Br
Br
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Periodic Trends in Electronegativity

• electronegativity tendency of an atom to attract
  electrons.
• noble gases do not have electronegativity values
• chemical bonds are determined by
  electronegativity differences between the bonding
  partners

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• electronegativity trends are not completely
  regular
• fluorine most electronegative element with a
  value of 4.0 (smallest anion formed)
• cesium least electronegative element (largest
  cation formed)

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• electronegativity decreases from top to bottom in
  a group

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• electronegativity increases from left to right in
  a period

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RECAP
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Atomic Radius

• Atomic Radius

• Increases to the LEFT and DOWN

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• First Ionization Energy

• Increases UP and to the RIGHT

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• electronegativity

• Increases UP and to the RIGHT

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• Electron affinity

• Increases UP and to the RIGHT

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Part 3 Electron Configuration
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Electrons in Atoms

• Niels Bohr.
• Electrons are arranged in orbits around the
  nucleus
• The energy level of an electron is the region
  around the
• nucleus where the
• electron is likely to
• be moving.

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modern 3-D electron-cloud model - probability
model

• Heisenberg Uncertainty Principle it is not
  possible to know both the exact position and
  velocity of an object simultaneously

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Modern electron cloud model

• orbitals are areas of high probability (95) of
  finding electrons

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• Electrons can change energy level, by absorbing
  energy. When an electron absorbs a quantum of
  energy, it moves up to a higher energy level.
• When the electron falls from a higher energy
  level to a lower energy level, energy is
  released, and we see light

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• Energy levels have sublevels divisions within
  an energy level
• 1) many similar energy states grouped together in
  a level
• 2) different shapes spherical, dumbbell,
  cloverleaf

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• There are 4 sublevels
• s, p, d, f
• (s p d f stand for sharp, principal, diffuse,
  fundamental)
• maximum number of e- in a principal
• energy level 2n 2
• n principal quantum number
• electron energy level or ? shell (period)
  number
• n 1, 2, 3, 4, 5, 6, 7

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• electron maximums in the sublevels
• s can hold 2 e-
• p can hold 6 e-
• d can hold 10 e-
• f can hold 14 e-

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• Electrons fill orbitals in a certain way
• electron configuration - a specific electron
  arrangement in orbitals

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Electron configurationGeneral Rules

• Pauli Exclusion Principle
• Each orbital can hold 2 electrons with opposite
  spins.

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• Aufbau Principle
• Electrons fill the lowest energy orbitals
  first.
• Lazy Tenant Rule

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• Hunds Rule
• Within a sublevel, place one e- per orbital
  before pairing them.
• Empty Bus Seat Rule

WRONG
RIGHT
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Different sections of the periodic table
correspond to the different sublevels Groups IA
IIA s block Groups IIIA VIIIA p
block Transition d block Inner transition f
block
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Diagonal rule -to help us remember the order in
which energy level subshells fill - follow the
arrows

• 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f
  6s 6p 6d 6f7s 7p

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Example

• Orbital Diagram

O 8e-

• Electron Configuration

1 s2 2s2 2p4
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1s2 2s2 2p4

• the sum of the superscripts the atomic number
  of the element
• superscripts are NOT exponents (nothing is being
  squared, etc.)

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• valence configurations will be
• s OR s and p

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• Condensed (Abbreviated) Electron Configurations
• use the previous Noble Gas as the starting point
  in brackets, then finish the configuration

• Longhand Configuration

1s2
2s2
2p6
3s2
3p4
S 16e-

• Shorthand Configuration

S 16e- Ne 3s2 3p4
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• Example Indium 49
• 1) complete 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
  5p1
• 2) condensed Kr 5s2 5p1
• 3) valence 5s2 5p1

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• I Heart Electron Configuration
• by Mark Rosengarten UF

http//www.youtube.com/watch?vVb6kAxwSWgU
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• THE END !