Chapter 8 Periodic Properties of the Elements

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Chapter 8Periodic Properties of the Elements
Chemistry A Molecular Approach, 1st Ed.Nivaldo
Tro
Roy Kennedy Massachusetts Bay Community
College Wellesley Hills, MA
2007, Prentice Hall
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Discovery

• Elements are organized according to their
  physical and chemical properties in the Periodic
  Table
• Historically, several people contributed to this
  effort in the late 19th century
• Dobereiner- triads Ca,Sr,Ba Li, Na, K (1817)
• Newlands- similarity between every eighth element
  (1864)
• Mendeleev (1870) - Arranged elements according to
  atomic mass. Similar elements were arranged
  together in a group

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Question

• Dobereiners triads (1817)
• Determine the Atomic mass of Sr by averaging the
  masses of Ca and Ba
• 87.62 amu

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Newlands

• In 1863, he suggested that elements be arranged
  in octaves because he noticed (after arranging
  the elements in order of increasing atomic mass)
  that certain properties repeated every 8th element

Law of Octaves
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Newlands Law of Octaves

• Newlands (1864) arranged the elements as follows
• Did not contain the Noble gases (yet to be
  discovered)
• Not valid for elements with Z gt 20
• His table had 7 groups instead of 8

Li Be B C N O F
Na Mg Al Si P S Cl
K Ca
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Mendeleev

• ordered elements by atomic mass
• saw a repeating pattern of properties
• Periodic Law When the elements are arranged in
  order of increasing atomic mass, certain sets of
  properties recur periodically
• put elements with similar properties in the same
  column
• used pattern to predict properties of
  undiscovered elements
• where atomic mass order did not fit other
  properties, he re-ordered by other properties
• Te I

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Mendeleev's Predictions
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The Periodic Table

• Periodic Law - Both physical and chemical
  properties of the elements vary periodically with
  increasing atomic mass
• Exceptions Te/I, Ar/K, Co/Ni
• Placed Te (M 127.6) ahead of I (M126.9)
  because Te was similar to Se and S, and I was
  similar to Cl and Br
• Moseley showed listing elements based on atomic
  no. resolved these issues

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What vs. Why

• Mendeleevs Periodic Law allows us to predict
  what the properties of an element will be based
  on its position on the table
• it doesnt explain why the pattern exists
• Laws summarize behavior while theories explain
  them!
• Quantum Mechanics is a theory that explains why
  the periodic trends in the properties exist

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Electron Spin

• experiments by Stern and Gerlach showed a beam of
  silver atoms is split in two by a magnetic field
• the experiment reveals that the electrons spin on
  their axis
• as they spin, they generate a magnetic field
• spinning charged particles generate a magnetic
  field
• if there is an even number of electrons, about
  half the atoms will have a net magnetic field
  pointing North and the other half will have a
  net magnetic field pointing South

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Electron Spin Experiment
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Energy Shells and Subshells
n principal quantum no. (overall size) l
angular momentum quantum no. (shape of
orbital) ml magnetic quantum no. (orientation
of orbital)
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Spin Quantum Number, ms

• spin quantum number describes how the electron
  spins on its axis
• clockwise or counterclockwise
• spin up or spin down
• spins must cancel in an orbital
• paired
• ms can have values of ½

Orbital diagram for H
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Pauli Exclusion Principle

• no two electrons in an atom may have the same set
  of 4 quantum numbers
• therefore no orbital may have more than 2
  electrons, and they must have with opposite spins
• knowing the number orbitals in a sublevel allows
  us to determine the maximum number of electrons
  in the sublevel
• s sublevel has 1 orbital, therefore it can hold 2
  electrons
• p sublevel has 3 orbitals, therefore it can hold
  6 electrons
• d sublevel has 5 orbitals, therefore it can hold
  10 electrons
• f sublevel has 7 orbitals, therefore it can hold
  14 electrons

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Allowed Quantum Numbers
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Quantum Numbers of Heliums Electrons

• helium has two electrons
• both electrons are in the first energy level
• both electrons are in the s orbital of the first
  energy level
• since they are in the same orbital, they must
  have opposite spins

n l ml ms
first electron 1 0 0 ½
second electron 1 0 0 -½
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Electron Configurations

• the ground state of the electron is the lowest
  energy orbital it can occupy
• the distribution of electrons into the various
  orbitals in an atom in its ground state is called
  its electron configuration
• the number designates the principal energy level
• the letter designates the sublevel and type of
  orbital
• the superscript designates the number of
  electrons in that sublevel
• He 1s2

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Orbital Diagrams

• we often represent an orbital as a square and the
  electrons in that orbital as arrows
• the direction of the arrow represents the spin of
  the electron

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Sublevel Splitting in Multielectron Atoms

• the sublevels in each principal energy level (n
  2) of Hydrogen all have the same energy (empty
  in lowest state) called degenerate
• for multielectron atoms, the energies of the
  sublevels are split
• caused by electron-electron repulsion
• the lower the value of the l quantum number, the
  less energy the sublevel has
• s (l 0) lt p (l 1) lt d (l 2) lt f (l 3)

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Penetration Shielding
Consider bringing an e- close to Li 1s2
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Penetrating and Shielding

• the radial distribution function shows that the
  2s orbital penetrates more deeply into the 1s
  orbital than does the 2p
• the weaker penetration of the 2p sublevel means
  that electrons in the 2p sublevel experience more
  repulsive force, they are more shielded from the
  attractive force of the nucleus
• the deeper penetration of the 2s electrons means
  electrons in the 2s sublevel experience a greater
  attractive force to the nucleus and are not
  shielded as effectively by 1s e-
• the result is that the electrons in the 2s
  sublevel are lower in energy than the electrons
  in the 2p

2p is shielded
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Energy

• Notice the following
• because of penetration, sublevels within an
  energy level are not degenerate
• penetration of the 4th and higher energy levels
  is so strong that their s sublevel is lower in
  energy than the d sublevel of the previous energy
  level
• the energy difference between levels becomes
  smaller for higher energy levels

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Order of Subshell Fillingin Ground State
Electron Configurations
start by drawing a diagram putting each energy
shell on a row and listing the subshells, (s, p,
d, f), for that shell in order of energy,
(left-to-right)
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d
7s
next, draw arrows through the diagonals, looping
back to the next diagonal each time
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Filling the Orbitals with Electrons

• energy shells fill from lowest energy to high
• subshells fill from lowest energy to high
• s ? p ? d ? f
• Aufbau Principle
• orbitals that are in the same subshell have the
  same energy
• no more than 2 electrons per orbital
• Pauli Exclusion Principle
• when filling orbitals that have the same energy,
  place one electron in each before completing
  pairs
• Hunds Rule

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Example 8.1 Write the Ground State Electron
Configuration and Orbital Diagram and of
Magnesium.

• Determine the atomic number of the element from
  the Periodic Table
• This gives the number of protons and electrons in
  the atom
• Mg Z 12, so Mg has 12 protons and 12 electrons

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Example 8.1 Write the Ground State Electron
Configuration and Orbital Diagram and of
Magnesium.

• Draw 9 boxes to represent the first 3 energy
  levels s and p orbitals
• since there are only 12 electrons, 9 should be
  plenty

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Example 8.1 Write the Ground State Electron
Configuration and Orbital Diagram and of
Magnesium.

• Add one electron to each box in a set, then pair
  the electrons before going to the next set until
  you use all the electrons
• When pair, put in opposite arrows

??
??
?
??
?
?
?
?
?
1s
2s
2p
3s
3p
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Example 8.1 Write the Ground State Electron
Configuration and Orbital Diagram and of
Magnesium.

• Use the diagram to write the electron
  configuration
• Write the number of electrons in each set as a
  superscript next to the name of the orbital set
• 1s22s22p63s2 Ne3s2

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Valence Electrons

• the electrons in all the subshells with the
  highest principal energy shell are called the
  valence electrons
• electrons in lower energy shells are called core
  electrons
• chemists have observed that one of the most
  important factors in the way an atom behaves,
  both chemically and physically, is the number of
  valence electrons

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Electron Configuration of Atoms in their Ground
State

• Kr 36 electrons
• 1s22s22p63s23p64s23d104p6
• there are 28 core electrons and 8 valence
  electrons
• Rb 37 electrons
• 1s22s22p63s23p64s23d104p65s1
• Kr5s1
• for the 5s1 electron in Rb the set of quantum
  numbers is n 5, l 0, ml 0, ms ½
• for an electron in the 2p sublevel, the set of
  quantum numbers is n 2, l 1, ml -1 or
  (0,1), and ms - ½ or (½)

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Electron Configurations
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Electron Configuration the Periodic Table

• the Group number corresponds to the number of
  valence electrons
• the length of each block is the maximum number
  of electrons the sublevel can hold
• the Period number corresponds to the principal
  energy level of the valence electrons

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s2
1 2 3 4 5 6 7
d1 d2 d3 d4 d5 d6 d7 d8 d9 d10
f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12
f13 f14 f14d1
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Electron Configuration fromthe Periodic Table
8A
1A
1 2 3 4 5 6 7
3A
4A
5A
6A
7A
2A
Ne
P
3s2
3p3
P Ne3s23p3 P has 5 valence electrons
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Transition Elements

• for the d block metals, the principal energy
  level of the d orbital being filled is one less
  than valence shell
• one less than the Period number
• sometimes s electron promoted to d sublevel

Zn Z 30, Period 4, Group 2B Ar4s23d10

• for the f block metals, the principal energy
  level is two less than valence shell
• two less than the Period number they really
  belong to
• sometimes d electron in configuration

Eu Z 63, Period 6 Xe6s24f 7
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Electron Configuration fromthe Periodic Table
8A
1A
1 2 3 4 5 6 7
3A
4A
5A
6A
7A
2A
Ar
3d10
As
4s2
4p3
As Ar4s23d104p3 As has 5 valence electrons
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Practice Use the Periodic Table to write the
short electron configuration and orbital diagram
for each of the following

• Na (at. no. 11)
• Te (at. no. 52)
• Tc (at. no. 43)

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Practice Use the Periodic Table to write the
short electron configuration and orbital diagram
for each of the following

• Na (at. no. 11) Ne3s1
• Te (at. no. 52) Kr5s24d105p4
• Tc (at. no. 43) Kr5s24d5

3s
5s
5p
4d
5s
4d
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Properties Electron Configuration

• elements in the same column have similar chemical
  and physical properties because they have the
  same number of valence electrons in the same
  kinds of orbitals

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Electron Configuration Element Properties

• the number of valence electrons largely
  determines the behavior of an element
• chemical and some physical
• since the number of valence electrons follows a
  Periodic pattern, the properties of the elements
  should also be periodic
• quantum mechanical calculations show that 8
  valence electrons should result in a very
  unreactive atom, an atom that is very stable
  and the noble gases, that have 8 valence
  electrons are all very stable and unreactive
• conversely, elements that have either one more or
  one less electron should be very reactive and
  the halogens are the most reactive nonmetals and
  alkali metals the most reactive metals
• as a group

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Electron Configuration Ion Charge

• we have seen that many metals and nonmetals form
  one ion, and that the charge on that ion is
  predictable based on its position on the Periodic
  Table
• Group 1A 1, Group 2A 2, Group 7A -1,
  Group 6A -2, etc.
• these atoms form ions that will result in an
  electron configuration that is the same as the
  nearest noble gas

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Electron Configuration of Anions in their Ground
State

• anions are formed when atoms gain enough
  electrons to have 8 valence electrons
• filling the s and p sublevels of the valence
  shell
• the sulfur atom has 6 valence electrons
• S atom 1s22s22p63s23p4
• in order to have 8 valence electrons, it must
  gain 2 more
• S2- anion 1s22s22p63s23p6

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Electron Configuration of Cations in their Ground
State

• cations are formed when an atom loses all its
  valence electrons
• resulting in a new lower energy level valence
  shell
• however the process is always endothermic
• the magnesium atom has 2 valence electrons
• Mg atom 1s22s22p63s2
• when it forms a cation, it loses its valence
  electrons
• Mg2 cation 1s22s22p6

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Electron Configuration of Atoms in their Ground
State

• Electrons in some elements are promoted to form
  more stable configurations, e.g.
• Cu 29 electrons
• Expected 1s22s22p63s23p64s23d9
• Actual 1s22s22p63s23p64s13d10
• Cr 24 electrons
• Expected 1s22s22p63s23p64s23d4
• Actual 1s22s22p63s23p64s13d5

The usual reason given for this configuration is
that the half-filled or completely filled
d-orbitals are more stable as compared to those
which are not.
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Trend in Atomic Radius Main Group

• Different methods for measuring the radius of an
  atom, and they give slightly different trends
• van der Waals radius nonbonding
• covalent radius bonding radius
• atomic radius is an average radius of an atom
  based on measuring large numbers of elements and
  compounds
• Atomic Radius Increases down group
• valence shell farther from nucleus
• effective nuclear charge fairly close
• Atomic Radius Decreases across period (left to
  right)
• adding electrons to same valence shell
• effective nuclear charge increases
• valence shell held closer

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Effective Nuclear Charge

• in a multi-electron system, electrons are
  simultaneously attracted to the nucleus and
  repelled by each other
• outer electrons are shielded from full strength
  of nucleus
• screening effect
• effective nuclear charge is net positive charge
  that is attracting a particular electron
• Z is nuclear charge, S is electrons in lower
  energy levels
• electrons in same energy level contribute to
  screening, but very little
• effective nuclear charge on sublevels trend, s gt
  p gt d gt f
• Zeffective Z - S

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Screening Effective Nuclear Charge
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Trends in Atomic RadiusTransition Metals

• increase in size down the Group
• atomic radii of transition metals roughly the
  same size across the d block
• must less difference than across main group
  elements
• valence shell ns2, not the d electrons
• effective nuclear charge on the ns2 electrons
  approximately the same

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Example 8.5 Choose the Larger Atom in Each
Pair
N or F? C or Ge? N or Al? Al or Ge?
N is further left so larger
Ge is further down so larger
Al is further down and left so larger
Opposing trends
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Electron Configuration of Cations in their Ground
State

• cations form when the atom loses electrons from
  the valence shell
• for transition metals electrons, may be removed
  from the sublevel closest to the valence shell
• Al atom 1s22s22p63s23p1
• Al3 ion 1s22s22p6
• Fe atom 1s22s22p63s23p64s23d6
• Fe2 ion 1s22s22p63s23p63d6
• Fe3 ion 1s22s22p63s23p63d5
• Cu atom 1s22s22p63s23p64s13d10
• Cu1 ion 1s22s22p63s23p63d10

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Magnetic Properties of Transition Metal Atoms
Ions

• Ferromagnetism certain mateirals form permanent
  magnets
• electron configurations that result in unpaired
  electrons mean that in the presence of an
  externally applied magnetic field the atom or ion
  will have a net magnetic field this is called
  paramagnetism
• will be attracted to a magnetic field
• electron configurations that result in all paired
  electrons mean that the atom or ion will have no
  magnetic field this is called diamagnetism
• slightly repelled by an externally applied
  magnetic field
• both Zn atoms and Zn2 ions not magnetic
  (diamagnetic), showing that the two 4s electrons
  are lost before the 3d
• Zn atoms Ar4s23d10
• Zn2 ions Ar4s03d10

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Example 8.6 Write the Electron Configuration
and Determine whether the Fe atom and Fe3 ion
are Paramagnetic or Diamagnetic

• Fe Z 26 (elemental iron is magnetic!)
• previous noble gas Ar
• 18 electrons

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Trends in Ionic Radius

• Ions in same group have same charge
• Ion size increases down the group
• higher valence shell, larger
• Cations smaller than neutral atom Anions bigger
  than neutral atom
• Cations smaller than anions
• except Rb1 Cs1 bigger or same size as F-1 and
  O-2
• Larger positive charge smaller cation
• for isoelectronic species
• isoelectronic same electron configuration
• Larger negative charge larger anion
• for isoelectronic series

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Periodic Pattern - Ionic Radius (Å)
1A
2A 3A 4A 5A
6A 7A
3
1.71
0.68
0.31
1.40
1.33
0.23
1
0.51
2.12
1.81
0.66
0.97
1.84
0.62
1.96
1.33
0.99
1.98
2.22
0.81
0.71
1.13
2.20
2.21
1.47
0.84
0.95
1.69
1.35
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Ionization Energy

• minimum energy needed to remove an electron from
  an atom
• gas state
• endothermic process
• valence electron easiest to remove
• M(g) IE1 ? M(g) 1 e-
• M(g) IE2 ? M2(g) 1 e-
• first ionization energy energy to remove
  electron from neutral atom 2nd IE energy to
  remove from 1 ion etc.

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General Trends in 1st Ionization Energy

• larger the effective nuclear charge on the
  electron, the more energy it takes to remove it
• the farther the most probable distance the
  electron is from the nucleus, the less energy it
  takes to remove it
• 1st IE decreases down the group
• valence electron farther from nucleus
• 1st IE generally increases across the period
• effective nuclear charge increases

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Example 8.8 Choose the Atom in Each Pair with
the Higher First Ionization Energy
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Irregularities in the Trend

• Ionization Energy generally increases from left
  to right across a Period
• except from 2A to 3A, 5A to 6A

Which is easier to remove an electron from B or
Be? Why?
Which is easier to remove an electron from N or
O? Why?
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Irregularities in the First Ionization Energy
Trends
Be
Be
1s
2s
2p
1s
2s
2p
To ionize Be you must break up a full sublevel,
cost extra energy
When you ionize B you get a full sublevel, costs
less energy
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Irregularities in the First Ionization Energy
Trends
To ionize N you must break up a half-full
sublevel, cost extra energy
When you ionize O you get a half-full sublevel,
costs less energy
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Trends in Successive Ionization Energies

• removal of each successive electron costs more
  energy
• shrinkage in size due to having more protons than
  electrons
• outer electrons closer to the nucleus, therefore
  harder to remove
• regular increase in energy for each successive
  valence electron
• large increase in energy when start removing core
  electrons

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Trends in Electron Affinity

• energy released when an neutral atom gains an
  electron
• gas state
• M(g) 1e- ? M-(g) EA
• defined as exothermic (-), but may actually be
  endothermic ()
• alkali earth metals noble gases endothermic,
  WHY?
• more energy released (more -) the larger the EA
• generally increases across period
• becomes more negative from left to right
• not absolute
• lowest EA in period alkali earth metal or noble
  gas
• highest EA in period halogen

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Metallic Character

• Metals
• malleable ductile
• shiny, lusterous, reflect light
• conduct heat and electricity
• most oxides basic and ionic
• form cations in solution
• lose electrons in reactions - oxidized
• Nonmetals
• brittle in solid state
• dull
• electrical and thermal insulators
• most oxides are acidic and molecular
• form anions and polyatomic anions
• gain electrons in reactions - reduced
• metallic character increases left
• metallic character increase down

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Example 8.9 Choose the More Metallic Element
in Each Pair
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Trends in the Alkali Metals

• atomic radius increases down the column
• ionization energy decreases down the column
• very low ionization energies
• good reducing agents, easy to oxidize
• very reactive, not found uncombined in nature
• react with nonmetals to form salts
• compounds generally soluble in water ? found in
  seawater
• electron affinity decreases down the column
• melting point decreases down the column
• all very low MP for metals
• density increases down the column
• except K
• in general, the increase in mass is greater than
  the increase in volume

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2 Na(s) 2 H2O(l) ? 2 NaOH(aq) H2(g)
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Trends in the Halogens

• atomic radius increases down the column
• ionization energy decreases down the column
• very high electron affinities
• good oxidizing agents, easy to reduce
• very reactive, not found uncombined in nature
• react with metals to form salts
• compounds generally soluble in water ? found in
  seawater
• reactivity increases down the column
• react with hydrogen to form HX, acids
• melting point and boiling point increases down
  the column
• density increases down the column
• in general, the increase in mass is greater than
  the increase in volume

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Example 8.10 Write a balanced chemical reaction
for the following.

• reaction between potassium metal and bromine gas
• K(s) Br2(g) ?
• K(s) Br2(g) ? K Br?
• 2 K(s) Br2(g) ? 2 KBr(s)
• (ionic compounds are all solids at room
  temperature)

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Example 8.10 Write a balanced chemical reaction
for the following.

• reaction between rubidium metal and liquid water
• Rb(s) 2H2O(l) ?
• Rb(s) H2O(l) ? Rb(aq) OH?(aq) H2(g)
• 2 Rb(s) 2 H2O(l) ? 2 Rb(aq) 2 OH?(aq)
  H2(g)
• (alkali metal ionic compounds are soluble in
  water)

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Example 8.10 Write a balanced chemical reaction
for the following.

• reaction between chlorine gas and solid iodine
• Cl2(g) I2(s) ?
• Cl2(g) I2(s) ? ICl
• write the halogen lower in the column first
• assume 11 ratio, though others also exist
• 2 Cl2(g) I2(s) ? 2 ICl(g)
• (molecular compounds found in all states at room
  temperature)

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Trends in the Noble Gases

• atomic radius increases down the column
• ionization energy decreases down the column
• very high IE
• very unreactive
• only found uncombined in nature
• used as inert atmosphere when reactions with
  other gases would be undesirable

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Trends in the Noble Gases

• melting point and boiling point increases down
  the column
• all gases at room temperature
• very low boiling points
• density increases down the column
• in general, the increase in mass is greater than
  the increase in volume

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