Title: Chapter 8 Periodic Properties of the Elements
1Chapter 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
2Mendeleev
- order 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
3Periodic Pattern
nm O2 16 H2O
m Al2O3 a/b 27 (AlH3)
nm/m SiO2 a 28 SiH4
m metal, nm nonmetal, m/nm metalloid
a acidic oxide, b basic oxide, a/b
amphoteric oxide
4Mendeleev's Predictions
5What 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
- Quantum Mechanics is a theory that explains why
the periodic trends in the properties exist
6Electron 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
7Electron Spin Experiment
8Spin 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
9Pauli 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
10Allowed Quantum Numbers
11Quantum 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
12Electron 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
13Orbital 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
14Sublevel Splitting in Multielectron Atoms
- the sublevels in each principal energy level of
Hydrogen all have the same energy we call
orbitals with the same energy degenerate - or other single electron systems
- 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)
15Penetrating 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 - the result is that the electrons in the 2s
sublevel are lower in energy than the electrons
in the 2p
16Penetration Shielding
17Energy
- 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
18Order 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
19Filling 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
20Example 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
21Example 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
22Example 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
23Example 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
24Valence 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
25Electron 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)
26Electron Configurations
27Electron 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
28(No Transcript)
29s2
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
30Electron 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
31Transition Elements
- for the d block metals, the principal energy
level 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
32Electron 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
33Practice 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)
34Practice 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
35Properties 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
36Electron 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
37Electron 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
38(No Transcript)
39Electron 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
40Electron 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
41Trend 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
42Effective 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
43Screening Effective Nuclear Charge
44Trends 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
45(No Transcript)
46(No Transcript)
47Example 8.5 Choose the Larger Atom in Each
Pair
48Electron 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
49Magnetic Properties of Transition Metal Atoms
Ions
- electron configurations that result in unpaired
electrons mean that 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 a magnetic field
- both Zn atoms and Zn2 ions are diamagnetic,
showing that the two 4s electrons are lost before
the 3d - Zn atoms Ar4s23d10
- Zn2 ions Ar4s03d10
50Example 8.6 Write the Electron Configuration
and Determine whether the Fe atom and Fe3 ion
are Paramagnetic or Diamagnetic
- Fe Z 26
- previous noble gas Ar
- 18 electrons
51Trends 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
52Periodic Pattern - Ionic Radius (Ã…)
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
53(No Transcript)
54(No Transcript)
55Ionization Energy
- minimum energy needed to remove an electron from
an atom - gas state
- endothermic process
- valence electron easiest to remove
- M(g) IE1 ? M1(g) 1 e-
- M1(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.
56General 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
57(No Transcript)
58(No Transcript)
59Example 8.8 Choose the Atom in Each Pair with
the Higher First Ionization Energy
60Irregularities 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?
61Irregularities 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
62Irregularities 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
63Trends 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
64(No Transcript)
65Trends in Electron Affinity
- energy released when an neutral atom gains an
electron - gas state
- M(g) 1e- ? M-1(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
66(No Transcript)
67Metallic 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
68(No Transcript)
69(No Transcript)
70Example 8.9 Choose the More Metallic Element
in Each Pair
71Trends 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
722 Na(s) 2 H2O(l) ? 2 NaOH(aq) H2(g)
73Trends 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
74(No Transcript)
75Example 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)
76Example 8.10 Write a balanced chemical reaction
for the following.
- reaction between rubidium metal and liquid water
- Rb(s) H2O(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)
77Example 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)
78Trends 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 undersirable - 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
79(No Transcript)