1. Chapter 8 Periodic Properties of the Elements

2. Electron Spin Experiment

3. 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”

4. 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

5. 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

6. Quantum Numbers of Helium’s 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 +½
second -½

7. 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

8. 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
unoccupied
orbital
orbital with
1 electron
orbital with
2 electrons

9. Sublevel 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) < p (l = 1) < d (l = 2) < f (l = 3)

10. 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
the result is that the electrons in the 2s sublevel are lower in energy than the electrons in the 2p

11. 6d 7s 5f 6p 5d 6s 4f 5p 4d 5s 4p 3d Energy 4s 3p 3s 2p 2s 1s
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 2s 2p 1s

12. 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
Hund’s Rule

13. Chapter 5: Periodicity and Atomic Structure
4/24/2017
Copyright © 2008 Pearson Prentice Hall, Inc.

14. Electron Configurations of Multielectron Atoms
Chapter 5: Periodicity and Atomic Structure
Electron Configurations of Multielectron Atoms
4/24/2017 H:
1s1
1 electron
s orbital (l = 0)
n = 1
He:
1s2
2 electrons
s orbital (l = 0)
n = 1
Lowest energy to highest energy
Electron configurations show the distribution of the electrons between the subshells.
Li:
1s2 2s1
1 electrons
s orbital (l = 0)
n = 2
Copyright © 2008 Pearson Prentice Hall, Inc.

15. 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

16. Examples For the following atom, write:
the Ground State Electron Configuration
Use short hand notation to write orbital Diagram
Determine the core electrons and valence electrons
Carbon
Sulfur
Potassium

17. Electron configuration of transition metal and atoms in higher energy state
For the following atom, write:
the Ground State Electron Configuration
Use short hand notation to write orbital Diagram
Determine the core electrons and valence electrons Cr Br Bi

18. 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

19. 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 > p > d > f
Zeffective = Z - S

20. Screening & Effective Nuclear Charge

21. Trends in Atomic Radius Transition 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

22. Example – Choose the Larger Atom in Each Pair
N or F
C or Ge
N or Al,
N or F
C or Ge,
N or F
C or Ge
N or Al
Al or Ge?
N or F,

23. Ionization 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-
M+1(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.

24. 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

25. 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 Rb+1 & Cs+1 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

27. 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
Al+3 ion = 1s22s22p6
Fe atom = 1s22s22p63s23p64s23d6
Fe+2 ion = 1s22s22p63s23p63d6
Fe+3 ion = 1s22s22p63s23p63d5
Cu atom = 1s22s22p63s23p64s13d10
Cu+1 ion = 1s22s22p63s23p63d10

28. Example 8.8 – Choose the Atom in Each Pair with the Higher First Ionization Energy
Al or S
As or Sb
N or Si
O or Cl?
Al or S
As or Sb
N or Si,
Al or S
As or Sb,
Al or S,

29. Irregularities in the Trend
Ionization Energy generally increases from left to right across a Period
except from 2A to 3A, 5A to 6A  N 1s 2s 2p  O Be  1s 2s 2p B 
Which is easier to remove an electron from B or Be? Why?
Which is easier to remove an electron from N or O? Why?

30. 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 B  1s 2s 2p  B+  1s 2s 2p
When you ionize B you get a full sublevel, costs less energy

31. Irregularities in the First Ionization Energy Trends 1s 2s 2p  N+  1s 2s 2p 
To ionize N you must break up a half-full sublevel, cost extra energy  O 1s 2s 2p   O+  1s 2s 2p
When you ionize O you get a half-full sublevel, costs less energy
Tro, Chemistry: A Molecular Approach

32. 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

33. Trends 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

34. 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

36. Example – Choose the More Metallic Element in Each Pair
Sn or Te
P or Sb
Ge or In
S or Br?

37. 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

38. 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

39. 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 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

40. Example– Write a balanced chemical reaction for the following.
reaction between potassium metal and bromine gas
K(s) + Br2(g) 
(ionic compounds are all solids at room temperature)
reaction between rubidium metal and liquid water
Rb(s) + H2O(l) 
reaction between chlorine gas and solid iodine
Cl2(g) + I2(s) 

41. Magnetic properties
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

42. Examples Al and Al3+ O and O-2 Ag and Ag+
Write the Electron Configuration and Determine whether the following atoms or their ions are Paramagnetic or Diamagnetic
Al and Al3+
O and O-2
Ag and Ag+