1. Before we start…
On your formula chart, highlight the formula used for calculating wavelength and frequency.
Highlight the formula used for calculating energy.
Highlight the speed of light constant and Plancks constant
On the periodic table, color the s,p,d,f block (each in different color)

2. Pre-AP Review: Light 1. Draw the EM spectrum.
2. Draw a blue wave. Describe the frequency, wavelength, energy of this wave.
3. Draw a red wave. Describe the frequency, wavelength, energy of this wave.
4. What formula do you use to calculate wavelength? What is the constant for this formula?
5. What formula do you use to calculate energy? What is the constant for this formula?
6. What are the units for frequency, wavelength, and energy?
7. How would you convert nanometers to meters for the following number: 589 nm

3. Atomic Structure and Periodicity
AP Chemistry: Chapter 7
Atomic Structure and Periodicity

4. Atomic Structure
In this unit, we will ask ourselves the same question that scientists asked once they believed in the atom; “What is the nature of the atom?” We will look not only at how the atom is structured, but also how this relates to the arrangement of the periodic table.

5. Electromagnetic Radiation
radiant energy that exhibits wavelike behavior and travels through space at the speed of light in a vacuum.

6. Waves have three Primary Characteristics:
Wavelength () -distance between two consecutive peaks or troughs in a wave.
Frequency () -number of waves per second that pass a given point in space.
Speed – all EM radiation travels at the speed of light in a vacuum

7. As the wavelength decreases, frequency increases.
 measured in m
 measured in s-1 or Hz
c speed of light 3.0x108 m/s

8. Ex. The yellow light given off by a sodium lamp has a wavelength of 589 nm. What is the frequency of the radiation?
589 nm 1 m = x 10-7m
109nm
= c/
= x 108m/s = 5.09 x 1014s-1
5.89 x 10-7m

9. ΔE =h  measured in cycles/second, simplified to s-1 or Hz
Until the early 1900’s, it was believed that matter and energy were very different. Matter was composed of particles and energy was composed of waves.
In 1901, Max Planck found that when solids were heated strongly, they absorbed and emitted energy.
He determined that energy can be gained or lost only in integer multiples of h.
h is Planck’s constant = x J.s
This showed that energy was quantized or made of “packets”.
7.2 The Nature of Matter
ΔE =h
 measured in cycles/second, simplified to s-1 or Hz

10. Ex. Calculate the smallest increment of energy (the quantum of energy) that an object can absorb from yellow light whose wavelength is 589 nm.
E = h

11. Calculation Review
1. Determine the wavelength of a photon whose frequency is 3.55x1017 Hz.
2. Find the energy of a red photon with a frequency of 4.57  1014 Hz
3. The yellow light given off by a sodium lamp has a wavelength of 589 nm. What is the frequency of the radiation? (1 nm = 10-9 m)
4. Calculate the smallest increment of energy (the quantum of energy) that an object can absorb from yellow light whose wavelength is 589 nm.

12. Review Warm-up 1. What is a line emission spectrum?
2. How do atoms transition between ground state and excited state?
3. What are the three rules that govern electron configuration?
4. Write the electron configuration for: sodium, germanium, oxygen, and helium
5. Write the orbital notation for calcium and nitrogen
6. Write the noble gas notation for rubidium, iron, and magnesium
*if you don’t remember how to do any of the above, work with a partner
* you have 15 minutes

13. The next 19 (blue) slides are a review of Pre-AP and why we see light
The next 19 (blue) slides are a review of Pre-AP and why we see light. (THEY ARE NOT ON YOUR NOTES FOR CHAPTER 7).
Recall: You did the flame test lab during this unit. You need to know why atoms burn a certain color and be able to explain what a line emission spectrum is. We will see this again in the next few units of this chapter.
Also Recall: you need to know how to write the electron configuration for this unit. There are three rules you need to remember.

14. Some vocab… Ground state-lowest energy state of an atom
Excited state-a state in which an atom has a higher potential energy that it has in its ground state
So, how do atoms transition between their ground state and their excited state???

15. Line-Emission Spectrum
excited state
ENERGY IN
PHOTON OUT
ground state
Emission spectrum of H2 gas

16. Line Emission Spectra
Classical theory-atoms would be excited by any amount of energy added to them, and should give off a continuous spectrum of EM radiation.
Attempts to explain further developed the quantum theory of the atom, and led to the Bohr model of the hydrogen atom

17. Bohr Model Energy of photon depends on the difference in energy levels6
Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1

18. Other Elements
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen! 

19. Electron Configuration
Specifies the “address” of each electron in an atom
UPPER LEVEL

20. General Rules Pauli Exclusion Principle
Each orbital can hold TWO electrons with opposite spins.

21. General Rules Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”

22. General Rules WRONG RIGHT Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
“Empty Bus Seat Rule”
WRONG
RIGHT

23. Representing e- configurations
3 ways to represent e- configurations
Orbital notation
Electron configuration notation
Noble gas notation

24. Some vocab…
Valence electrons-electrons that are found in the highest occupied energy level
Highest occupied energy level-electron-containing main energy level with the highest principal quantum number
Inner shell or core electrons-electrons that are not in the highest occupied energy level

25. Noble Gas Notation “Abbreviated” form of e- configuration notation
Use the noble gas (Gp. 18) that immediately precedes the element of interest as shorthand for the majority of the notation
Write the e- configuration notation for Ne and compare it to S. What do you notice?

26. S 16e- 1s2 2s2 2p6 3s2 3p4 Ne 10e- 1s2 2s2 2p6 S 16e- [Ne] 3s2 3p4
Noble Gas Notation
S 16e-
1s2
2s2
2p6
3s2
3p4
Ne 10e-
1s2
2s2
2p6
Only difference is the valence electrons in sulfur. Use [Ne] as an abbreviation for all parts that are the same!
S 16e- [Ne] 3s2 3p4

27. Stability Full energy level Full sublevel (s, p, d, f)
Half-full sublevel

28. Exceptions to AUFBAU principle
Copper
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.

29. Exceptions to Aufbau principle
Chromium
EXPECT: [Ar] 4s2 3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel and decreased electron repulsion

30. Configuration of Ions 1+ 2+ 3+ NA 3- 2- 1- Ion Formation
Atoms gain or lose electrons to become more stable.
Isoelectronic with the Noble Gases. 1+ 2+ 3+ NA 3- 2- 1-

31. Configuration of Ions O2- 10e- [He] 2s2 2p6
Writing Ion Electron Configuration
Write the e- config for the closest Noble Gas
EX: Oxygen ion  O2-  Ne
O e [He] 2s2 2p6

32. Practice
1. What element has the electron configuration of 1s22s22p63s23p1. Box the highest energy level and circle the valence electrons.
2. Write the electron configuration for Mn, Cl, and Sn.
2. Draw the orbital notation AND the noble gas notation for the following P, Ag, and C.
3. Write the electron configuration for the nitrogen ion, the cadium ion, and the potassium ion.
4. Write an isoelectronic series which includes argon.
5. Using your phone, google the photoelectric effect and explain this concept AND explain ionization energy

33. PES graph activity
Pick up a piece of paper, glue, and scissors from the side table
Cut out the elements and graphs from the paper you picked up.
Glue the matching element and graph next to each other.
Write the electron configuration, the orbital notation, and the noble gas notation for each element.
Label the peaks of each graph.
Draw a Bohr model for each element.

34. On the clean piece of paper that you picked up, draw the following…
Trend
Definition
Why trend for group
Why trend for period/row
Exceptions Explained
Atomic Radius
N/A
Ionization Energy (I.E.)
*Group 2 to 13
*Group 15 to 16
Electron Affinity
Electronegativity
Ionic Radius:
Cation/Anion

35. #1 OLD AP QUESTION: 1s22s22p63s23p3
Atoms of an element, X, have the electronic configuration shown above. The compound most likely formed with magnesium.
(A)MgX
(B) Mg2X
(C) MgX2
(D) Mg2X3
(E) Mg3X2

36. #2Questions 6-9 (OLD AP QUESTION) (A) 1s22s22p53s23p5
(B) 1s22s22p63s23p6
(C) 1s22s22p62d103s23p6
(D)1s22s22p63s23p63d5
(E)1s22s22p63s23p63d34s2
An impossible electron configuration
The ground-state configuration for the atoms of a
transition element
The ground-state configuration of a negative ion of
a halogen
The ground-state configuration of a common ion
of an alkaline earth element C D B E

37. # 3. Which of the following conclusions can be drawn from J. J
# 3. Which of the following conclusions can be drawn from J. J. Thomson's cathode ray experiments?
(A) Atoms contain electrons. (B) Practically all the mass of an atom is contained in its nucleus. (C) Atoms contain protons, neutrons, and electrons. (D) Atoms have a positively charged nucleus surrounded by an electron cloud. (E) No two electrons in one atom can have the same four quantum numbers.

38. #4 (OLD AP QUESTION)
Which of the following represents the ground state electron configuration for the Mn3+ ion? (Atomic number Mn = 25)
(A)1s22s22p63s23p63d4
(B)1s22s22p63s23p63d54s2
(C)1s22s22p63s23p63d24s2
(D)1s22s22p63s23p63d84s2
(E)1s22s22p63s23p63d34s1

39. #5 Which of the following atoms or ions has three unpaired electrons?
A) N
B) O
C) Al
D) S2–
E) Ti2+

40. #6. The complete electron configuration of element is
A) 1s22s22p63s23p64s23d104p65s24d105d105px
B) 1s22s22p63s23p64s23d104d104px
C) 1s22s22p63s23p64s24p65s24d105d105px
D) 1s22s22p63s23p64s23d104p65s24d105px
E) none of these

41. # 7. Which group contains species that are isoelectronic with each other?
A) P, S, Cl
B) Ag, Cd, Ar
C) Na, Ca, Ba
D) P, As, Se
E) none of these

42. #8. Which of the following electron configurations is different from that expected?
A) Ca
B) Sc
C) Ti
D) V
E) Cr

43. D) More information is needed to answer this question.
# 9. The first ionization energy of Mg is 735 kJ/mol. The second ionization energy is
A) 735 kJ/mol
B) less than 735 kJ/mol
C) greater than 735 kJ/mol
D) More information is needed to answer this question.
E) None of these.

44. Questions 11-14 (OLD AP QUESTION) (A) Heisenberg uncertainty principle
(B) Pauli exclusion principle
(C) Hund’s rule (principle of maximum multiplicity)
(D) Shielding effect
(E) Wave nature of matter
11. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic.
12. Explains the experimental phenomenon of electron diffraction.
13. Indicates that an atomic orbital can hold no more than two electrons.
14. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron. C E B A

45. Einstein studied the photoelectric effect whereby light of sufficient frequency shining on a metal causes current to flow. The amplitude of the radiation was not important, the frequency was.

46. This told him that the light must be in particles, each having a given energy. Einstein proposed that electromagnetic radiation can be viewed as a stream of particles called photons.
Einstein’s special theory of relativity: E = mc2
Matter and energy are different forms of the same special entity.
Energy of a photon: E = h
λ = h/m
m = mass in kg
v = velocity in m/s
h = Planck’s constant = 6.62x10-34 J.s
Ephoton =hc/λ
 measured in m
c speed of light 3.00x108 m/s
h = Planck’s constant = 6.626x10-34 J.s

48. We can combine this equation with the equation for waves to determine the mass of a photon. This is called the deBroglie equation. DeBroglie’s equation is used to find the wavelength of a particle. It was determined that matter behaves as though it were moving in a wave. This is important in small objects such as electrons but is negligible in larger objects such as baseballs. Since mass is in the denominator or the equation, the larger the mass, the shorter the wavelength. Heavy objects have very short wavelengths.

49. Einstein studied the photoelectric effect whereby light of sufficient frequency shining on a metal causes current to flow. The amplitude of the radiation was not important, the frequency was.

50. Einstein’s photoelectric effect is now being utilized in another manner for chemists. This is called photoelectron spectroscopy, or PES.
PES uses the energy from electrons emitted via the photoelectric effect to gain information about the electronic structure of a substance.

51. The Photoelectric Effect and PES
When light of a certain frequency is shone onto a sample, there are a limited number of electrons emitted. The energy of these emissions reflect the energy (and energy levels) in an atom
When the kinetic energy of the emitted electron is subtracted from the energy of the light photons, the resulting value is the ionization energy of that particular electron.
The photons used for photoelectron spectroscopy range from ultraviolet light to X-rays. Analysis of photoelectron spectra for a single element can give us information about the shells and subshells of electrons and the number of electrons in each.

52. The diagram to the right is a PES image for the element boron.
The x-axis has units of ionization energy or binding energy of the electron.
Electrons closer to the nucleus have greater Coulombic attraction and thus have a higher ionization energy.
The energy scale at the bottom may change and units vary, so all that is needed is interpretation of the diagram, not details.
The height of the peak indicates the number of electrons in the atom having that energy (the number of electrons in a sublevel).

53. Lithium Potassium Let’s Practice! What element is represented here?
What about this one?

54. Krypton Phosphorus Let’s Practice!
Let’s try another one! Note: the 1s peak has been omitted.
Phosphorus
Last one!

55. Let’s look at the PES for gold.
These peaks all occur at energy level 4. What does each peak represent?
How many electrons are shown by each peak?
Discuss the type of energy present at each peak. Why is it high or low?

56. 7.3 The Atomic Spectrum
When EM radiation is separated into its single wavelength components a spectrum is produced. White light produces a continuous spectrum (rainbow) when passed through a prism.

57. Bright Line Spectrum
Some emitters of light radiate only certain colors and wavelengths. This produces a bright line spectrum.
When various gases at low pressures are put in tubes and a high voltage is applied, they glow in various colors.
If this light is passed through a prism, a series of lines of color is produced. This series identifies the element.

58. Bright Line Spectrum of Helium

60. Niels Bohr found that the absorptions and emissions of light by hydrogen atoms correspond to energy changes of electrons within the atom. Bohr proposed that the electron in hydrogen travels only in certain allowed orbits.
Bohr calculated the energy differences between these orbits and predicted the wavelengths at which lines would be found.
7.4 The Bohr Model

61. Bohr determined that when an electron moved from an outer orbital to an inner orbital, or to its ground state (lowest energy level of an electron), it would release a predictable amount of energy.
His research worked great with the hydrogen atom but did not work correctly for polyelectronic atoms.
Polyelectronic atoms- atoms with more than one electron (anything beyond hydrogen)
Bohr developed an equation that could be used to find the change in energy of a hydrogen electron as it goes from one energy level to another:
2.178 x is called the Rydberg constant.

62. E=-2.178 x 10-18J = -4.084 x 10-19J    = 4.864 x 10-7m or 486.4 nm
E = hc= x 10-19= x 10-34(2.998x 108)
 
 = x 10-7m or nm

63. 7.5 The Quantum Mechanical Model of the Atom
THIS UNIT IS NOT ON THE AP EXAM

64. 7.7 Orbital Shapes and Energies
THIS IS NOT ON THE AP EXAM

65. 7.8 Electron Spin and the Pauli Exclusion Principle
In a given atom, no two electrons can have the same set of four quantum numbers. An orbital can only hold two electrons, and they must have opposite spins (creating two magnetic moments).

66. 7.9 Polyelectronic Atoms
When an atom has more than one electron, these electrons tend to repel each other. The effect of the electron repulsions is to reduce the nuclear charge (pull of the nucleus on the electron).

67. The apparent nuclear charge or effective nuclear charge, Zeff, is the charge felt by a particular electron. Electrons in inner shells shield the electrons in higher shells quite effectively from the nuclear charge (shielding effect). Electrons in the same shell are much less effective at shielding each other.
Zeff = Zactual - effect of e- repulsions
Each electron in an atom has its own value of Zeff which can be calculated from the experimental energy required to remove that electron from the atom.
Shielding- the effect by which the other electrons screen or shield a given electron from some of the nuclear charge.
Ex. Zeff for the 3s electron in Na is 1.84, the Zeff for a 1s electron is 10.3.

68. Penetration effect- the effect whereby a valence electron penetrates the core electrons, thus reducing the shielding effect and increasing Zeff.
A 3s electron has a small but significant chance of being close to the nucleus.
Most penetration ns > np > nd > nf least penetration
The Radial Probability Distribution for the 3s, 3p, and 3d Orbitals

69. Electrons fill orbitals in order of increasing energy
Electrons fill orbitals in order of increasing energy. Because of the penetration effect, electrons fill (n+1)s before nd. ((n+1)s has lower energy). For example, if “n” is 4, then an electron would fill 5s before filling 4d.

70. Electrons sharing an orbital do not shield each other as well as core electrons shield outer electrons.
Zeff increases for a 1s electron going form H to He but decreases form He to Li because 1s electrons are effective in shielding the 2s electron.
Zeff increases from Li to Be
Zeff decreases from Be to B
Zeff increases from B to N
Zeff decreases from N to O because one 2p orbital is doubly occupied
Zeff increases from O to Ne

72. 7.10 The History of the Periodic Table
Originally, four elements were suggested by the Greeks; earth, air, fire and water.
Mendeleev, a Russian chemist, was the first to organize elements by their atomic mass and to organize them by their physical characteristics.
Moseley later organized the periodic table by atomic number instead of mass.

73. 7.11 The Aufbau Principle and the Periodic Table
Aufbau principle- As protons are added one by one to the nucleus, electrons enter orbitals of lowest energy first.
In its ground state, atoms have electrons in the lowest energy orbitals. The order that the orbitals fill is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

74. Exceptions to Aufbau principle: Cu and Cr are two elements that have exceptional electron configurations. Cu has a configuration that ends in 4s1 3d10 instead of 4s2 3d9. Cr is 4s1 3d5 instead of 4s2 3d4. This gives half filled sublevels. It is more stable because of decreased electron repulsion.

75. Hund's rule- All orbitals in a sublevel must be half-filled with electrons having parallel spins before any may be completely filled.
Valence electrons- electrons in the outermost principle quantum number
Core electrons- inner electrons

76. TRICKS TO HELP EXPLAIN MOST TRENDS:
In a period/row: use effective nuclear charge (Zeff) and columbic attraction
In a group: use SHIELDING, ADDED ENERGY LEVELS, DISTANCE
Use energy where necessary

77. (this is highly tested on the AP exam)
7.12 PERIODIC TRENDS
(this is highly tested on the AP exam)

78. Trends Tricks:
“A trend is an observation NOT an explanation”. Say it over and over until it sticks in your head. (Stating the trend doesn’t justify an answer, you must EXPLAIN the reason the trend occurs. What you learned in pre-AP is no longer enough information to get full credit for a question)
Talk about BOTH atoms involved on the AP exam or you will lose credit.
Mention Coulombic attraction where it pertains to a trend.
Emphasize energy
Use effective nuclear charge (Zeff) when in a period/row AND use size, distance, and SHIELDING when in a group as an explanation

79. AP WORDS: Zeff (What is it? What is the trend) Coulombic attraction
Shielding
Distance
Energy

80. Remember the Periodic Law
When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

81. Ionization Energy
First Ionization Energy-energy required to remove one electron from a neutral atom
Increases UP and to the RIGHT

82. Ionization Energy Successive Ionization Energies 2nd I.E. 1,445 kJ
Large jump in I.E. occurs when a CORE e- is removed.
Mg 1st I.E kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ

83. Ionization Energy Successive Ionization Energies 2nd I.E. 1,815 kJ
Large jump in I.E. occurs when a CORE e- is removed.
Al 1st I.E kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ

84. Units are usually kJ/mol
Ionization energy (IE)- energy required to remove an electron from a gaseous atom or ion.
X(g)  X+(g) + e-
A measure of how much work it is to remove an electron or how tightly the electron is held by the nucleus.
Units are usually kJ/mol
Ionization Energy depends mainly on two factors:
The effective nuclear charge
The average distance of the electron from the nucleus

85. Ionization Energy Why more energy across a period?
Effective nuclear charge increases the attraction of the nucleus therefore holds electrons more tightly (WITH EXCEPTIONS from GROUP 2 to 13 AND from GROUP 15 to 16)
Why more energy up a group?
Decreased number of energy levels, decreases the distance over which nucleus must pull and therefore reduces attraction for electrons
Full energy levels provide some shielding between nucleus and valence electrons so closer to the top of group, less energy levels and less shielding

86. Practice: Draw a Bohr Model for lithium and Francium
Draw a Bohr Model for Lithium and Neon

87. Ionization Energy~EXCEPTIONS
Exceptions to these general trends occur at group 2 to 3 (shielding of p electron by s electrons)
Less energy needed to remove p because it doesn’t penetrate the nucleus and is not held as tightly
Example: boron
AND from group 5 to 6 (paired electron repulsion).
Increase in repulsion outweighs increase in Zeff and less energy is needed to remove the electron
Example: oxygen

88. Ionization Energy
Ionization energy takes a large jump going from a valence electron to a core electron. The second ionization energy of sodium is much higher than the first. The third ionization energy of magnesium is much higher than that of the first and second electron. Core electrons are much more tightly held than valence electrons.

89. Ionization Energy
The first ionization energy generally increases going from left to right across a period (very similar to Zeff). It decreases going down a group because electrons being removed are farther from the nucleus.
Exceptions to these general trends occur at group 3 (shielding of p electron by s electrons) and group 6 (paired electron repulsion).
Atoms with a low IE1 tend to form cations during reactions, whereas those with a high IE1(except noble gases) tend to form anions.
Metals have low IE. Nonmetals have high IE and very negative EA. Density generally increases going down the periodic table because the atomic number increases faster than the atomic size.

90. Activity: Ionization Energy
1. Draw the orbital notation for an element in group 3 and group 16
2. Highlight the electrons that cause the exception to this trend
3. Explain the exception to the trend from 2 to 3
4. Explain the exception to the trend from 5 to 6
5. Why does rubidium have a smaller I.E. than Na
6. Justify chlorine having a higher I.E. than potassium
7. Which group:
Ist I.E. 352 kJ/mol
2nd I.E. 658 kJ/mol
3rd I.E kJ/mol

91. Electron Affinity
Energy change that occurs when an electron is acquired by a neutral atom (addition of electron in gaseous atom or ion).
Tends to become less negative (less energy released) DOWN and to the LEFT

92. Electron Affinity (EA)
The energy associated with the addition of an electron to a gaseous atom -If the addition is exothermic, EA is negative. The more negative, the greater the amount of energy released.
Electron Affinity (EA)

93. Electron Affinity Explained:
Why down a group:
Change little moving down a group.
Why more negative across a row/period towards noble gases:
Become increasingly negative from left to right. More positive, the less attractive to electrons around them. Careful: addition or subtraction can be exothermic (-) or endothermic (+). As you more toward the noble gases, the affinities become more negative.
Trend explained because of octet rule. (atoms close to full valence will tend to gain electrons and have very negative affinities (give off great deal energy when gaining electrons) NOBLE GASES DO NOT CONFORM TO THIS. They have very positive values.

94. In going down a group, EA generally becomes more positive (less energy released) because the electron is farther from the nucleus. There are many exceptions. EA generally becomes more negative going across a period (again, many exceptions)

95. Atomic Radius
½ the distance between the nuclei of identical atoms that are bonded together
Increases to the LEFT and DOWN

96. Atomic Radius Why bigger across a period/row: Why larger down a group:
Effective nuclear charge decrease attraction of nucleus and therefore the pull of the nucleus results in larger radius
Why larger down a group:
Increase number of energy levels down a group, increase distance over which nucleus must pull electrons and therefore because of shielding of inner core, reduces attraction for electrons

97. Atomic Radius K Na Li Ar Ne

98. Atomic Radius-obtained by measuring the distances between atoms in chemical compounds.
Decreases going from left to right across a period because Zeff increases from left to right.
Increases going down a group because of larger orbitals. Zeff stays about the same because of shielding.

100. Ionic Radius Ionic Radius (WHY) Cations (+) Anions (–) lose e- smaller
gain e-
larger
© 2002 Prentice-Hall, Inc.

101. Ionic Radius Trend
GROUP 1, 2, 3, and D block form cations and in general they will get small because they are LOSING electrons
NOTE: ratio of protons to electrons increases when an electron is lost so electrons are held closer with more strength
GROUP 5, 6, 7 form anions and in general will get bigger because they are GAINING electrons
NOTE: ratio of protons to electrons decrease and the electrons are NOT held closely. Increased electron/electron repulsion also plays a role in expanding the electron cloud

102. ISOELECTRONIC SERIES
An atom or ion containing the same amount of electrons
Ex: Ar, Ca2+, S2-, K+, Cl-
Practice questions:
List all atoms and common ions of representative (main group) elements that are isoelectronic with the nitrogen ion. Which is the smallest?

103. Electronegativity
A measure of the ability of an atom in a chemical compound to attract electrons
Most electronegative element is fluorine
Given arbitrary value of 4; all others relative

104. Electronegativity Explained
Why higher up a group:
Decreased number of energy levels will decrease the distance over which the nucleus must pull and therefore cause an increase for the attraction of the electrons.
Full energy levels AT THE bottom of a group provides shielding for valence electrons and the nucleus
Why higher across a period/row:
Effect nuclear charge increases attraction of the nucleus and therefore strengthens the attraction of the electrons

105. Examples
Which atom has the larger radius?
Be or Ba
Ca or Br Ba Ca

106. Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne N Ne

107. Examples K or Li Al or Cl Li Cl
Which has the greater electonegativity?
K or Li
Al or Cl Li Cl

108. Examples S or S2- Al or Al3+ S2- Al
Which particle has the larger radius?
S or S2-
Al or Al3+
S2- Al

109. Metals have low IE. Nonmetals have high IE and very negative EA
Metals have low IE. Nonmetals have high IE and very negative EA. Density generally increases going down the periodic table because the atomic number increases faster than the atomic size.

110. Transition Metal Size and Lanthanide Contractions The variation in size as we go across a row of transition metals is much less than among the representative elements. This is because electrons are being added to an inner shell as the nuclear charge gets larger.

111. The Lanthanide Contraction
A similar phenomenon occurs among the inner transition metals (Ex. lanthanides) The lanthanides fall between La and Hf, La last fills the 5d1 electron, Ce ends in 5d1 4f1. In this 6th period we have a much larger decrease in size occurring between La and Hf because of the intervening lanthanide elements. This additional decrease in size is known as the lanthanide contraction. It causes Hf to be the same size as Zr, even though Hf is below Zr.
All of the rest of the transition elements in the 6th period are nearly the same size as the elements above them in the fifth period. This causes the 6th period transition metals to be extremely dense. This even influences lead and bismuth beyond the transition metals.