Atomic Theory and Perioidicity
Views of the Atom





Dalton – One piece of matter, indestructible and indivisible, all atoms of same element are same
JJ Thompson – Used cathode ray tube to find electrons are inside atom (Chocolate-Chip Cookie Model)
Rutherford – Used Gold-Foil Experiment to find massive center has protons and neutrons and the mostly empty
space has the essentially mass-less electrons.
Bohr – Electrically charged Hydrogen gas to find that there are energy levels called Orbits (n=1, n=2).
Electron-Cloud Theory – Used Heisenberg’s Uncertainty Theory to show that electrons were in orbitals.
Atomic Number = number of protons
Atomic Weight = number of protons and neutrons
Ionic Charge = number of protons minus electrons
Isotope = same number of protons (same element), different number of neutrons (a different mass)
Particle
Proton
Neutron
Electron
Mass
1
1
0
Charge
1
0
-1
c = λν (where c = speed of light = 3 x 108 m/sec, λ = wavelength in m, ν = frequency in Hz or sec-1)
E = h ν (where E = Energy in J, h = 6.63 x 10-34 J sec, ν = frequency in Hz or sec-1)
Debroglie wavelength (wavelength of a particle)
h
λ = mv
Electron Configurations
Exceptions: Cu: [Ar] 4s1 3d10 (to be fully-filled)
Cr: [Ar] 4s1 3d5 (to be half-filled)
Atoms will generally lose electrons from the s-orbital before the d-orbital
Orbital Diagrams (Boxes)



Aufbau Principle – Electrons are filled in lowest energy orbitals first
Pauli-Exclusion Principle – Each orbital can have a maximum of two electrons with opposite spins
Hund’s Rule – one electron in each box before pairing (helps to explain ionization energy)
Diamagnetic (electrons are paired), Paramagnetic (electrons are not paired, next to each other)
Unpaired electrons in orbital diagrams make a transition metal attracted to a magnet (example: Fe)




n = Principal Quantum Energy Level or shell
l = Orbital denotation (l = 0 is an s, l = 1 is a p, l = 2 is a d, etc.)
ml = How many boxes are in each orbital (ml = 0 or 1 box; ml = 1, 0, -1 or three boxes; etc.)
ms = Spin of the electron in the box (ml = +1/2 or -1/2)
Example: If n=3; there is 3s, 3p, and 3d available (so l could be 2, 1, 0); for 3s it could have 1 box (ml=0); for 3p it could
have 3 boxes (ml=1,0,-1); for 3d it could have 5 boxes (ml=2,1,0,-1,-2); each box could have ms= +1/2 or -1/2.
Diamagnetic – all electrons are spin paired. Not affected by a magnetic field. (Example: Zinc)
Paramagnetic – having electrons that are not spin paired. Strongly affected by a magnetic field. (Example: Iron)
Periodic Trends
All periodic trends can be explained using these three basic principles:
1. Electrons are attracted to the protons in the nucleus. This is called the Effective Nuclear Charge (Zeff)
a. The closer the electron is to the nucleus, the more strongly an electron is attracted.
b. The more protons in the nucleus, the more strongly an electron is attracted.
2. Electrons are repelled by other electrons.
a. Valence electrons are shielded by completed shells of electrons.
3. Completed shells are very stable. Atoms will lose or gain valence electrons to create completed shells if possible.

Atomic Radius Decreases
o Up due to the n = # becoming smaller (less shells). There is less shielding, therefore more attraction
between the valence electrons and the nucleus.
o To the Right due to the Effective Nuclear Charge (Zeff) increasing. A greater number of protons creates
more attraction of the valence electrons to the nucleus.

Ionic Radius
Cations (+ ions) are smallest. Anions (- ions) are largest. Neutral atoms are in the middle.
o Cations have more attraction between nucleus and valence electrons because the electron-electron
repulsions are reduced (lose e- gets smaller)
o Anions has greater electron-electron repulsion (gain e- gets bigger)

Ionization Energy Increases (Energy required to remove an electron)
o Up due to the n= # becoming smaller. Electrons are closer to the nucleus due to less shielding therefore
will be held more tightly requiring more energy to be removed.
o To the Right due to more protons being added which creates a greater attraction between the valence
electrons and the nucleus. Greater attraction requires more energy to remove an electron.
o The second ionization energy is always greater than the first because when an electron has been removed,
the electron-electron repulsion decreases and the attractive force between the valence electron and the
nucleus increases requiring more energy to remove a second electron.
o Chemical Equation: Na + energy  Na+ + 1eNote: B has smaller 1st Ionization Energy than Be since it has 1 e- in the 2p and Be has filled s.
Note: O has a smaller 1st IE than N since O is 1 e- away from half-filled and N is already half-filled.
Note: Low 1st IE, low 2nd IE, high 3rd IE – will only lose 2 e-, therefore Group IIA (Mg, Ca, etc.)

Electronegativity Increases (Ability to attract an electron in a bond – 0 EN for Noble Gases)
Same reasons as Ionization Energy and Electron Affinity. Greatest EN is Fluorine.

Electron Affinity – When an atom (usually nonmetal) gains an electron, it gives off energy (exothermic) when it
completes a shell and becomes more stable. Adding an electron to a metal will make the element unstable,
therefore one would have to put in energy (endothermic) to gain an electron.
o Chemical Equation: F + 1e-  F- + energy

Reactivity
Alkali Metals : Greater Reactivity going down Group IA due to lower Ionization Energy
Halogens: Greater Reactivity going up Group VIIA due to greater Electron Affinity
Atomic Theory and Perioidicity Problem Set
Atomic Theory
1. Complete the table below, assuming each column represents a neutral atom:
52
Cr
Symbol
25
Protons
30
64
Neutrons
48
Electrons
Atomic Mass
2. Complete the table below:
59
Co3+
Symbol
Protons
Neutrons
Electrons
Oxidation Number
34
46
36
82
86
222
76
116
207
80
120
78
2+
3. There are only two known isotopes of Chlorine, chlorine-35 and chlorine-37. If the atomic mass is 35.5, what is
the percentage of the isotope chlorine-35.
(A) 25%
(B) 35%
(C) 50%
(D) 65%
(E) 75%
4. A certain element has two stable isotopes. The mass of one of the isotopes is 62.93 amu and the mass of the other
isotope is 64.93 amu.
a. Identify the element. Justify your answer.
b. Which isotope is more abundant? Justify your answer.
5. Determine how many valence electrons and unpaired electrons for:
a. Phosphorous
b. Oxygen
c. Sodium
d. Fluorine
6. Write the abbreviated electron configuration for the following:
a. Chlorine
b. K
c. Titanium
d. Cu
e. Zn2+
f. Ag+
7. Identify the group or element that corresponds to each of the following electron configurations:
a. [noble gas] ns2np5
b. [noble gas] ns1
c. [noble gas] ns1 (n-1)d9
d. 1s22s22p63s23p4
e. 1s22s23p24p1 (in addition to finding this element, what state is this atom?)
8. If sodium metal is irradiated with light of 450 nm,
a. What is the frequency of this light?
b. What is the energy of this light?
9. Light used to break down hydrogen peroxide is at 100 nm of wavelength.
a. What is the frequency of this light?
b. What is the energy of this light?
c. Why do you not notice this wavelength of light?
10. Explain why solid sodium conducts electricity, but solid sodium chloride does not.
11. What is the maximum number of electrons that can be held in a shell with a principal quantum number of n=3?
(A) 2
(B) 8
(C) 32
(D) 24
(E) 18
12. Which of the following elements is diamagnetic?
(A) H
(B) Li
(C) Be
(D) B
(E) C
13. Which of the following is an impossible set of quantum numbers (n, l, ml, ms)?
(A) 4, 0, 0 ½
(B) 4, 0, 1, ½
(C) 4, 1, 0, ½
(D) 4, 1, 1, ½
(E) 4, 2, 1, ½
Questions 14-17 refer to the orbital diagrams below.
14.
15.
16.
17.
Represents an atom with six valence electrons.
Is a violation of Hund’s Rule.
Represents a noble gas.
Is in an excited state.
Periodicity
Questions 18-20 refer to the following terms.
(A) Ionization Energy
(B) Electron Affinity
(C) Electronegativity
(D) Atomic Radius
(E) Ionic Radius
18. The change in energy when an electron is added to an element in the gas phase to form a gaseous anion.
19. The amount of energy that is required to remove an electron from an element in the gas phase to form a gaseous
cation.
20. Decreases consistently, from element to element, while moving from left to right across the periodic table.
Questions 21-23 refer to the atoms of the following elements.
(A) Gallium
(B) Potassium
(C) Fluorine
(D) Neon
(E) Selenium
21. Is the most electronegative element.
22. Its highest energy electron in the ground state has the quantum numbers: n=4, l=1, m=-1, and ml=1/2.
23. Has two unpaired electrons.
24. What mass of hydrogen gas is displaced when 6.0 mol of iron react with excess hydrochloric acid according to the
process outlined below.
2 Fe + 6 HCl  2 FeCl3 + 3 H2
(A) 9.0 g
(B) 6.0 g
(C) 12.0 g
(D) 18.0 g
(E) 36.0 g
IE1
738 kJ/mol
IE2
1451 kJ/mol
IE3
7733 kJ/mol
IE4
10540 kJ/mol
IE5
13630 kJ/mol
IE6
17995 kJ/mol
25. The successive ionization energies for a certain element are outlined in the table above. Predict the identity of the
element from the information given.
(A) Na
(B) Mg
(C) Al
(D) Si
(E) P
26. Which of the following ions has the smallest ionic radius?
(A) O2(B) F(C) Na+
(D) Mg2+
(E) Al3+
27. As one moves from element to element, from left to right, across a period in the periodic table, the atomic radius
generally
(A) Decreases
(B) Increases
(C) Remain Constant
(D) Increase and then Decrease
(E) Decrease and then Increase
28. As one moves from element to element, from left to right, across a period in the periodic table, the
electronegativity generally
(F) Decreases
(G) Increases
(H) Remain Constant
(I) Increase and then Decrease
(J) Decrease and then Increase
29. A compound is found to contain 92.6% Hg and 7.4% O by mass. What is the empirical formula for the
compound?
(A) Hg2O
(B) Hg3O2
(C) HgO
(D) HgO2
(E) Hg3O
30. Which element has the largest third ionization energy?
(A) Mg
(B) Al
(C) Ga
(D) In
(E) Tl
31. Which of the following groups of species are considered to be isoelectronic?
(A) Cl-, I-, F(B) Kr, Xe, Rn
(C) Cl-, Ar, Na+
(D) I-, Xe, Ba2+
(E) K+, Ca2+, Fe2+
32. A sample of KNO3 (FW = 101) is mixed with a small amount of some impurity. It was found that the overall
sample consisted of 20% potassium by mass. What was the approximate percentage of KNO3 in the sample?
(A) 25%
(B) 39%
(C) 52%
(D) 63%
(E) 75%
Explain each of the following in terms of principles of atomic structure, including the number, properties and
arrangements of subatomic particles.
33. Rank the following atoms in order of increasing effective nuclear charge experienced by the valence electrons:
K, Mg, H, I
34. Write the chemical equation that shows the processes that describe the first, second, and third ionization energies
of Magnesium. Use relative values for energy (example: lowest energy, highest energy, etc.)
35. Why are ionization energies always positive values?
36. When are electron affinity values positive and when are they negative?
37. Why is the second ionization energy always larger than the first?
38. Lithium has a larger first ionization energy than Sodium.
39. The difference between the first and second ionization energies of potassium is much larger than the difference
between the first and second ionization energies of calcium.
40. The first ionization energy of Mg is 738 kilojoules per mole and that of Al is 578 kilojoules per mole. Account for
this difference.
41. Lithium has a larger second ionization energy than Beryllium.
42. The radius of the Ca atom is 0.197 nanometer; the radius of the Ca2+ ion is 0.099 nanometer. Account for this
difference.
43. The electron affinity value for bromine is negative, whereas it is positive for kypton.
44. Cesium is much more reactive in water than lithium.
45. Consider the two chemical species S and S2a. Write the electron configuration (e.g., 1s2 2s2 . . .) of each species.
b. Explain why the radius of the S2- ion is larger than the radius of the S atom.
c. Which of the two species would be attracted into a magnetic field? Explain.
46. As shown in the table below, the first ionization energies of Si, P, and Cl show a trend.
Element First Ionization Energy
(kJ mol-1)
Si
786
P
1012
Cl
1251
a. For each of the three elements, identify the quantum level (e.g., n=1, n=2, etc.) of the valence electrons in
the atom.
b. Explain the reasons for the trend in first ionization energies.
47. Which has a larger radius and why, O or O2-?
48. Atomic radius decreases from Na to Cl across the period.
49. The difference between the atomic radii of Na and K is relatively large compared to the difference between the
atomic radii of Rb and Cs.