Chemistry 12
Ch 9:
Electrons and Periodic Table
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Chapter 9: Electrons and the Periodic Table
Bonus: 29, 31, 41, 45, 49, 51, 53, 61, 67, 75, 79, 83, 87, 89
Work on MasteringChemistry assignments
What we have learned:
Dalton’s Indivisible Atom explained the Law of Constant Composition and the Law
of Conservation of Mass and led to the Law of Multiple Proportions.
J.J. Thomson, through Cathode Ray Tube experiments, discovered electrons are
small negatively charged particles inside a divisible atom and came up with the
Plum Pudding Model of the atom.
Rutherford came up with the gold foil experiment shooting alpha particles through
thin gold foil to test the Plum Pudding Model and discovered that some alpha
particles were deflected. This led to Rutherford’s Nuclear Model of the atom in
which a heavy positive nucleus is surrounded by a cloud of electrons.
Now we will further our knowledge of the atom by describing how electrons
behave inside the atom. The Bohr model and the Quantum-Mechanical model of
the atom explains emission/absorption spectra and the periodic behavior of the
elements such as why He is inert.
Constants and Equations:
Below are several constants and equations.
Mass of an electron (me = 9.11 x 10-31 kg)
c = ;
wavelength (), frequency or hertz (), speed of light (c=3.00 x 108m/s)
sidenote: sound waves move about 340 m/s
E = h;
Planck’s constant (h=6.6262 x 10-34 J s)
Energy of a Bohr Hydrogen orbit (En = -B/n2)
Bohr’s constant (B = 2.179 x 10-18 J)
[Maximum number of electrons allowed in a quantum number level n] = 2n2
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Quantum Mechanics explains the behavior of the electrons inside atoms and leads to the
explanation of the periodic law and chemical bonding.
Electrons are very small.
Electron behavior determines much of the behavior of atoms
Directly observing electrons is impossible; observing an electron would
change its behavior
Nature of Light:
Light is wave-like (c = ).
Electromagnetic radiation has a magnetic field component and perpendicular to
that an electric field component.
c = ;
speed of light= (wavelength)(frequency)
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Ch 9:
Electrons and Periodic Table
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The electromagnetic spectrum is continuous starting with the low energy, long
wavelength, low frequency radio waves and increasing in energy through
microwaves, IR, VIS (ROYGBIV), UV, Xrays, and gamma rays which are high
energy, short wavelength, high frequency).
Visible light ranges around 750nm-400nm
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Try this 1:
c = ;
wavelength (), frequency or hertz (), speed of light (c=3.00 x 108m/s)
E = h;
Planck’s constant (h=6.6262 x 10-34 J s)
a) Solve for the frequency of green light with a wavelength of 540nm.
b) Solve for the wavelength of a radio signal whose frequency is 100.7MHz.
Hz = s-1
c) Solve for the energy of a photon with a frequency of 2.55 x 1014s-1.
Light is particle-like (E = h). Light has photons of quantized energy. Quantized
energy can explain the emission of light from hot bodies, the emission of electrons
from metal surfaces on which light shines (the photoelectric effect), and the
emission of light from excited gas atoms (emission spectra)
Electromagnetic radiation is a form of energy that travels through empty space at 3.00 x
108 m/s (186,000 mi/s) and has wavelike and particle like properties.
Photoelectric Effect: Emission of an electron from the surface of a metal caused by
impinging electromagnetic radiation of certain minimum energy; current increases
with increasing intensity of
radiation. Experiment led
to photons and E = h
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Emission Spectra: the
emission of light from
excited gas atoms
Mercury (blue)
Hydrogen (pink)
Emission Spectra are like
fingerprints, each element or
compound has a unique emission
spectrum. This allows scientist to
investigate what matter makes up
the stars without going to the sun
and bringing back a sample to test.
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Electrons and Periodic Table
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Problems with Rutherford’s nuclear atom:
• Electrons are moving charged particles
• According to classical physics, moving charged particles give off energy
• Therefore electrons should constantly be giving off energy; they should glow!
• The electrons should lose energy, crash into the nucleus, and the atom should
collapse; but it doesn’t!
Neils Bohr (1885–1962)
1913 Bohr’s Model (electrons move around the nucleus in circular orbits):
Emission spectra of hydrogen gave experimental evidence of quantized energy
states for electrons within an atom.
Quantum Theory:
Explains the emission and absorption spectra
1. An atom has discrete energy levels (orbits) where e- may exist without
emitting or absorbing electromagnetic radiation.
2. An electron may move from one orbit to another. By doing so the
electromagnetic radiation is absorbed or emitted.
3. An electron moves in circular orbits about the nucleus and the energy
of the electron is quantized.
Chemistry 12
Ch 9:
Electrons and Periodic Table
Hydrogen Series:
nfinal = 1, Lyman series (UV),
nfinal = 2, Balmer series
(Visible)
nfinal = 3, Paschen series (IR)
nfinal = 4, Brackett series (IR)
nfinal = 5, Pfund series
Balmer series which gives visible light is the
only one to remember
Problem: Bohr’s theory is
limited. It only explained
spectra from anything with
only one electron.
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Bohr’s atomic theory only worked for 1 electron systems, to explain further the next
theory involves orbitals not orbits…
Quantum Mechanical Model of the Atom (orbitals):
Electrons can be treated as waves or particles (just as in light)
Weakness: Heisenberg’s Uncertainty Principle. It is impossible to determine both
the momentum and position of an electron simultaneously. This means that the
more accurately you know the position of a small particle, such as an electron, the
less you know about its speed (momentum) and vice-versa
Quantum Mechanical Model:
Use 90% probability maps (orbitals not orbits)
volume of space.
1. Electrons have quantized energy states
(orbitals).
2. Electrons absorb or emit
electromagnetic radiation when
changing energy states.
3. Allowed energy states are described by
four quantum numbers which describe
size, shape, position, and spin
respectively.
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Orbitals:
The orbitals in an atom are all centered around the
same center nucleus. As orbitals are filled, some seem to
overlap, creating an electron cloud type appearance.
Orbitals are quantized and exist at discrete energy levels.
Shapes: within an energy level
s-1, p-3, d-5, f-7
Amounts:
Maximum 2 electrons in any one orbital,
Maximum 2n2 electrons for any n level
n= 1; 2 electrons in 1s2
n = 2; 8 electrons 2s22p6
n = 3; 18 e-‘s 3s23p63d10
Electron configuration:
Aufbau Principle: method that fills
electrons in the ground state
Ground state electrons fill lowest
energy and up
Excited state occurs when electrons
have jumped to higher states.
Hund’s Rule of Multiplicity: Electrons
fill up singly with the same spin
within an energy sublevel before
doubling up
Pauli’s Exculsion Principle: No two electrons in the same atom may have the
same four quantum numbers (they cannot occupy the same orbital with
the same spin)
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Electrons:
core [noble gas element]
pseudocore
valence electrons
Isoelectronic series: Atom and ions having the same number of electrons
Electron Configurations:
Long form:
Condensed form:
Condensed configurations
start with a noble gas in
brackets.
N [He]2s22p3
Energy vs Size when d, or
f orbitals are involved.
Orbital diagram: labels and a drawing of the electronic configuration in order to show
each orbital as filled with two electrons, half filled with one electron or empty for
each energy sublevel.
Paramagnetic: (weakly
attracted to a
magnetic field) The
electron configuration
has unpaired
electrons.
Diamagnetic: (weakly
repelled by a
magnetic field) All
electrons are paired.
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Electron Configuration of atoms and ions
Understand the general trends.
Some exceptions (Cr, Mo s1d5, Cu, Ag, Au s1d10), Nb is 5s14d4, Ce, Th,
Cations lose valence p,s orbital electrons before d orbital electrons.
Sn
Sn+2
Sn+4
Ions and their Electron Configurations:
Metals: Give up electrons and become positive cations
Nonmetals: Accept electrons and become negative anions
Main group atoms give up or accept electrons toward the goal of establishing a
core (noble gas) electron configuration. Metals will first lose the valence p
electrons before the valence s electrons.
Transition metals and Inner Transition metals first lose the largest s electrons,
then the d electrons and for inner transitions last are the f electrons.
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Quantum Numbers
Each electron in an atom is defined by four Quantum Numbers.
Name
Symbol
Values
Significance
principle
n
1, 2, 3, ...
size
azimuthal

0, 1, ...(n-1)
shape
magnetic
ml
0, 1, 2,.... 
orientation
spin
ms
½
electron spin
Size: n
1, 2, 3, 4, 5, 6, 7
s, p orbital n = period number
d orbitals n-1 of period
f orbitals n-2 of period
Shape: l
0, 1, 2, 3
s = 0, p = 1, d = 2, f = 3
Orientation: ml
-l…+l
s=0,
p=-1, 0, +1,
d=-2,-1, 0, +1, +2,
f=-3,-2,-1,0,+1,+2,+3
Spin: ms -1/2, +1/2
Try this:
Predict the ground state short electron configuration for each.
Ba
Ba+2
Ag
Ag+1
Fe
Fe+2
Fe+3
P
P-3
P+3
I
I-1
P+5
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Periodic Table:
The periodic table gives us information in an organized manner. If one knows the
properties of some elements in a group, one can predict the properties of
another element in the same group. Many patterns can be seen or predicted
following groups and periods.
Effective Nuclear Charge:
Negatively charged electrons are attracted to the positively charged nucleus.
The attraction depends on the net nuclear charge acting on an electron and the
average distance between the nucleus and the electron.
Zeff (effective nuclear charge), is always smaller than the total charge of the
nucleus. Zeff = Ztotal - Sscreening constant
S is generally close to the # of core electrons, (i.e. for Na, 10 core electrons, 1
valence electron. The Zeff = 11-10 = +1 for the valence electron 3s1)
Periodic Trends:
Size (atomic and ionic radii):
Atomic radii increase from right to left; top to bottom of periodic table.
Along a period the Zeff increases left to right pulling in electrons closer to the
nucleus and causing the atoms to decrease in size.
Vertically, the size of the orbitals (quantum number n) increases from top to
bottom.
Ionic radii and isoelectronic series.
Cations lose electrons and are therefore smaller than the original atom
Anions gain electrons and are larger than the original atom
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In an isoelectronic series (all have the same number of electrons, same
electron configuration) the size increases as the charge of the nuclei
decreases. (smallest Sr+2, Rb+1, Kr, Br-1, Se-2 largest)
Ionization Energy (Energy required to remove the outermost ground state electron,
endothermic):
Ionization Energy generally decreases from right to left; top to bottom of
periodic table.
The small nonmetals require the highest ionization energy; they do not want to
lose electrons
Large metals have lowest ionization energy, they want to lose electrons and
become positively charged cations.
Oxidation is a term for an atom losing electrons and increasing its charge (its
charge is also known as its oxidation state).
Electron Affinity (Energy given away when adding an electron, exothermic):
The greatest negative value (most preferred) electron affinity is for fluorine,
(small nonmetals, ignore noble gases they do not want to add electrons)
decreases from right to left; top to bottom of periodic table.
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Metallic Character:
Metallic character increases from right to left; top to bottom of periodic table.
Nonmetal character increases from left to right bottom to top.
Electronegativity (Ch 10 page 341):
The ability of an element to attract electrons within a covalent bond.
Electronegativity decreases from right to left; top to bottom of periodic table.
It does not include the noble gases. The strongest one is F.