Chemistry 12 Ch 9: Electrons and Periodic Table Page | 1 Chapter 9: Electrons and the Periodic Table 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 examining the Bohr model and the quantum-mechanical model which describe how electrons behave inside the atom and how those electrons affect the chemical and physical properties of elements. These models explain the observed emission/absorption spectra and the periodic behavior of the elements such as why He is inert. Pictured below are Niels Bohr (left) and Erwin Schrödinger (right)
Transcript
C h e m i s t r y 1 2 C h 9 : E l e c t r o n s a n d P e r i o d i c T a b l e P a g e | 1
Chapter 9: Electrons and the Periodic Table
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 examining the Bohr model and
the quantum-mechanical model which describe how electrons behave inside the
atom and how those electrons affect the chemical and physical properties of
elements. These models explain the observed emission/absorption spectra and the
periodic behavior of the elements such as why He is inert.
Pictured below are Niels Bohr (left) and Erwin Schrödinger (right)
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Quantum Mechanics explains the behavior of the electrons inside atoms, the periodic
law and expectations in chemical bonding.
Electrons are extremely tiny.
Electron behavior determines much of the behavior of atoms
Problem: Directly observing electrons is impossible. Observing an electron
would change its behavior
Light has properties of both waves and particles
Wave-like Property of Light
Light is a form of Electromagnetic Radiation
Electromagnetic radiation is a form of energy that travels through empty space at
3.00 x 108 m/s = c = speed of light. (186,000 mi/s)
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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Visible light: The components of
white light R O Y G B I V separate
when passed through a prism
Practice 1:
c = ;
wavelength in meters (,
frequency or hertz in 1/s (),
speed of light (c=3.00 x 108m/s)
a) Solve for the frequency of
green light with a wavelength
of 540nm.
b) A radio signal has a frequency of 100.7MHz. Solve for its wavelength.
Hz = s-1
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Particle-like Property of Light (photons)
Light is particle-like: Light has photons of quantized energy.
Photoelectric Effect: Emission of an electron from a metal surface caused by
shining light (electromagnetic radiation) of certain minimum energy. The current
increases with increasing intensity of radiation. This experiment from Albert
Einstein led to the idea of photons and E = hhc/
(Planck’s constant = h = 6.6262 x 10-34 J s).
The shorter the wavelength
the higher the energy.
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 red - 400nm violet
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Emission Spectra: the
emission of light from
excited gas atoms. This
observation led to Bohr’s
Atomic Model to explain it.
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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Problems with Rutherford’s nuclear atom:
Electrons are moving charged particles. According to the known classical physics
at that time, moving charged particles give off energy, therefore electrons should
constantly be giving off energy; they should glow, 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.
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Hydrogen Series:
nfinal = 1, Lyman series (UV)
nfinal = 2, Balmer series (Visible)
nfinal = 3, Paschen series (IR)
Balmer series which gives visible light is the
only one to remember
Problem: Bohr’s theory
is limited. It only
explained spectra for an
item with one electron,
the element H,
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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): Replaced Bohr model
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,
exceptions are not tested in Introductory chemistry
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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Try this: Predict the ground state short electron configuration for each.
Ca Ca+2
F F-
Ag Ag+1
Fe Fe+2 Fe+3
Periodic Table:
The periodic table gives us information in an organized manner. Patterns and
properties can be 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 smaller than the total charge of the nucleus.
Zeff = Ztotal - Sscreening constant
S is 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 radii generally
increase from
right to left
top to bottom
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.
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Size: Ionic radii and isoelectronic series.
Cations lose electrons and are smaller than the original atom
Anions gain electrons and are larger than the original atom
In an isoelectronic series (all have the same number of electrons, same
electron configuration) the size increases as the charge of the nuclei
