Electrons in Atoms and Periodicity
I. Atomic Models (Review)
A. Rutherford - Gold Foil Experiment
1. Nucleus - hard dense small center of atom containing protons & neutrons
2. Most of the atom is empty space in which the electrons move randomly
1. Did not explain how the electrons filled the empty space around nucleus.
B. Bohr – Solar System Model
1. Electrons orbit the nucleus maintaining a fixed amount of energy, at
constant speeds (goes against the laws of physics)
2. Energy Levels are the circular paths that the electrons travel.
3. Energy levels are arranged like a ladder. (Rungs are not evenly spaced.
They get closer together the further from the ground, nucleus.)
4. Electrons at higher energy levels have greater energy.
5. The amount of energy required by an electron to move from one level to
another is a quantum of energy.
6. A quantum leap is the abrupt change in level by the electron.
C. Schrodinger – Quantum Mechanical Model
1. Also called The Electron Cloud Model.
2. The electrons travel randomly with no definite path.
3. Electrons are particles that act like waves.
4. This model is based on a mathematical 90% probability of finding an
electron within a certain volume of space, called an orbital.
5. Heisenberg Uncertainty Principle – it is impossible to determine both the
position and velocity of an electron simultaneously
II. Wave-Particle Duality of Motion – Electrons act like both waves and particles.
A. Wave Behavior
1. Electromagnetic Radiation - Energy lost or gained as the electron moves
from one energy level to another. (Only some are visible.)
2. Electromagnetic Spectrum - consists of radiation over a broad range of
wavelengths; includes radio waves, radar waves, microwaves, infrared
waves, visible light waves, ultraviolet waves, x-rays, gamma rays, and
cosmic rays. (listed from longest λ to the shortest λ)
3. Electromagnetic radiation has the ability to interfere and diffract.
4. Amplitude - the height of the wave from origin to crest
5. Wavelength - the distance between wave crests. Units are the same as
length (cm). Symbol is λ
6. Frequency - the number of wave cycles to pass a given point per unit of
time. Units of frequency: cycles/second or hertz (SI). Symbol is ν
7. Wavelength and frequency are inversely proportional, when one goes up
the other goes down.
Electrons in Atoms and Periodicity
8. Speed of all forms of electromagnetic radiation (Light) in a vacuum
(empty space) is constant c = 3.0 x 108 m/sec
9. Formula c = λ ν
B. Particle Behavior
1. Photoelectric Effect - electrons are ejected by metals when light shines
on them - light waves act like particles (Einstein). The energy of the light has
to be at a certain minimum. Light can be absorbed or emitted by matter.
Example: Calculators that require light to work.
2. The Quantum Concept - energy increases in small discrete units (or steps)
called quanta. Example: An hourglass. Electromagnetic Energy – E
3. Photon - discrete units of light energy
4. Frequency - energy per photon
5. Intensity - number of photons
6. Plank’s Constant - showed mathematically that the electromagnetic energy
is directly proportional to frequency. h = 6.626 x 10-34 J*sec
7. Formula E = h ν
8. Atomic Emission Spectra - every element emits a unique atomic spectrum,
the pattern of frequencies obtained by passing energy emitted by atoms of an
element in the gas phase through a prism.
9. Ground State - n = 1 is the lowest energy level
1. Higher Energy Levels - when the electron moves from lower energy levels
to higher energy levels it must gain energy. When moving from high to low
energy levels, the electrons release energy in the form of photons = light.
10. Electrons in n = 2, 3, 4 ... are in an excited state
11. Each energy level emits a particular kind of light. Each element emits a
unique combination.
Example: Hydrogen
n = 1 ultraviolet
n = 2 visible
n = 3 infrared
III. Electron Configurations - most probable places to find electrons, locations
A. Levels n=1,2,3,4…to infinity (similar to Bohr's model)
1. Ground State (n=1) lowest energy level, nearest to the nucleus
2. Excited State (n=2,3,4…) higher and higher energy, further and further
from the nucleus
B. Sublevels – reside within the energy levels
1. The number of sublevels corresponds to the energy level number.
Examples: n = 1 has one sublevel, n = 4 has four sublevels
2. The sublevels are represented with letters rather than numbers. s,p,d,f,
then through the alphabet g, h, i, j, k, etc.
Electrons in Atoms and Periodicity
3. All known elements only have enough electrons to fill up to the (f)
sublevel.
C. Orbitals – each sublevel has a certain amount of orbitals
1. The number of orbitals increases by odd numbers.
a) The (s) sublevel has 1 orbital.
b) The (p) sublevel has 3 orbitals.
c) The (d) sublevel has 5 orbitals.
d) The (f) sublevel has 7 orbitals.
2. Each orbital holds a maximum of two electrons each.
a) The (s) sublevel has 1 orbital = 2 electrons.
b) The (p) sublevel has 3 orbitals = 6 electrons.
c) The (d) sublevel has 5 orbitals = 10 electrons.
d) The (f) sublevel has 7 orbitals = 14 electrons.
3. The total number of orbitals on an energy level can be calculated with n2.
The total number of electrons possible in each energy level can be calculated
with 2n2.
D. Atom - Office Building
1. Energy Level - Floors of Building
2. Sublevels - Offices on each Floor
3. Orbitals - the Desks in each Office
4. Electrons - the Two People who share the Desk
E. Aufbau Principle - electrons fill the lowest energy level first.
1. 1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p ....
2. A superscript indicates the number of electrons occupying a sublevel.
3. The sum of all the superscripts will equal the total number of electrons for
that element.
Example: Oxygen - Atomic number is 8
(eight electrons if it is a neutral atom) 1s2 2s2 2p4
4. Valence Electrons are the electrons furthest from the nucleus. They
determine the properties of the element. (how it reacts to external stimuli).
Example: Our skin.
a) Outer most energy level (highest level, furthest away) electrons.
b) Electrons that are lost, gained, or shared between atoms to form
compounds.
c) Matches the Group number for the Representative (A) elements.
5. Exceptions to the Aufbau Principle - Chromium and Copper
F. Pauli’s Exclusion Principle - an atomic orbital can hold a maximum of two
electrons with opposite spins because electrons repel each other. Electrons are
designated with arrows in the Orbital Notation - diagram.
Electrons in Atoms and Periodicity
G. Hund’s Rule - when electrons enter orbitals within the same sublevel, one
electron enters each orbital until all orbitals contain one electron with parallel spins.
The second electron is then added to each orbital so that their spins are paired
oppositely with the first electron. Visualize dealing cards.
H. Procedure for calculating the electron configuration…
1. The atomic number is the proton number which will be the electron
number for the neutral atom.
2. Fill in the diagram using your rules.
3. Write out the sublevels with superscripts representing the number of
electrons.
Example: What is the electron configuration of Tin
I.Noble Gas Notation – To simplify the electron configuration…
1. Put the previous periods noble gas symbol in square brackets
2. Then add only the last periods electron configuration thereafter
Example: Silicon (AN=14) 1s22s22p63s23p2
The closest noble gas prior to Silicon is Neon (AN=10)
so the Noble Gas Configuration would be [Ne]3s23p2
IV. Periodic Table
A. History – Mendeleev
1. Mid-1800’s Russian chemist and high school teacher who developed the
very first periodic table.
2. He studied the atomic masses of the 70 known elements and listed them in
order in columns.
3. He noticed a regular (periodic) recurrence of their chemical and physical
properties, and arranged it so that similar properties were side by side.
4. His arrangement left blanks in his PT where no known element fit. This
enabled him to predict the approximate weight and properties of currently
undiscovered elements.
B. Current - Moseley
1. British physicist who died at 28 in WWII. Rearranged Mendeleev’s PT
into the currently used PT of today.
2. His focus was on nuclear charge, he arranged the elements according to
atomic (proton) number going across the PT.
3. He noticed that the element’s properties lined up in columns.
C. Electron Configurations and Periodicity
1. Periodicity in properties matches the outer electron configuration.
2. Elements in the same group react same way under similar conditions.
3. Block Diagram – identifies groups of elements according to the sublevel
that are filled with electrons.
(1) s block – Group 1A (s1) ands 2A (s2)
Electrons in Atoms and Periodicity
(2) p block – Group 3A (p1) through 8A (p6)
(3) d block – the Representative Elements 3B (d1) - 2B
(d10) period - 1 = level
(4) f block – the Inner Transition Metals (f1 f14)
period - 2 = level
4. Notice the trends in the group, same number of valence electrons in all
elements in a group. Listing the highest sublevel filled:
1A
1s1
2s1
3s1
4s1
5s1
6s1
7s1
3A
8A
2p1
3p1
4p1
5p1
6p1
7p1
2p6
3p6
4p6
5p6
6p6
7p6
Example Group 2A:
Noble Gas Configuration
12 Mg
[Ne]3s2
2 2 6 2
6 2
20 Ca 1s 2s 2p 3s 3p 4s
[Ar]4s2
38 Sr 1s22s22p63s23p64s23d104p65s2
[Kr]5s2
56 Ba 1s22s22p63s23p64s23d104p65s24d105p66s2
[Xe]6s2
Note: row number = period = energy level for the s and p block elements
5. Atomic Radii (size) - half of the distance between 2 nuclei of the same
element. Since atoms are mostly space with no fixed outer boundary, the only
solid part to measure against was the nucleus.
a) Periodic Trend - Going across the PT the radius decreases.
(1) As you go across a period, more electrons are added into
the same energy level.
(2) The more electrons, the more the protons build up in the
nucleus.
(3) The greater the positive charge of the nucleus, the stronger
it will be to pull in on the electrons.
(4) The stronger the nucleus pulls, the closer the electrons get.
b) Group Trends - Going down the PT the radius increases.
(1) As you go down a group, you add energy levels.
(2) Example: Think of layers of an onion. The more layers the
fatter (larger the radius).
6. Ionization Energy - measures the energy required to remove an electron
from the outermost energy level. Corresponds to the size of the atom.
Removing an electron from an atom creates a positive ion = cation.
a) Periodic Trend - Going across the PT the ionization energy
increases as it becomes more difficult to remove electrons.
(1) A small atom holds its electrons very tightly. Strong
attraction of e- to positive nucleus.
Electrons in Atoms and Periodicity
(2) These electrons need a lot of additional energy to escape
their nuclear attraction.
(3) When an atom loses an electron it becomes ionized (ion)
with a positive charge because there are now more protons
than electrons.
b) Group Trends - Going down the PT the ionization energy
decreases.
(1) A big atom has electrons that are far away from the
nucleus. These electrons only need a little more energy to
escape the attraction of the nucleus. 1st ionization energy – the
energy needed for one electron to escape.
(2) After the first electron has escaped, the nucleus is holding
on tighter to the remaining electrons.
(3) 2nd ionization energy must be even greater than 1st
ionization energy for the next electron to escape. 2nd ionization
energy – energy needed for a second electron to escape.
(4) Again, the nucleus holds even tighter to the remaining
electrons.
(5) So 3rd ionization energy must be even greater than 2nd for
another electron to escape. 3rd ionization energy –energy
needed for a third electron to escape.
7. Ionic Radii - the size of a charged “atom”.
a) Ions – charged “atoms” are called ions.
b) Cations – Lost electrons. Same number of protons.
(1) Removing electrons forms positive ions (requires the
addition of ionization energy)
(2) Cations are smaller than their neutral atoms. Less electrons
than protons, so the positive nucleus holds on tighter to the
remaining electrons, pulling them in closer.
(3) Example: Sodium
c) Anions - Gained electrons. Same number of protons
(1) Electron Affinity – Adding electrons to negative ions
(requires the release of energy)
(2) Anions are larger than their neutral atoms. More electrons
than protons, so the electrons gang up and pull further away
from their attraction to the positive nucleus.
(3) Example: Chlorine
d) Periodic Trends – Going down the PT the ionic size increases. The
more energy levels the fatter the ion so the nucleus cannot hold the
electrons as tightly.
Electrons in Atoms and Periodicity
e) Group Trends - Going across the PT the ionic size decreases.
Remaining electrons can be held tighter by the nucleus whose
strength has remained the same.
8. Electronegativity – the tendency to attract electrons from another elements
atom when chemically combining.
a) Corresponds to the ionization energy
b) If hard to lose an electron from self (high ionization energy) then
easy to pull an electron away from another (high electronegativity)
c) If easy to lose an electron from self (low ionization energy) then
hard to pull an electron from another (low electronegativity)
d) Omit the noble gases (8A) and all B Elements when considering
electronegativity.
e) Periodic Trend - Going across the PT the electronegativity
increases.
(1) If the element has a high electronegativity it will pull
electrons from an element with a lower electronegativity.
Think of a tug of war.
(2) Fluorine has the highest electronegativity.
f) Group Trend - Going down the PT the electronegativity decreases.
(1) If the element has a low electronegativity it will NOT be
able to pull electrons away from an element with a higher
electronegativity.
(2) Cesium has the lowest electronegativity.