Download File

Class Notes: Unit 2 – Atomic Structure and the Periodic Table
(Ch. 4, 5, and 6)
Ch. 4 The Atom and Subatomic Particles
Substance - Matter with a single composition and a single set of properties
Elements – A pure substance that is the simplest form of matter with a unique set of
properties
Atoms – The smallest piece of an element that could no longer be divided without losing
the properties of the element
The atom is the smallest particle of an element that retains the properties of the element.
The name atom comes from the Greek word “atomos” meaning indivisible. The atom is
very small. If you have a copper penny and you keep dividing it in half until you can’t
divide it anymore and have it still be copper, that smallest piece is an atom. A copper
penny has 2.4x1024 atoms. By comparison, the earth’s population is 7x109 people.
Dalton’s Atomic Theory:
-All elements are composed of tiny indivisible particles called atoms
-Atoms of the same element are identical
-Atoms of one element are different from atoms of every other element
-Atoms of different elements can combine to form compounds
-Chemical reactions occur when atoms are joined, separated, or rearranged
-Atoms of one element are never changed into atoms of another element
Dalton’s atomic theory is still accepted today, except for one important change – atoms
are divisible into smaller particles. Atoms are made up of three types of particles:
protons, neutrons, and electrons. These three are collectively known as the “subatomic
particles” of an atom. The properties of the subatomic particles are all the same,
independent of what atom their in. For example, an oxygen atom is very different from a
lead atom; however, the electrons in both the oxygen and lead atoms are exactly the same
(have exactly the same properties).
Protons (p+)
- positively charged particles (charge = +1)
- 1 amu (amu = atomic mass unit) - same mass as a neutron
- identify element
- never change for a given element
Neutrons (n0)
- neutrally charged particles (charge = 0)
- 1 amu - same mass as a proton
- affect mass of element
- all elements have versions of their atoms with different number of neutrons
(called isotopes)
Electrons (e-)
- negatively charged particles (charge = -1)
~ 0 amu - much much smaller in mass than protons and nuetrons
- affect charge and properties (other than mass) of element
- can be gained or lost to form a charged particle (aka an ion)
The actual mass of a proton or neutron is on the order of 10-24 g and an electron is about
10-28 g. Relative masses are used to get simpler, more manageable (practical) numbers
like 1 for protons and neutrons and 0 for electrons.
The atom can be thought of as a sphere. In the middle is a cluster of protons and neutrons
(the nucleus). The neutrons act like a glue. If there were only protons in the nucleus, their
charges would repel each other. The nucleus is positively charged since it is made up of
only protons (+ charge) and neutrons (no charge). The negatively charged electrons are
orbiting around the nucleus. Individual atoms of an element are neutral, meaning they
have a balanced charge, or the same number of protons (+1 charge) and electrons (-1
charge). In a neutral atom, the number of protons in the nucleus is equal to the number of
electrons in orbit (+ charges balance the - charges so overall the atom is neutral). The
electrons are moving very fast. The exact locations of electrons can’t be pinpointed – you
can only determine the probability of the region where they will be.
The protons and electrons are responsible for the chemical behavior of the atom. The
protons identify the atom as a specific element. The electrons determine the chemical
properties.
The periodic table holds a lot of useful information, if you know how to read it. Each box
contains an element symbol, the name of the element, and two numbers. The number
above the symbol is the atomic number. The number below the symbol is the atomic
mass. The atomic number is always the number of protons in the atom. The protons are
what identify the atom as a specific element. For example, if the atom has 6 protons, the
atom will always be carbon. If it has 1 proton, it will be hydrogen.
Ions
Everything wants to be in a state of equilibrium- meaning that everything tries to be as
balances as possible; as stable as possible; at the lowest energy as possible. Atoms in
their individual state are neutral, meaning they have a balanced charge, or the same
number of protons (+1 charge) and electrons (-1 charge). The number of protons can
never change or else the element the atom is will change. However, it is possible when
one atom chemically bonds with another atom for the number of electrons to change.
When the number of electrons is different from the neutral balanced state, the atom is
known as an ion.
When the protons and electrons are not equal, the atom is charged. The charge describes
the difference in the number of protons and electrons. For example, an atom with a +2
charge has 2 more protons than electrons (because protons are the subatomic particle with
a positive charge). An atom with a -2 charge has 2 more electrons than protons (since
electrons are the particle with a negative charge). **The number of protons is fixed for an
element and is always equal to the atomic number, so the number of electrons is the
number that changes in a charged atom.
For example:
Sodium (Mg)
Atomic number = 12
For a neutral atom of Mg: protons = 12, electrons = 12
For Mg+2: protons = 12, electrons = 10
Chlorine (Cl)
Atomic number = 17
For a neutral atom of Cl: protons = 17, electrons = 17
For Cl-1: protons = 17, electrons = 18
Isotopes
In the boxes on the periodic table, the number below the symbol is the atomic mass (also
known as the weighted average mass number). Every element naturally has atoms with
different numbers of neutrons. These are called isotopes. An isotope is one of two or
more atoms of an element that have the same number of protons but a different number of
neutrons; they have the same atomic number, but different mass numbers. The atomic
masses on the periodic table are not whole numbers. Why not? They are weighted
averages that take into account all the different forms of the atoms of an element. In
nature, the elements occur as a mixture of isotopes.
The mass number of an atom is the number of protons and neutrons the atom has. Each
individual isotope has its own whole number mass number. It is always a whole number
because for an individual atom you cannot have a fraction of a proton or neutron.
Because the mass number is the number of protons and neutron, it is also the mass of the
nucleus (the only place in the atom where protons and neutrons are located).
Other than the different number of neutrons, isotopes are chemically identical. Why?
Because it is the number of protons that determines the identity of the atom and the
number of electrons that determines the properties of the atoms. Isotopes only vary in the
number of neutrons which affect nothing but the mass.
Isotopes are shown in shorthand by stating the symbol, then a dash, then the mass number
of the isotope. The most common/most abundant isotope is usually represented just by
the element symbol.
Ex:
C has 6p+, 6n0, 6e- (could also be referred to as C-12)
C-13 has 6p+, 7n0, 6eC-11 has 6p+, 5n0, 6eEach individual isotope has a fixed mass number (number of protons and neutrons) and a
natural percent abundance. Abundance describes how common something is. The percent
abundance in this case describes how common each isotope is in nature. One isotope (one
atomic form of the element) is usually predominantly abundant over the others.
Atomic Mass = (Mass 1)(%1) + (Mass 2)(%2) + ….
The percent abundances when plugged into this equation should be put in decimal form
(so 90% becomes 0.90).
Ex:
Potassium has two naturally occurring isotopes. K-39 has a percent abundance of 93.5%
and K-41 has a percent abundance of 6.5%. What is the atomic mass of potassium?
Atomic Mass = (Mass 1)(%1) + (Mass 2)(%2) + ….
Atomic Mass = (39)(.935) + (41)(.065)
Atomic Mass = 36.465 + 2.665
Atomic Mass = 39.13 amu
What is the mass of something usually measured in? Grams. Is the average mass of atoms
measured in grams? No. Remember how the real masses of electrons, protons, and
neutrons were really small and so they were turned into relative masses of 0, 1, and 1.
The mass of atoms is really small too, so it is easier (more practical) to talk about them in
relative terms. All the masses on the periodic table are relative to the mass of the carbon12 atom. The masses are all measured in “amu” or “atomic mass units”. Carbon-12 was
picked to have a mass of exactly 12.00000 amu.
1 amu = 1/12 the mass of carbon-12.
The atomic of He is 4.003 amu. This average is a weighted average meaning it takes into
account both the mass of the isotopes as well as their percent abundances. (He-3 occurs
0.0001% of the time and He-4 occurs 99.999% of the time). In the case of He, which
isotope has the highest percent abundance? He-4. How could you tell, even without being
given the exact percent abundances? Out of the 2 isotopes of He (He-3 and He-4), He-4 is
the isotope with a mass number closest to the atomic mass, meaning it is the isotope with
the highest percent abundance. It is the isotope with the biggest contribution to the
weighted average. You can usually round the atomic mass to the nearest whole
number to tell which mass number has the highest percent abundance.
Ex:
Carbon has two major contributing isotopes: C-12 with a percent abundance of 98.89%
and C-13 with a percent abundance of 1.11%.
Regular average mass for C:
(12+13.003)/2 = 12.502
(not correct process - doesn’t take into account %s)
Weighted average mass for C:
(12)(.9889)+(13.003)(.0111) = 12.011 amu
(correct process – because the average is weighted to show which mass
occurs more frequently)
You can also use the atomic mass equation to determine the percent abundances of your
isotopes. Since you do not know the percent abundances, they can be expressed in your
equation as variables (like x and y). You are now in a situation where you have one
equation to solve two variable – a problem. The thing that helps you in this situation is
that you know the percent abundances have to add to 100%. So you actually have a
second equation ( x + y = 100% or x + y = 1 since 100% in decimal form is 1). You
can rearrange this equation so that you are solving for one of the variables (so it turns in
to x = 1 – y or y = 1 – x ) and then plug that back into your atomic mass equation.
This makes your atomic mass equation only involve one variable instead of two.
Ex:
Potassium has two naturally occurring isotopes. K-39 and K-41. The atomic mass of
potassium is 39.13 amu. What are the percent abundances of the isotopes of potassium?
Atomic Mass = (Mass 1)(%1) + (Mass 2)(%2) + ….
39.13 = (39)(x) + (41)(y)
39.13 = (39)(x) + (41)(1-x)
39.13 = 39x + 41 – 41x
-1.87 = -2x
0.935 = x
You now have to interpret what this value of x is telling you. It means that whichever
isotope you had the original percent represented by x, this decimal (returned to
percentage form) is its percent. The other percentage is 1-this decimal or 100%-the x
percentage.
Therefore, K-39 = 93.5% and K-41 = 6.5%
Atomic Number
- The number located above the element on the periodic table
- Always equal to the number of protons in the atom (determines which atom it is)
Atomic Mass
- The number located below the element on the periodic table
- The weighted average mass of all the isotope masses (the weighted average of all the
isotope mass numbers taking into account each isotope’s percent abundance)
Mass Number
- The mass of the nucleus of one atom (number of protons + neutrons)
- If just given the element symbol (for example: C) you are referring to the most
abundant/most common isotope of that element. The mass number then is the atomic
mass (the number below the element symbol on the periodic table) rounded to the
nearest whole number. There are 2 commonly used elements that are exceptions to
this rule: Cl and Cu (because their mass numbers are rounded to .5 it is more difficult
to tell which whole number is actually the most abundant mass number).
- If you are given the element symbol in a shorthanded form (for example: C-14 or 146C )
the number after the symbol in the first case or the larger of the two in the second
case is the mass number - even if this number is not close to the atomic mass on the
periodic table
Which isotope is it?
- All elements have more than one isotope so whenever you are referring to an atom it
is by default also an isotope. What you are usually trying to determine is whether it is
the more common/more abundant isotope or not.
- Isotopes vary based on the number of neutrons in the atom. Therefore, the number that
you need to focus on to determine what isotope is the mass number (mass number
includes the number of protons and the number of neutrons).
- If you are just given a symbol (Ex: C) or given the symbol with a whole number very
close to the atomic mass (Ex: C-12) that atom is the most abundant isotope.
- If you are given a symbol with a whole number not very close to the average mass
number (Ex: C-13 or 136C ) then you are dealing with one of the lesser abundant
isotopes.
Ch. 6 The Periodic Table
The periodic table is a very organized arrangement of elements. It is arranged in order of
increasing atomic number (number of protons).
Horizontal rows of elements are called “periods” (7 total periods). The element properties
vary (either increasing or decreasing) as you move across a row and this variation pattern
repeats as you move to a different row.
Vertical columns of elements are called “groups” (18 total groups). All elements in a
group have similar chemical and physical properties.
Mendeleev published the first widely accepted periodic table in 1869. He arranged the
elements in order of increasing atomic mass (different from how we do it today). He left
spaces in his table for elements that hadn’t yet been discovered but he knew fit into the
periodic patterns he had developed. He was able to predict the properties for the
undiscovered elements based on their future locations on the periodic table.
Some problems arose, however, with organizing by atomic mass. For example,
127.6
52
Te
and 12653.9 I . Iodine has properties very similar to bromine (Br) and chlorine (Cl) and
therefore should be placed in a group with them. This conflicted with Mendeleevs
ordering by atomic mass because iodine weighs less than tellurium and therefore should
have been placed one group to the left. A similar problem arose with other pairs of
elements.
In 1913, the atomic number (number of protons) was found for elements. It was realized
that the elements fit being arranged in order of increasing atomic number.
Why is atomic number a better organizational system than atomic mass? Atomic number
is unique for each element, while atomic mass can vary. There is a potential for two
elements to have very close atomic masses.
Periodic Law states that when the elements are arranged in order of increasing atomic
number, there is a periodic repetition of their physical and chemical properties. This is
seen in the current periodic table as many of the physical and chemical properties either
increase or decrease as you move across a row.
Looking at the periodic table, there are numbers above each group. The groups numbers
1-18 are the international standard for group numbers. All international standards in
chemistry are set by IUPAC (International Organization for Pure and Applied
Chemistry). Previously scientists have used 1-8A and 1-8B to number the groups.
The elements can be grouped into 3 broad categories based on general properties: metals,
nonmetals, and semimetals (or metalloids).
Metals
-80% of elements
-man-made above Uranium
Physical Properties:
-good conductors of heat and electricity
-high luster of sheen (ability to reflect light); shiny
-solids at room temperature (except Mercury)
-ductile (can be shaped into wires)
-malleable (can be formed into thin sheets or other shapes)
Chemical Properties:
-reactivity (the ease and speed at which the metal combines or reacts with other elements;
metals tend to react with nonmetals)
-corrosion (the destruction of a metal by its reaction with oxygen)
Nonmetals
-15% of elements
-upper right corner of periodic table
Physical Properties:
-most are gases at room temperature, a few are solids, one is a liquid (bromine)
-because there are many varieties of nonmetals, the properties vary as well
-tend to have properties opposite of metals
-poor conductors of heat and electricity (insulators)
-solid nonmetals tend to be brittle (will shatter easily)
-dull in appearance
Chemical Properties:
-some are reactive (F is the most reactive element known) and some are unreactive (noble
gases hardly ever form compounds)
-can react with either metals or other nonmetals
-7 nonmetal (H, N, O, F, Cl, Br, and I) occur in nature as diatomic molecules (a molecule
made up of two of the same atom)
Metalloids (or Semi-Metals)
-most elements that border the step line between metals and nonmetals
-have properties similar to both metals and nonmetals
-behavior can be controlled by changing the condition they are in
Ex: Si is a poor conductor of electricity but when you add a small amount of B, it
becomes a good conductor of electricity
-most useful property is their varying ability to conduct electricity (they are good
semiconductors and can have conductivity between an insulator and a conductor)
Names of Groups Commonly Referred To
Hydrogen
-the chemical properties of hydrogen differ very much from those of the other
element and therefore it is not classified in any 1 group
-rarely found on Earth as a pure element (most is in the compound H2O)
Group 1, Alkali Metals
-from Arabic “al aqali” meaning the ashes; ashes are rich in Na and K which are
both in this group
-very reactive (never found as uncombined elements in nature; they are only
found in compounds)
-very shiny and soft (can cut with a knife)
Group 2, Alkali Earth Metals
-not as reactive as group 1, but more reactive than most other metals
-fairly hard, grey-white in color, and good conductors of electricity
Groups 3 – 12, Transition Metals
-hard and shiny
-good conductors of electricity
-slow to react (gold is the least reactive)
- Includes inner transition metals:
- Lanthanide series (elements 57 – 71)
-often mixed with more common metals to make alloys (an alloy is
a just a mixture of a metal with at least one other element)
- Actinide series (elements 89 – 103)
- all but 4 are manmade (everything after U is synthetic)
- the synthetic elements are have unstable nuclei and they last for
only a fraction of a second after they are made
- synthetic elements are made by forcing nuclear particles to crash
into one another
Group 17, Halogens
-the name is from the Greek word “hals” meaning salt; these are common
elements in salts
-the most reactive nonmetals
-uncombined elements in this group are very dangerous to humans
Group 18, Noble Gases
-also called inert gases because they rarely take part in reactions (unreactive)
-all exist in the Earth’s atmosphere but in very small amounts
Periodic Trends
There are definite trends that can be seen when looking at the properties of the elements
on the periodic table. Three that we are going to focus on are atomic size, ionization
energy, and electronegativity.
Atomic Size (aka Atomic Radius)
The size of an atom increases as you move right to left across a period and top to bottom
down a group (↓←).
Why is this the trend?
- (size decreases left to right across a period) - As you move to the right across a period,
the atomic number (also the protons) increases. This increase in positive charge causes
the electrons orbiting the nucleus to be pulled in closer to the nucleus. Therefore, with a
tighter orbit of electrons, the overall size of the atoms is smaller.
- (size increases going down a group) - As you move down a group you have more and
more energy levels being filled with electrons. Each energy level is a little further away
from the nucleus. The more energy levels an atom has, the less the electrons in the
outermost level can feel the pull from the positive nucleus so the atom gets bigger.
Ionization Energy
The energy required to remove an electron from an atom (thus making it an ion) is called
the ionization energy. Ionization energy increases as you move left to right across a
period and up a group (↑→).
Why is this the trend?
- (ionization energy increases left to right across a period) - As you move left to right
across a row, the number of electrons in the outermost energy level increases. Elements
on the left, which only have 1 or 2 electrons in the outermost energy level, can easily give
those up. Elements on the right, with 7 or 8 electrons in the outermost energy level, have
a much harder time just giving up 1 or 2 of them.
- (ionization energy increases up a group) - The less electrons the atom has, the closer
they are held to the positive nucleus and the harder it is to remove them.
Electronegativity
Electronegativity is the ability (or desire) of an atom to attract electrons when the atom is
in a compound. Electronegativity increases as you move left to right across a period and
up a group (↑→).
The most electronegative element on the periodic table is F (the upper right corner). The
least electronegative element is Cs (the lower right corner). The noble gases have an
electronegativity of 0 because they have completely filled energy levels and absolutely no
desire to gain any more electrons.
Ch. 5 Atomic Models
Can you see an atom? They are on the scale of 10-11 and 10-10 m.
What could you do if someone asks you to describe something that is too small to see?
How would you find out what it looks like?
Scientists did different experiments on atoms and used their findings to develop models.
A model is a representation to show the construction or appearance of something that is
either too small to see easily or too dangerous to experiment on. Models of the atom tried
to explain the observations scientists were making.

Dalton (1803): indestructible particle; solid sphere; “the atom”

Thompson (1897): discovered the electron; electrons on a solid sphere;
“plum pudding”
After the discovery of electons, scientists continued to experiment to determine how they
were put together in the atom. New atomic models emerged.

Rutherford’s Gold experiment (1911)
- shot positively charges (alpha) particles at a thin sheet of gold foil
- most went through, some bounced off at angles or came straight back
- Rutherford found this unusual since it should have been like “shooting bullets
through tissue paper”
- Proposed that the atom is mostly empty space and that’s why most of the
particles went through but that sometimes the particles hit a spot of concentrated
mass and positive charge causing them to repel back.
- named this spot the nucleus (the tiny central core of the atom)
Rutherford’s model is known as the “nuclear atom” model.
How is this model inadequate? It doesn’t explain a lot of the properties of atoms. When
you heat some metals, like for a horseshoe, the metal changes color as it heats (black to
red to yellow to white). Each of these colors represents a different level of energy. As it
gets hotter, other color appear.
What is it in the atom that relates to this energy? What changes in the atom depending on
how much energy it has? The electron. The behavior of electrons is what determines the
properties of atoms.
1912 - discovery of the proton
 Bohr (1913)
Student of Rutherford. Wanted to develop a model that would include how the atom’s
energy changes when it absorbs or emits light.
Electron is found in specific orbits around the nucleus. Each orbit has a fixed energy. The
orbits are called energy levels. An electron can gain or lose energy by changing its orbit.
It can only move from one orbit to another by gaining or losing the exact amount of
energy between the two orbits.
Just like when you’re climbing a ladder, you have to move completely to the next rung
up; you have to put in the amount of energy required to get to that next rung. You can’t
only move half a step up. Electrons can’t only move half and energy level. They need to
go to the next level or they don’t move at all.
A “quantum” of energy is the amount of energy required to move an electron from one
level to another. Quantum from Latin “quantus” meaning how much. The energy of an
electron if said to be quantized; can only be at distinct/discrete energies. There is a certain
amount of distinction between the levels due to the electrons in each orbit repelling each
other.
The energy levels in an atom are not evenly spaced. The further from the nucleus, the
close together the levels get. Why do you think that is? You are dealing with two charges
– a positive nucleus and a negative orbit (due to the electrons in it). Think of if you had
two magnets. When they are close together it takes a lot of energy to separate. Once they
are further apart, it is much easier to move.
When you put energy into an atom (in the form of light, heat, etc), as long as it is the
right amount of energy (one of the quantums of energy) the electron will increase its
energy to an energy level further from the nucleus from its normal position. Everything in
the universe likes to be neutral, stable, at a lower energy. An electron will return to its
original lower energy level as soon as possible after being excited to a higher one. When
this happens, the energy is released from the atom – usually in the form of light or heat.
For example, when you use glow in the dark stars, you are exciting the electrons to higher
energy levels using light energy. When the lights are turned off, the electrons relax back
to their lower energy levels. When this relaxation happens, the energy is released in the
form of light and the stars “glow” for a little while.
The Bohr model describes an exact path for the motion of the electron. This was ok in his
case because he developed his model based on the hydrogen atom which only has one
proton and one electron. What about when there are lots of electrons – do they all follow
the same paths?
 Schrodinger (1926)
Developed the electron cloud model (also known as the quantum mechanical model).
Still the modern description of the atom.
His model still has electrons in distinct/discrete energy levels however is does not
describe and exact path of motion, but instead describes how likely it is to find an
electron in various locations around the nucleus. Describes the probability of finding an
electron in different regions.
His model looks like a “cloud” around the nucleus with a high density right around the
nucleus and the density decreases the further from the nucleus you travel. The areas of
higher density correspond to areas of higher probability of finding an electron.
The model tells the probability of being in an area but doesn’t pinpoint position at any
instant. For example, when a fan is spinning, all you see is a blur – the blade is in there
somewhere but you can’t pinpoint an exact location.
Most of the electrons in the atom are going to be in the areas of the highest probability
(closest to the nucleus – at the lowest energy possible)
Will all electrons follow the same paths? What if you’re dealing with an atom with 100
electrons? Why don’t they follow the same paths? Repulsion.
When describing an electron in an atom you can describe 2 things: its energy level (the
energy the electron has) and its atomic orbital (the region of space where there is a high
probability of finding that electron)
Atomic orbital – region of probability (not necessarily such distinct boundaries)
There are 4 types of atomic orbitals:
Type
Shape
# of orbitals
s
spherical
1
p
figure-8/dumbbell
3
d
(don’t worry about it)
5
f
(don’t worry about it)
7
Total electrons it can hold
2
6
10
14
**do not confuse “atomic orbital” with “orbital”
Atomic orbital – region of probability of finding an electron
Orbit –distinct path of electron (more specific region within general region of probability)
Pauli Exclusion Principle - each orbital can hold up to 2 electrons and no more.
If you had to choose, would you rather be standing on 1 foot, 2 feet, or sitting? Why?
What takes less energy? Everything in the universe wants to be in a state of lowest
possible energy. Lowest energy = most stable.
In the atom, the electrons and the nucleus interact to make the most stable arrangement
possible.
Aufbau Principle – electrons occupy and fill the lowest energy levels first
Electron configuration – the way in which the electrons are arranged in the orbitals
around the nucleus
Aufbau performed experiments to determine the exact order of increasing energy for
atomic orbitals.
Aufbau Diagram: (for the lower energy levels)
3p
3s
2p
Energy
2s
1s
Nucleus
Each
represents one of the orbitals that type of atomic orbital contains.
1s2 1 = the energy level
s = the type of atomic orbital
2 = the number of electrons the orbital has in it
The orbitals listed in increasing energy:
1s22s22p63s23p64s23d104p65s24d105p6…….
This same increasing orbital energy can be found on the periodic table, and so does not
have to be memorized. The first 2 columns’ elements (group 1 and 2) have their
outermost energy level being s1 and s2. Groups 13 – 18 elements have their outermost
energy levels being p1, p2, p3, p4, p5, and p6 respectively. Groups 3-12 elements have their
outermost energy levels being d1 through d10 respectively. The inner transition metals
(Lanthanide and Actinide series) elements have their outermost energy levels being f1
through f14.
To determine the electron configuration, start at the top left of the periodic table and read
across (left to right) when you get to the end of the row go all the way back to the left and
down a row and read across (left to right).
Ex: H: 1s1
He: 1s2
Electrons want to be at the lowest possible energy. They will completely fill the lower
energy orbitals before starting to fill the higher ones. Once you move on to a higher
orbital, the ones below it will be completely full so you just write the full orbital (s2, p6,
etc).
Ex: Li: 1s22s1
Be: 1s22s2
B: 1s22s22p1
C: 1s22s22p2
Cl: 1s22s22p63s23s5
When you get to the part of the table where you start to include the d atomic orbitals (in
row 4), the d orbitals start with an energy level of 3 – even though they are in row 4. d
orbitals are in the energy level 1 less than the number of the row they’re in.
Ex: Mn: 1s22s22p63s23p64s23d5
Ni: 1s22s22p63s23p64s23d8
Zr: 1s22s22p63s23p64s23d104p65s24d2
When you get to the part of the table where you start to include the f atomic orbitals (in
row 6), the f orbitals start with an energy level of 4 – even though they are in row 6. f
orbitals are in the energy level 2 less than the number of the row they’re in.
Ex: La: 1s22s22p63s23p64s23d104p65s24d105p66s24f1
U:1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f4
There is a shorthand way to write the electron configurations. For whatever element you
are writing it for, you can shorten it to start with the noble gas before it (noble gases are
in group 18) and then continue with the configuration.
Ex: Cl: 1s22s22p63s23p5 = [Ne]3s23p5
Mn: 1s22s22p63s23p64s23d5 = [Ar]4s23d5
Zr: 1s22s22p63s23p64s23d104p65s24d2 = [Kr]5s24d2
Electrons are spinning and have either a clockwise or counter-clockwise spin. Each
orbital can hold up to 2 electrons (Pauli Exclusion Principle), one of each spin direction.
The two spin directions are represented by arrows:
and
.
The electron configuration is also represented on the Aufbau Diagram.
Hund’s Rule – one electron is placed in each orbital until all orbitals of the type contain
one electron of the same spin direction. Then a second electron is added with the opposite
spin to each orbital.
Ex. C: 1s22s22p2
Ex. F: 1s22s22p5
3p
3p
3s
3s
2p
E.
2s
1s
2p
E.
2s
1s
How many total electrons can each energy level hold?
Energy Level
Orbitals Included
1
1s2
2
2s22p6
3
3s23p63d10
4
4s24p64d104f14
5
5s25p65d105f14
6
6s26p66d10
7
7s27p6
Total Electrons
2
8
18
32
32
18
8
7 rows in the periodic table, 7 energy levels in the atom
The atom is the most stable when it has a completely full energy level (2 electrons for the
1st level; 8 electrons for the 2nd level; 18 electrons for the 3rd level; etc).
Even though the current model is an electron cloud of probability with all of the atomic
orbitals overlapping, we still draw orbital energy levels as circular orbits around the
nucleus for ease of discussion in class (and for illustration in your Flash representation).
The energy levels increase the further the distance from the nucleus.
For most elements, the electron configuration can be determined just by reading it off the
periodic table, like we have been doing. There are some instances however, where the
electrons fill the outermost energy level without the level below it being filled
completely.
Ex: Cr: from reading off periodic table:1s22s22p63s23p64s23d4
In actuality, it’s: 1s22s22p63s23p64s13d5
Cu: from reading off periodic table:1s22s22p63s23p64s23d9
In actuality, it’s: 1s22s22p63s23p64s13d10
In these rare cases (mainly with higher atomic number d elements), having the outermost
d energy level ½ or completely filled makes the atom more stable. Cr and Cu are the two
more common elements that have this occurrence. This phenomenon leads to the
sometimes unexpected behaviors of transition metal elements.
Electrons play a key role in determining properties of elements. There is a connection
between electron configuration and location on the periodic table. For example, the noble
gases are also called inert gases because they rarely take part in reactions. Why? The
outermost energy level of noble gases is full (it is stable) – there no desire to gain or lose
electrons.
The older group numbers (1A-8A) are above the columns collectively known as the
“representative elements.” The group number (1A-8A) represents the number of electrons
in the outermost occupied energy level.