Chapter 7: Ionic Bonding
• Background (chapter 4): The 3 major
subatomic particles: protons, neutrons,
and electrons.
• The identity of an element is determined
by the number of protons, (the positivelycharged particles) that are found in each
atoms' nucleus .
• The number of protons is referred to as
the "atomic number" of the atom. To
make the math easy for us, we say that
each proton has a charge = __+1__.
• The element hydrogen is the simplest
element, and has the simplest atoms: An
atom of the element hydrogen contains
only _1__ proton in its nucleus;
hydrogen's atomic number is 1.
• An atom of the element helium contains
__2__ protons in its nucleus; helium's
atomic number is 2. Each element from
atomic number 3 through 118 has a
consecutive increase in the number of its
protons by ___1_____.
• The nucleus is the high density part of an
atom, taking up very little of the
atom’s_volume_; but, accounting for
practically all of its _mass_. This is
because, in addition to all the protons, the
nucleus also contains neutrons (_neutral_
particles having no electrical charge);
and, both the protons and the neutrons
are each about 2000 x _more_ massive
than the very tiny electron.
• While the nucleus by itself is positivelycharged (due the protons within it),
atoms as a whole are electrically
_neutral_ due to their having a number
of _negatively_-charged electrons
_equal_ to their number of protons.
We say that each electron has a charge
= -1. The charge of each 1 electron
exactly cancels the charge of each 1
proton.
• As long as an atom remains neutral, the
number of its electrons will equal the
number of its protons.
• To be neutral means that, although there
may be lots of positively– charged and
lots of negatively– charged particles
present in each atom, overall the charges
all cancel out, thus the “net (overall)”
charge of a neutral atom equals 0.
Background, Chapter 5:
• The electron is the smallest of the 3 major subatomic
particles; and, these very energetic particles, though
_strongly attracted_ to the positively-charged nucleus,
are able to travel quite far away from the nucleus, flying
around the nucleus at a speed close to the _speed_ _of_
_light_.
• If the nucleus were the size of a marble, the electrons
would occupy about the volume of a football stadium.
Due to the large volume of space the tiny electrons fly
through, atoms are actually mostly _empty_ _space_.
• Niels Bohr suggested that electrons “orbit” the nucleus
of an atom at exact distances away from the nucleus.
New information, Chapter 5:
• Electrons do NOT_ reside in specific orbits!
There is no such thing as an “orbit” (no such
thing as an exact distance away from the
nucleus) to which electrons are limited!
• What _Niels_ _Bohr_ was correct about was
this: Electrons have certain “allowable
energies_”. It’s not that they are found in
certain allowable orbits away from the
nucleus; instead, they are found with certain
allowable energies, and they are always
outside the nucleus.
• When we refer to an electron residing in an
_energy_ _level_ (later on in the notes), we are
really referring to an amount of energy that an
electron has at that moment of time.
• And, the electron can’t just have any ole’ amount
of energy….The atom itself seems to dictate
specific allowable quantities of energy that its
electrons might have.
• In much the same way, you may have a friend that
is always either asleep in class (low energy state),
or driving every one crazy with his/her exuberance
(high energy state); but, you’ve never seen your
friend exhibit a moderate amount of energy.
• It’s like that with the electrons….They jump
from energy level to energy level, showing
sometimes dramatic increases in energy and
sometimes showing dramatic decreases in
energy; but, never showing an amount of
energy, or a change in energy, placing them “
in between” their routinely observed energies.
• We believe that electrons, based on the atom
that they’re found in, have certain “allowable”
energies: This is what is meant by the
statement, “The energy of electrons is
quantized”. It comes and goes in _packets_
called _quanta_.
• When all of the electrons in an atom are
at the lowest possible energy state, we
say that the atom is at _ground_
_state__ conditions.
• When even just one of the electrons in
an atom is at a higher energy state
(than the lowest possible), we say that
the atom is in the _excited__ state_.
• Since there is no such thing as an “orbit”, we
can’t ever pinpoint a specific location that a
specific electron will be found in at a
_specific_ moment in time; instead, we can
only refer to there being a 90% or
better_probability_ of finding an electron in a
given region of space at any time.
• In much the same way, I can’t predict exactly
where I will be at 6:30 am, as I drive in to
school; but, I can say that there is a 90%
probability I will be somewhere along route 81.
• During the early days of the developing
“quantum theory of electrons”, De Broglie
and Schrodinger called this “probable
region”, that is, the region where there is
a > 90% chance of finding an electron, an
“atomic_ _orbital_".
• Atomic orbitals can be envisioned as
“electron _clouds_” surrounding the
atom's nucleus, in that, like a regular
cloud, the edges of any given atomic
orbital are _NOT_ distinct boundaries.
• Atomic orbitals don’t have defined edges;
and, you might have anticipated, there is a
10% chance that an electron will be found
outside of its most probable region.
• In much the same way, there is a small
probability that I might be found
somewhere along the shoulder of route
81, or even still, on route 114, at 6:30 am.
• Atomic orbitals, like clouds, form different
_shapes_.
• Some atomic orbitals look like _spheres_,
others look like _dumbbells_ (hand weights),
and some look like four-leafed_ clovers.
– The spherically-shaped orbitals are called "_s___
_orbitals__",
– the dumb-bell shaped orbitals are called "__p__
_orbitals_",
– the four-leafed clover-shaped orbitals are called
"__d__ orbitals_".
– The last group of orbitals have complex-shapes and
are called "__f__ _orbitals_".
• Electrons found in different types of
orbitals (s, p, d, or f) will have
_different_ amounts of energy.
• And, all of the electrons found in the
same type of orbital will have about the
_same_ amount of _energy_ AS LONG
AS _the electrons are in the same
energy level.
• Niels Bohr was the person who
introduced the concept of the “_energy__
__level_”.
– Think of an energy level as being similar to one
layer of growth in an onion, or one layer of
growth in a tree.
– Onions and trees are 3–dimensional objects
which grow in layers that extend outwards:
– In a similar manner, energy levels are 3
dimensional _spaces___ or regions extending
outwards from a nucleus. That’s the end of the
similarity; but, hold onto that thought.
• Each energy level holds electrons
which have a range of similar energies.
• Electrons having a given amount of
energy will be found in a given energy
level; and, as electrons gain energy,
they will move up to “_higher_” energy
levels which are found out _farther
away_ from the atom’s nucleus.
•
Most energy levels (in fact, all but one of
the energy levels in an atom) have, during
the establishment of _quantum theory__,
been sub–divided into what are called
_sublevels__, based on the type of
orbital(s) found in that energy level.
– “s” sublevels contain electrons in “s” orbitals.
– “p” sublevels contain electrons in “p” orbitals.
• As long as electrons are in the _same_ energy
level, the electrons residing in the s-orbitals
(the s sublevel) have the _same approximate_
energy.
• More energetic electrons (in that same energy
level) will be found in the p-orbitals (the p
sublevel), even more energetic electrons (in
that same energy level) will be found in the dorbitals (the d sublevel), and the __most__
energetic electrons (in that same energy level)
will be found in f-orbitals (the f sublevel).
• Memory aid for the order, from least
energy to most energy within an energy
level: spdf ... __some people don’t fart_
•
In class we will need to become comfortable
working with sublevels; and, we will find out
that the two most important sublevels (when
trying to understand_chemical_ changes) are
the __s_ sublevel and the _p_ sublevel.
– The symbol, “1s”, is our symbol for the “1 s
_sublevel_”, and is our way of saying, “first__ energy
level”, “s orbitals only”.
– The symbol, “2p”, is our symbol for the “2 p
sublevel”, and is our way of saying 2nd energy level,
___p___ _orbitals__ only.
– Etc. What is our symbol for saying, “3rd energy level,
d orbitals only”? ___3d__
• There are specific numbers of electrons that
an orbital can hold maximally, specific
numbers of electrons that sublevels can hold
maximally, and specific numbers of electrons
that energy levels can hold maximally.
• Thus, there is a specific organization of
orbitals, sublevels, and energy levels, within
each atom.
– There are several different patterns that occur.
– Understanding these patterns will help you keep
the organization straight in your mind.
• – Here is a pattern between energy levels and sublevels.
Can you see the pattern?
– 1st energy level has ____1_____ sublevel,
identified as: __1s__
– 2nd energy level has _2__sublevels,
– identified as: __2s__ __2p__
– 3rd energy level has __3__ sublevels,
– identified as: __3s__ _3p___ _3d___
– 4th energy level has _4_ sublevels,
– identified as: _4s___ __4p__ _4d___ _4f___
– 5th energy level has _4_ sublevels,
– identified as: __5s__ _5p___ _5d___ _5f___
– 6th energy level has ___3__ sublevels,
– identified as: _6s__ __6p__ _6d___
– 7th energy level has __2__ sublevels,
– identified as: _7s__ __7p__
• There are also patterns in the number
of orbitals in a sublevel, and in the
maximal number of electrons in a
sublevel. Can you see the patterns?
Orbital
Type
# of orbitals in
one sublevel:
Each orbital
may hold this
many electrons
maximally:
s
1
p
3
2
d
5
2
10
f
7
2
14
X
2
=
Maximal number
Of electrons
that fits in
this sublevel
2
6
– By combining the two patterns, you can see that
the formula for determining the number of
electrons held maximally within the first four
energy levels = __2n2___
• where “n” = _number of the energy level (1
through 7)
“n” =
energy
level
1st
2nd
3rd
4th
5th
6th
7th
Identity of the
sublevels in the
energy level.
1s
2s 2p
3s 3p 3d
4s 4p 4d
4f
5s 5p 5d
5f
6s 6p 6d
7s 7p
# of electrons maximally found in
the first 4 energy levels = 2n2
The element with atomic
118
number _______ is the
2
2
2
2
2
2
2
= 2
+ 6
= 8
+ 6 + 10
= 18
+ 6 + 10 + 14 = 32
+ 6 + 10 + 14 = 32
+ 6 + 10
= 18
+ 6
= 8
Sum of the electrons: 118
most complex element in
our known universe, and
this system of organization
can account for all
______of its electrons!
118
• The “__diagonal_ rule” (next page) uses the
sublevel pattern to help scientists generate
what is called an “_electron_ _configuration_”
and an _orbital_ _diagram_ when the atom is
in its _ground state_ (lowest energy) condition.
– You are always allowed to use your reference
packet, which includes a diagram of the diagonal
rule (and the periodic table, of course, is always
available);
– but, while the energy ladder is displayed in our
classroom – it is not guaranteed to be visible to you
during the unit exam, the midterm and/or the final
exams.
• To use the diagonal rule, just read off
the sublevels one at a time, as you
follow along the slanted arrows (from
_right_ to _left_), starting at the top of
the diagram and ending at the bottom.
Additionally, state the number of
electrons found in each sublevel.
•
•
•
•
•
•
•
•
•
•
•
•
•
s-2e
1s
p-6e
d-10e
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p
f-14e
•
When reading from the diagonal rule correctly, you generate a correctly ordered list of sublevels:
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
6s
4f
5d
6p
7s
5f
6d
7p
When working with an atom of neon, which only has 10 electrons, its electron configuration would
be this: 1s2 2s2 2p6
*as you follow the diagonal arrows, add up the exponents till you get to 10, then stop. See note
below.
In other words, an electron configuration is a list of sublevels, each being “filled” with the maximum number of
electrons for that given sublevel.
The list of maximum electrons is given to the right side of the diagonal rule diagram above.
What would be the electron configuration for an atom of argon?
Atomic number = 18. so, 18 electrons.
1s2 2s2 2p6 3s2 3p6
2 + 2 + 6 + 2 + 6 = 18
Follow the diagonal arrows,
Maximum exponent for every “s” = 2. Max every “p” = 6.
max “d” = 10.
max “f” = 14”
Orbital Filling Ladder
7p
___ ___ ___
6d
___ ___ ___ ___ ___
5f
___ ___ ___ ___ ___ ___ ___
7s
___
6p
___ ___ ___
5d
___ ___ ___ ___ ___
4f
___ ___ ___ ___ ___ ___ ___
6s
___
5p
___ ___ ___
4d
___ ___ ___ ___ ___
5s
___
4p
___ ___ ___
3d
___ ___ ___ ___ ___
4s
___
3p
___ ___ ___
3s
___
2p
___ ___ ___
2s
___
1s
___
Neon, 10 electrons:
Read directions (next page in notes)
Try Argon, 18 electrons on your own:
When reading from the orbital filling
ladder correctly, you generate a
correctly ordered list of sublevels, and
you can also visually see the number of
orbitals in each sublevel.
To read from the ladder correctly, start
at the bottom and proceed upwards!
• Lastly, as per the instructions on your last
page of notes, try writing out an electron
configuration and orbital diagram for nitrogen.
• If you don’t yet have the webquest handout,
now is the time to request one!