5.3 Electron Configuration and Periodic Properties (1/2)
Summary
TLDRThis educational video delves into electron configuration and its impact on periodic properties, focusing on atomic radius and ionization energy. It explains how atomic radius is determined by averaging the distance between nuclei in bonded atoms and trends across the periodic table, decreasing from left to right and increasing down groups. Ionization energy, the energy required to remove an electron, is highest for elements on the right and lowest for groups one and two, decreasing down groups due to increased atomic size and reduced effective nuclear charge. The video also discusses the concept of successive ionization energies, highlighting the significant increase when removing electrons from a noble gas configuration.
Takeaways
- 🔬 Atomic radius refers to the size of an atom, but due to the Heisenberg uncertainty principle, its boundaries are not definite.
- ⚛️ Chemists measure the atomic radius by binding two atoms, measuring the distance between their nuclei, and dividing that distance by two.
- 📉 Across a period on the periodic table, the atomic radius decreases as nuclear charge increases, pulling electrons closer to the nucleus.
- 📈 Going down a group, the atomic radius increases due to the addition of energy levels, which reduces the effective nuclear charge.
- 🎈 A 'snowman blowing bubbles' analogy helps illustrate how atomic size increases down a group and decreases across a period.
- ⚡ Ionization energy is the energy required to remove an electron from a neutral atom, resulting in the formation of an ion.
- 📊 Ionization energy trends: it increases across a period and decreases down a group due to changes in atomic radius and nuclear charge.
- 🔋 Removing additional electrons requires more energy, known as second ionization energy (IE2), which is higher than the first ionization energy (IE1).
- 🔒 Noble gas configurations are highly stable, and removing electrons from these configurations requires significantly more energy.
- 🌟 When an element reaches a noble gas configuration after ionization, the energy required to remove more electrons spikes drastically.
Q & A
What is atomic radius, and how is it measured?
-Atomic radius refers to the size of an atom, but due to the fuzzy boundary of an atom (as per Heisenberg's uncertainty principle), chemists measure it by binding two atoms together, measuring the distance between their nuclei, and then dividing that distance by two.
How does atomic radius change across a period on the periodic table?
-As you move across a period from left to right, the atomic radius decreases. This is because the number of protons in the nucleus increases, leading to a stronger nuclear charge that pulls the electrons closer to the nucleus.
Why does atomic radius increase as you move down a group in the periodic table?
-Atomic radius increases as you move down a group because more energy levels (electron shells) are added, making the atom larger. Although the nuclear charge increases, the additional energy levels and inner electrons reduce the effective nuclear charge felt by outer electrons, allowing them to be farther from the nucleus.
What is ionization energy, and how is it measured?
-Ionization energy is the energy required to remove an electron from a neutral atom to form a positively charged ion. It is measured in kilojoules per mole (kJ/mol), representing the energy needed to ionize one mole of atoms.
How does ionization energy trend across a period?
-Ionization energy increases as you move across a period from left to right. This is because atoms on the right side of the periodic table have a smaller atomic radius, meaning their electrons are closer to the nucleus and experience a stronger nuclear charge, making it harder to remove an electron.
Why does ionization energy decrease down a group?
-Ionization energy decreases as you move down a group because the outermost electrons are farther from the nucleus and experience less effective nuclear charge. This makes it easier to remove an electron, requiring less energy.
What is the second ionization energy, and why is it higher than the first ionization energy?
-The second ionization energy is the energy required to remove a second electron from an ion that has already lost one electron. It is always higher than the first ionization energy because, after the first electron is removed, the remaining electrons experience a stronger attraction to the nucleus due to the reduced electron shielding, making them harder to remove.
Why is there a large increase in ionization energy after removing electrons that result in a noble gas configuration?
-There is a large increase in ionization energy when removing electrons that result in a noble gas configuration because noble gases are very stable and unreactive. Their electron configurations are low-energy and stable, so removing an additional electron requires significantly more energy.
What analogy is used to explain the atomic radius trend down a group?
-The analogy of a snowman blowing bubbles is used to explain the trend. The bottom of the snowman represents the larger atoms at the bottom of a group, while the bubbles represent how atomic radius increases as you move down the group.
What is the relationship between atomic radius and ionization energy?
-Atomic radius and ionization energy are inversely related. As atomic radius increases (like down a group), the ionization energy decreases because the outer electrons are farther from the nucleus and easier to remove. Conversely, as atomic radius decreases (across a period), ionization energy increases because the electrons are closer to the nucleus and harder to remove.
Outlines
📏 Understanding Atomic Radius
The atomic radius refers to the size of an atom, but due to the uncertainty principle, its boundaries are unclear. To measure it, chemists often bind two atoms, measure the distance between their nuclei, and divide it by two. The atomic radius decreases across a period due to increased nuclear charge but increases down a group because added energy levels outweigh the attraction of nuclear charge. This trend can be visualized using the metaphor of a snowman blowing bubbles, where the atomic size gets larger as you go down a group.
⚡ Introduction to Ionization Energy
Ionization energy is the energy needed to remove an electron from a neutral atom, forming a positively charged ion. The process is called ionization, and the energy is typically measured in kilojoules per mole. Elements on the right side of the periodic table have higher ionization energies, while groups 1 and 2 on the left have lower ionization energies. This is because their outer electrons are farther from the nucleus, experiencing less nuclear charge, making them easier to remove.
Mindmap
Keywords
💡Atomic Radius
💡Heisenberg's Uncertainty Principle
💡Nuclear Charge
💡Ionization Energy
💡Effective Nuclear Charge
💡Periodic Trends
💡Electron Configuration
💡Ion
💡Noble Gas Configuration
💡Second Ionization Energy
💡Stable Noble Gas Formation
Highlights
Introduction to Chapter 5 Section 3, focusing on electron configuration and periodic properties.
Explanation of atomic radius: Chemists measure the distance between nuclei of two bonded atoms and divide it by two.
Trend of atomic radius across a period: The radius decreases as you move across a period due to increased nuclear charge.
Trend of atomic radius down a group: The radius increases as you move down a group due to added energy levels and electron shielding.
Analogy of a snowman blowing bubbles to help remember atomic radius trends across periods and groups.
Introduction to ionization energy: The energy required to remove an electron from a neutral atom.
Definition of ionization and formation of ions, such as the sodium ion (Na+).
Measurement of ionization energy in kilojoules per mole (kJ/mol) and its importance in comparing elements' ability to lose electrons.
Trend of ionization energy across a period: It increases from left to right due to decreasing atomic radius and stronger nuclear charge.
Trend of ionization energy down a group: It decreases as atoms get larger and electron shielding reduces effective nuclear charge.
Explanation of second ionization energy (IE2): Always higher than the first ionization energy (IE1) due to increased attraction between remaining electrons and nucleus.
Large jump in ionization energy when moving from IE1 to IE2, especially when the second electron is removed from a stable noble gas configuration.
Explanation of why removing electrons from noble gas configurations is challenging due to their stability and low energy state.
Key example: Removal of a second electron from lithium requires much more energy because it disrupts a stable configuration.
General conclusion: Ionization energies show a pattern across the periodic table, with significant spikes when removing electrons from noble gas configurations.
Transcripts
all right so in this video we're going
to be covering chapter 5 Section three
which is electron configuration and
periodic properties and the first of
these properties that we're going to be
covering is something called atomic
radius now as you may have guessed uh
the atomic radius has to do with the
size of an atom however because the
border of where an atom sort of ends
this is a very fuzzy line due to
Heisenberg's uncertainty principle and
the behavior of electrons so you can't
just assign a radius R definitively and
say that's the atomic radius so what
chemists do instead is they'll take two
atoms and bind them together let's say
This Is two hydrogen
atoms and then they take the distance
between the nucleus or nuclei
rather and then what they do to find the
radius is they just take this and divide
it by two and then you get what we call
the atomic
radius now if we look how this Trends
across a period on the table uh we can
see what you'll notice is that as you
go across the
period the radius of an atom starts out
bigger and then gets smaller and smaller
and smaller as you
go
because as you get down over towards
this end there's a stronger Nu nuclear
charge duee to the extra protons like
there's five protons and borons 6 7 Etc
as you go all the way
down and then as you go down a group you
may expect that it would get even
smaller because these atoms down here
have a large nuclear charge as well
however what you'll notice is that as
you go down the group they get larger
over and over again because as you add
energy
levels what ends up happening is that
takes so much
space that it overcomes the uh
attraction of the nuclear charge plus as
you add energy levels there are more and
more electrons inside each energy level
which reduces the effective nuclear
charge that is if you take the net
positive charge due to the nucleus and
then add the negative charge due to
these electrons
you'll find that they experience less
charge as you go further and further out
and an easy way to remember this is to
think of it as a snowman blowing bubbles
you make
the the lowest sphere of a Snowman the
largest with its head at the top being
the smallest sphere and then as you go
across you can see it's blowing bubbles
now I know the bubbles aren't getting
bigger however it makes it easy to
remember the group
Trend now the next thing we're going to
be covering is a property called
ionization energy and before we do that
what you have to know is that if you
take an atom given by this simplified
nucleus and electron cloud here and you
apply enough energy to it what you
conventionally get is you can remove an
electron from the
atom and I'll give you a mathematical
representation so let's take an atom a
you add some
energy and then what you get
is
A+ because you took away one negative
charge so it leans more positive now
that you've gotten one negative away
from
neutral plus an
electron now an ion which is what this
is called the
A+ uh is an atom or molecule that has a
positive or A negative charge basically
it's an atom or molecule that is
n neutral for example if you take a
sodium atom and then take away one of
its electrons what you'll end up with is
the sodium ion
na+ and the electron out in free space
now any process like this where you take
a neutral atom and end up with an ion at
the end is a process known as
ionization and and what chemists will do
is in order to compare how
easily U elements can give up this
electron they will measure the energy
required to remove one electron from a
neutral atom like
sodium and this is measured in
kles per
mole meaning the amount of energy in
kles required to remove elect R from one
mole of substance so now that we know
about ionization energy the next thing
we have to do is look at how it Trends
across the table and if you look at 15
figure 15 in your book which lists the
ionization energies for various elements
what you'll find is that they are
highest over here on the right and
lowest in groups one and two over here
on the left now this isn't just
coincidence because uh the these first
two groups if you'll remember have
electrons in just the very lowest levels
but of a new energy level so they are
farthest from the nucleus if you'll
remember the atomic radi you have the
biggest atoms over here and the smallest
over here now it's very hard to remove
an electron from uh a shell that is
closer to the nucleus because there's a
bigger positive charge but if an
electron is just sort of floating
Loosely out in a brand new energy level
what you'll find is that the net
positive charge it feels is much smaller
and it's much easier to take away an
electron and form an ion like na+ atomic
radius and nuclear charge also affect uh
ionization energy going down a group if
you'll remember from the Snowman blowing
bubbles uh atoms tend to get larger as
you go down a group which means that the
electrons are farther away with more
electrons between them and the nucleus
meaning their effective nuclear charge
that they feel is much smaller so what
you'll find is that the ionization
energy decreases as you go down a group
now you can also remove electrons from
ions like let's say you had taken away
one electron from lithium to make it
lithium plus which has three
protons and two electrons
giving it the net positive
charge now what you could do is you
could take away another
electron giving you three protons and
one electron
sorry and this process is called the
same thing this requires a huge amount
of energy called the ionization energy
however it's called the second
ionization energy represented by the
symbol
ie2 now ie2 2 is always going to be
greater than ie1 because let's say you
have the three electrons sort of around
lithium like this again not an accurate
model but it's it'll work for what we
need to do here if you take away this
electron right
here what you'll notice is that you have
the same nuclear charge in the middle
the positive
however these two electrons now feel a
greater net force because there's not
the extra electron around them and this
ends up decreasing the atomic
radius which leads to an increase in
ionization energy as we just cut now if
you look over at this plot of the
various ionization energies for lithium
you'll notice that from its initial
ionization energy to its second
ionization energy there is a huge jump
and this is because after you remove the
first
electron from lithium you have to come
all the way back over here to helium
which is covered up however helium has a
stable noble gas formation and what
you'll find is that removing an electron
from a stable from a noble gas uh
configuration is very difficult because
noble gases are in a very low energy
state they're very stable very
unreactive and nature wants to keep it
that way so removing this second
electron causes a huge
Spike and what you'll find is that
whenever you end up with a noble gas
uh configuration for an element after
you've ionized it for example after you
go to after you uh take away two
electrons from burum
and get its second ionization energy
there's a huge Spike to its third
because you're taking away from what was
then a noble gas configuration and this
news all across the period
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