5.3 Electron Configuration and Periodic Properites (2/2)
Summary
TLDRThis video explains key concepts in atomic physics, including electron affinity, ionization energy, and atomic/ionic radii. It discusses how atoms gain or lose electrons, with examples of halogens having the highest negative electron affinity due to their desire to form stable octets. The video also covers trends in the periodic table, explaining why electronegativity increases across a period and decreases down a group. Additionally, it introduces the concept of valence electrons and how chemical reactions are driven by electron interactions in the outermost energy levels of atoms.
Takeaways
- 🔋 Ionization energy is the energy required to remove an electron from an atom, while electron affinity is the energy released when an atom gains an electron.
- ⚛️ Adding an electron to an atom forms a negative ion and releases energy, which is expressed as negative because energy is released, not added.
- 🌀 Halogens, located next to noble gases on the periodic table, have the most negative electron affinity because they reach a stable octet by gaining electrons.
- 📉 Electron affinity becomes more negative across a period, with halogens having the most negative affinity, but group 6 has a larger electron affinity than group 7 due to orbital filling differences.
- 🧲 As you move down a group, it becomes harder to add electrons because they are farther from the nucleus, reducing the effective nuclear charge.
- 🔄 Second and third electron affinities require energy input, unlike the first affinity, because adding electrons to a negative ion creates repulsion.
- 📏 Cations (positive ions) are smaller than their neutral atoms due to fewer electrons and increased effective nuclear charge, while anions (negative ions) are larger because of additional electrons.
- 🧪 Metals tend to form cations, while non-metals form anions because metals lose electrons easily, and non-metals are close to achieving a stable noble gas configuration.
- 🔗 Valence electrons, located in the outermost energy levels, are responsible for chemical reactions, as they are most affected by external forces.
- 📊 Electronegativity measures an atom’s ability to attract electrons, with fluorine being the most electronegative element, and electronegativity increases across periods and decreases down groups.
Q & A
What is electron affinity, and how is it generally expressed?
-Electron affinity is the energy released when an atom acquires an electron, forming a negative ion. It is typically expressed as negative energy because it involves energy being released rather than put into the system.
Why do halogens have the most negative electron affinity?
-Halogens have the most negative electron affinity because they are one electron away from achieving a stable octet configuration. Adding an electron allows them to reach a lower energy state, releasing a significant amount of energy.
How does electron affinity change as you move across a period on the periodic table?
-As you move across a period, electron affinity generally becomes more negative, with halogens having the most negative electron affinity. This is because atoms get closer to filling their outer electron shells, releasing more energy when they acquire electrons.
Why does group 6 have a larger electron affinity than group 7?
-Group 6 elements have a larger electron affinity than group 7 because when you add an electron to a group 7 atom, it has to enter a sublevel that is already occupied, requiring more energy input, which reduces the overall energy released.
How does electron affinity change as you go down a group in the periodic table?
-As you go down a group, electron affinity becomes less negative. This is because the added electrons are farther from the nucleus, experiencing less nuclear attraction, and are shielded by inner electron levels.
Why does adding a second electron to an already negative ion require energy?
-Adding a second electron to a negative ion requires energy because the negative charge of the ion repels the incoming electron. This makes it energetically unfavorable, so energy must be supplied to overcome the repulsion.
How do the ionic radii of cations and anions compare to their neutral atoms?
-Cations have smaller ionic radii than their neutral atoms because they lose electrons, resulting in fewer electron-electron repulsions and a greater effective nuclear charge. Anions, on the other hand, have larger ionic radii due to the additional electrons increasing electron-electron repulsions, spreading them farther from the nucleus.
Why do metals tend to form cations and non-metals tend to form anions?
-Metals tend to form cations because they have low electron affinities, meaning they lose electrons easily. Non-metals tend to form anions because they are closer to achieving a stable noble gas configuration, so they gain electrons to complete their outer shell.
What are valence electrons, and why are they important in chemical reactions?
-Valence electrons are the outermost electrons of an atom and are involved in chemical reactions. They are the most susceptible to outside influence and interact with other atoms during reactions, as they are farthest from the nucleus and experience the least shielding from inner electrons.
How does electronegativity relate to electron affinity, and how is it measured?
-Electronegativity is a measure of how effectively an atom attracts electrons from another atom. It is related to electron affinity in that elements with high electronegativity also tend to have more negative electron affinities. Electronegativity is measured on a relative scale from 0 to 4, with fluorine assigned a value of 4 as the most electronegative element.
Outlines
🔬 Electron Affinity and Atomic Energy States
This paragraph discusses the concept of electron affinity, which is the energy released when an atom gains an electron. It explains that this energy is negative by convention, as it represents energy released rather than added to the system. The paragraph uses the example of halogens, which have the largest electron affinity (most negative) because adding an electron allows them to achieve a stable octet. The trend of electron affinity across the periodic table is also described, noting that it becomes more negative as you move from left to right across a period, with the exception of a decrease when moving from group six to group seven due to the energy required to add an electron to an already occupied sublevel. The paragraph also touches on the difficulty of adding electrons to atoms lower in a group due to increased distance from the nucleus and the presence of negative charge from other electrons.
🌐 Ionic Radius and Valence Electrons
The second paragraph delves into ionic radius, explaining that cations (positively charged ions) are smaller than their neutral atom counterparts due to fewer electrons and a greater effective nuclear charge. Conversely, anions (negatively charged ions) are larger because of additional electrons that spread out the electron cloud. The paragraph highlights that metals tend to form cations due to their low electron affinity, while non-metals form anions, aiming for a stable noble gas electron configuration. It also discusses how ionic radii increase down a group and decrease across a period, similar to atomic radii, due to changes in effective nuclear charge. The concept of valence electrons is introduced, emphasizing their role in chemical reactions and how they can be determined from an element's group number on the periodic table. The paragraph concludes with the idea that atoms strive for a stable octet of valence electrons, which explains the unreactivity of noble gases.
⚛️ Electronegativity and Chemical Bonding
The final paragraph focuses on electronegativity, which measures an element's ability to attract electrons in a compound. It mentions that electronegativity increases across a period and either decreases or remains constant down a group, with fluorine being the most electronegative element. The paragraph explains that electronegativity differences can lead to uneven charges in chemical bonds, with one atom becoming slightly positive and the other slightly negative. The scale of electronegativity is arbitrary, with fluorine assigned a value of four, and other elements' electronegativities are measured relative to it. The paragraph also notes that while the main group elements (s and p blocks) are the primary focus, the properties of d and f block elements related to electronegativity and other topics are not covered unless requested.
Mindmap
Keywords
💡Ionization energy
💡Electron affinity
💡Stable octet
💡Halogens
💡Cation
💡Anion
💡Electronegativity
💡Effective nuclear charge
💡Valence electrons
💡Noble gases
Highlights
Ionization energy refers to the energy required to remove an electron from an atom.
Electron affinity refers to the energy released when an atom acquires an electron, typically expressed as negative energy.
The halogens have the most negative electron affinity because they form a stable octet by filling their outer energy level.
Electron affinity generally becomes more negative as you move across a period, with halogens being the highest.
Group 6 has a larger electron affinity than group 7 because adding an electron to an already occupied sublevel requires more energy.
As you go down a group in the periodic table, the electron affinity decreases because added electrons are farther from the nucleus.
Second and third electron affinities are less convenient, often requiring energy to add additional electrons to a negatively charged ion.
Cations are smaller than neutral atoms due to fewer electrons and a greater effective nuclear charge.
Anions are larger than neutral atoms because they gain more electrons, expanding the electron cloud.
Metals tend to form cations, while non-metals tend to form anions due to differences in electron affinity.
Ionic radii increase down a group and decrease across a period, similar to atomic radii.
Valence electrons, found in the outermost energy level, are involved in chemical reactions.
The number of valence electrons can be determined by an element's group number in the periodic table.
Noble gases are unreactive because they already have a stable octet of eight valence electrons.
Electronegativity is the ability of an atom to attract electrons from another atom, with fluorine being the most electronegative element.
Transcripts
so just as atoms can have an electron
removed
which requires energy called the
ionization energy most atoms can also
acquire an electron
which releases energy
and this is called an atom's electron
affinity
now i'll give you an example that's just
a general equation for electron affinity
so you take
an atom you add an electron
and what you end up with is a negative
ion
and energy
and this energy is always expressed as
negative energy because by convention
energy is measured as energy put into
the system
so
if you add an electron to an atom and it
releases energy
that is the opposite of putting energy
in hence the negative sign which means
this energy is opposite putting energy
in so now if we look at electron
affinity in the periodic table
if you'll remember from our last video
that
the noble gases are in the lowest
possible
energy state
for that energy level
now
what this means is that the group right
next to them
the halogens
have the
largest electron affinity
but because it's negative they have the
least
electron affinity
so
this means that when they add
another electron
they form what's known as a stable octet
they fill their outer energy level
which
causes them to be in the lowest energy
possible
meaning that they go from a state of
very high energy to very low energy
relatively quick quickly which is why
they have
the most negative electron affinity that
is they release the most energy out of
any of the groups of elements
now
as a general rule electron affinity gets
more and more negative as you go across
a period eventually culminating with as
i mentioned the halogens being the
having the most negative electron
affinity
however when you go from group six to
group seven
you'll notice that
group six has a much larger electron
affinity than group seven and this is
because
in group six the carbon group
what you're doing
is you're filling up the final uh
sublevel within the p orbital with one
electron
now when you go to the seventh
when you go to the nitrogen group what
you have to do is you go from carbon's
uh
configuration
and then you add the electron into
a sublevel that's already occupied
and this requires a much larger energy
input which detracts from the energy
released
by this electron affinity
to form the
negative ion and as you go down groups
it gets harder and harder to add an
electron
because they are farther or farther away
from the nucleus so the effect of
nuclear charge they feel is much higher
up here where the electron cloud is much
smaller so they're closer to the nucleus
in the middle however when you get down
here
the nuclear charge they feel is
a lot smaller because
they're much farther from the nucleus
they're way out there
plus within that area there's a bunch of
negative charge from the other electrons
in the energy levels below it so what
you'll find is that as you get lowered
down a group
they release less and less energy as you
add electrons to form ions
so much like the second and third
ionization
the second and third electron affinities
aren't as convenient as the first one
so what you'll find is that the second
electron affinity that is when you take
a negative ion let's say f minus
and you try to add
another electron to it
it actually instead of releasing energy
and having a negative electron affinity
it requires energy because
this no longer has a neutral charge
instead it has a negative charge so this
electron
will repel that negative charge so it
requires positive energy to add
electrons for the second or third
etc
uh electron affinities
so now we're going to be covering ionic
radius and if you'll remember from
earlier an ion is just
an atom with a charge so either positive
or negative and now these each have a
name for example the positive
positive ions are called cations
which is easy to remember because the t
in cation looks like a positive sign
and then the negative is called an anion
and that is easy to remember because
it's not the cation
now the radius of cations is going to be
smaller
than that of the normal atom
because
they have fewer electrons and a greater
effective
nuclear charge on the outer electrons
since there are fewer electrons in the
middle sort of shielding them with
negative repulsion now the anions
oppositely
are going to be larger
than a neutral atom
because
they have more
electrons
in the electron cloud
meaning that they have to spread farther
and farther away from the nucleus
now if we go down and look at our table
what you'll find is that these metals
over here
will tend to form cations
because
they have a
very low electron affinity which means
they lose electrons very easily
and the non-metals which are over here
will tend to form
anions because they are very close to
the stable noble gas formation
over here on the far right
so what they want to do is add one more
or two more however many more electrons
to get to that stable formation
and just like atomic radii ionic radii
will tend to form
that snowman blowing bubble shape i was
talking about earlier that is the radius
will
increase as you go down the group
however it will decrease as you go
across a period again because of the
effective nuclear charge is much smaller
over here than it is over here
so now we'll be discussing valence
electrons
and the first thing you need to know is
that chemical reactions compounds and
molecules form by gaining losing or
sharing electrons it's not really about
the protons or neutrons in the nucleus
it's all about the electrons in the
cloud
that sort of engulf the nucleus
now
the electrons that are in the outermost
energy level that is if you
have an element say lithium which has
the configuration 1s2
2s1
the outermost energy level is the second
energy level and in this case there's
only the one electron which is in the s
sublevel
these are called the valence electrons
and these are the ones
that are involved
in chemical reactions
because
they are the most susceptible to outside
influence let me draw a quick lithium
atom let's say you have the nucleus
there and then you have the first energy
level
and the second energy level now on the
first energy level you have these two
electrons
and on the second you have just the one
now let's say
a proton was passing by
with its positive charge
this electron which is on the outside is
much closer to the proton
than these internal ones and these sort
of feel a repulsion from this electron
as well
so what ends up happening is that only
the electrons that are in this outer
energy level sort of feel the influence
of this proton or a different atom or an
ion or any sort of charged particle
which is why these
in the outer energy level are the ones
involved in reactions
and if you look at the periodic table
it's actually very easy to figure out
how many valence electrons an element
has
based on its group number
so i'm just going to go ahead and list
the number of valence electrons for each
group
and what you'll find is that first of
all for the s block it's just the group
number
so for groups one and two they have one
and two valence electrons respectively
however when you get over the p block
they're thrown off on group number by
this d block in here
so all you have to do is take the second
digit of each group for example
the boron group which is group 13
happens to have
three valence electrons or the carbon
group group 14
has four valence electrons etc and this
is true all the way across
and what you'll notice when you look at
chemical reactions is that
uh
all atoms are sort of trying to get to
this
number eight over here these eight
valence electrons called the stable
octet
and this is why the noble gases are so
unreactive
is because they already have this so
they don't need to give up
or share
or take on more electrons to get to this
stable eight valence electron
configuration so now we'll be discussing
a property called electronegativity and
as you'll remember
the valence electrons are the ones that
hold compounds and molecules together
now the problem with this is that
because these electrons are being shared
or given up
what happens
is that this can cause an uneven charge
across two atoms because if one atom has
taken another atom's electron this one
will become slightly positive and this
one will become slightly more negative
due to the difference in charge
now in order to measure how
easily
an element will take an electron from
another chemist created
a scale to measure how effectively an
element can attract
another electron
and this property of attraction
attracting
electrons from other elements is called
electronegativity so it's sort of an
arbitrary measurement because they just
do a scale from zero to four
and four is just given to
fluorine which through experimentation
they have determined to be the most
electronegative element and then
the electric the electronegativity
rather of all the other elements is
determined relative to fluorine
so now if we go down the gr go down to
the periodic table and look at how
electronegativity tends to uh go across
periods and groups what you'll find is
that
electronegativity is highest
over here on the right and lowest over
here on the left that is it tends to
increase as you go across a period and
it tends to either go down or remain
about the same
as you go down a group
and this of course makes sense because
as i mentioned before fluorine
which they determined to be the most
electronegative element
uh is assigned the arbitrary value
of four and then all the rest of them in
this vicinity
especially the halogens those uh gases
and liquids and solids within
fluorine's groove
tend to be very electronegative as well
so all the elements around fluorine are
very electronegative and then the ones
far away
tend to be
less electronegative which makes the
trends across the period and down the
group very easy to remember
now we work mostly with the main group
elements that is the s block and p block
in this course so i won't be covering
the properties of d and f block elements
as they relate to electronegativity
electron affinity and all the other
properties in this section unless it's
explicitly requested by you guys
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