5.3 Electron Configuration and Periodic Properites (2/2)

Summary

This 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

00:00

🔬 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.

05:02

🌐 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.

10:03

⚛️ 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.

Keywords

💡Ionization energy

Ionization energy is the energy required to remove an electron from an atom. In the video, it is mentioned as a contrast to electron affinity. Ionization energy is key in understanding how atoms lose electrons, influencing their chemical reactivity.

💡Electron affinity

Electron affinity refers to the amount of energy released when an atom gains an electron, forming a negative ion. This energy is always expressed as negative because energy is released rather than absorbed. The halogens are highlighted as having the most negative electron affinity, making them highly reactive in gaining electrons to achieve a stable octet.

💡Stable octet

A stable octet is a configuration where an atom's outermost energy level contains eight electrons, making it energetically stable. The video explains that halogens achieve this stable octet when they gain electrons, and this stability makes noble gases unreactive as they already possess this configuration.

💡Halogens

Halogens are elements in Group 17 of the periodic table, known for their high electron affinity and tendency to gain electrons to form negative ions. The video describes halogens as having the most negative electron affinity, which results in their strong ability to attract electrons and form stable compounds.

💡Cation

A cation is a positively charged ion formed when an atom loses electrons. The video explains that metals tend to form cations because they have low electron affinity, meaning they easily lose electrons, resulting in a smaller ionic radius compared to their neutral atomic form.

💡Anion

An anion is a negatively charged ion formed when an atom gains electrons. Non-metals, especially those close to the noble gases, tend to form anions. The video notes that anions have larger ionic radii than neutral atoms due to the increased electron repulsion in the electron cloud.

💡Electronegativity

Electronegativity measures how strongly an atom attracts electrons in a chemical bond. The video describes a scale from 0 to 4, with fluorine having the highest value. Electronegativity increases across periods and decreases down groups, affecting how atoms interact in molecules and compounds.

💡Effective nuclear charge

Effective nuclear charge is the net positive charge experienced by an electron in an atom, considering both the nucleus's attraction and the shielding effect of other electrons. The video explains that cations have a greater effective nuclear charge on their outer electrons because fewer electrons remain to shield the nucleus's pull.

💡Valence electrons

Valence electrons are the outermost electrons of an atom, involved in chemical reactions and bond formation. The video emphasizes that these electrons are the most affected by external forces like other atoms or ions, determining an element's chemical properties and reactivity.

💡Noble gases

Noble gases are elements in Group 18 of the periodic table, known for their full valence electron shell, making them chemically inert. The video highlights that noble gases already have a stable octet, which is why they do not need to gain or lose electrons and are very unreactive.

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