5.3 Electron Configuration and Periodic Properties (1/2)

Summary

This 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

00:00

📏 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.

05:02

⚡ 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.

Keywords

💡Atomic Radius

Atomic radius refers to the size of an atom, which is not precisely defined due to the fuzzy boundary where an atom ends. In the video, it is explained that chemists determine atomic radius by measuring the distance between the nuclei of two bonded atoms and dividing it by two. This concept is crucial for understanding periodic trends, as the atomic radius changes across the periodic table, generally decreasing across a period and increasing down a group.

💡Heisenberg's Uncertainty Principle

Heisenberg's Uncertainty Principle is a fundamental concept in quantum mechanics that states the position and momentum of a particle cannot both be precisely measured at the same time. In the context of the video, it is mentioned to explain why atomic radius cannot be defined with absolute precision, as the behavior of electrons is subject to this principle.

💡Nuclear Charge

Nuclear charge is the positive charge of an atomic nucleus due to the presence of protons. The video explains that as you move across the periodic table, the nuclear charge increases due to the addition of protons, which affects the atomic radius and ionization energy. A stronger nuclear charge pulls electrons closer, resulting in a smaller atomic radius and higher ionization energy.

💡Ionization Energy

Ionization energy is the energy required to remove an electron from an atom or molecule. The video discusses how ionization energy is measured in kilojoules per mole and is used to compare the ease with which elements can lose an electron. It is shown to decrease down a group due to increasing atomic size and decreasing effective nuclear charge.

💡Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. It is calculated by subtracting the charge of inner electrons from the total nuclear charge. The video uses the analogy of a snowman blowing bubbles to illustrate how effective nuclear charge decreases as you move down a group, leading to larger atomic radii and lower ionization energies.

💡Periodic Trends

Periodic trends refer to the patterns in the properties of elements as they are arranged in the periodic table. The video focuses on trends in atomic radius and ionization energy, explaining how these properties change across periods and down groups in the periodic table.

💡Electron Configuration

Electron configuration is the distribution of electrons in an atom's energy levels and sublevels. The video script mentions this concept as a foundational understanding for discussing atomic radius and ionization energy, as the arrangement of electrons influences these properties.

💡Ion

An ion is an atom or molecule that has a net electric charge due to the loss or gain of one or more electrons. The video explains the process of ionization, where a neutral atom becomes an ion by losing an electron, and how this relates to ionization energy.

💡Noble Gas Configuration

A noble gas configuration refers to an electron arrangement that resembles that of the noble gases, which are known for their stability and reluctance to react chemically. The video discusses how reaching a noble gas configuration after ionization can result in a significant increase in ionization energy due to the stability of such configurations.

💡Second Ionization Energy

Second ionization energy is the energy required to remove a second electron from an atom that has already lost one electron. The video script explains that the second ionization energy is typically higher than the first because the remaining electrons experience a greater effective nuclear charge after the first electron is removed.

💡Stable Noble Gas Formation

Stable noble gas formation refers to the electron configuration that an atom achieves when it resembles the noble gases in terms of electron count. The video mentions that removing an electron from an atom with a stable noble gas configuration is very difficult, as it requires a significant amount of energy, leading to a spike in ionization energy.

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