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Lecture Comments (9)

0 answers

Post by Parth Shorey on April 25 at 07:23:00 PM

I still don't understand why paramagnetic is attracted to external magnetic field?

1 answer

Last reply by: Professor Franklin Ow
Thu May 28, 2015 12:36 PM

Post by Iman Abbas on May 27, 2015

Why is that when we talk about electron affinity and ionization, we refer to just the gas phase only?

0 answers

Post by Saadman Elman on November 27, 2014

Thanks for the video. The explanation was very explicit.

2 answers

Last reply by: Shawn Freeman
Sun May 24, 2015 6:33 AM

Post by John Wadsworth on December 7, 2013

In the two exceptions, the expected configuration of Cr would be [Ar]4s^2 3d^4 and not [Ar]4s^2 3d^6.

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:12 PM

Post by Baoer Ye on October 8, 2013

why is Cl in the 7th column number?And O in the  6th ?

Related Articles:

Periodic Trends

  • A consequence of quantum mechanics was the establishment of certain periodic trends, including ionization energy, electron affinity, atomic size and electronegativity.
  • The periodic table is very organized and informative concerning an element’s physical and chemical properties.
  • An atom’s electron configuration can be determined following the Aufbau Principle.
  • Orbital box diagrams are drawn following the Pauli Exclusion Principle and Hund’s Rule.

Periodic Trends

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Lesson Overview 0:09
  • Introduction 0:36
  • Electron Configuration of Atoms 1:33
    • Electron Configuration & Atom's Electrons
    • Electron Configuration Format
  • Electron Configuration of Atoms Cont'd 3:01
    • Aufbau Principle
  • Electron Configuration of Atoms Cont'd 6:53
    • Electron Configuration Format 1: Li, O, and Cl
    • Electron Configuration Format 2: Li, O, and Cl
  • Electron Configuration of Atoms Cont'd 12:48
    • Orbital Box Diagrams
    • Pauli Exclusion Principle
    • Hund's Rule
  • Electron Configuration of Atoms Cont'd 17:35
    • Exceptions to The Aufbau Principle: Cr
    • Exceptions to The Aufbau Principle: Cu
  • Electron Configuration of Atoms Cont'd 20:22
    • Electron Configuration of Monatomic Ions: Al
    • Electron Configuration of Monatomic Ions: Al³⁺
    • Electron Configuration of Monatomic Ions: Cl
    • Electron Configuration of Monatomic Ions: Cl¹⁻
  • Electron Configuration Cont'd 24:31
    • Paramagnetism
    • Diamagnetism
  • Atomic Radii 26:08
    • Atomic Radii
    • In a Column of the Periodic Table
    • In a Row of the Periodic Table
  • Ionic Radii 27:30
    • Ionic Radii
    • Anions
    • Cations
    • Isoelectronic Species
  • Ionization Energy 29:00
    • Ionization Energy
  • Electron Affinity 31:37
    • Electron Affinity
  • Summary 33:43
  • Sample Problem 1: Ground State Configuration and Orbital Box Diagram 34:21
    • Fe
    • P
  • Sample Problem 2 36:38
    • Which Has The Larger Ionization Energy: Na or Li?
    • Which Has The Larger Atomic Size: O or N ?
    • Which Has The Larger Atomic Size: O²⁻ or N³⁻ ?

Transcription: Periodic Trends

Hi, welcome back to

Today's presentation from general chemistry is going to be on periodic trends.0003

We are going to go ahead and as always start off with a brief introduction.0009

Then we are going to get into something we call electron configuration of atoms.0014

After electron configuration of atoms, we are then going to see how we can use the periodic table0018

to help us determine electron configurations--that is what we call the periodic trend.0023

After that, we are going to go ahead and look at additional periodic trends.0028

Followed by summary and finally a few sample problems.0032

We in the previous presentation discussed quantum mechanics and how it gave rise to what was called an atomic orbital.0038

One of the grand outcomes of quantum mechanics was the establishment of periodicity.0048

What periodicity refers to is the following.0055

That patterns of elements and how they are organized throughout the periodic table can be quite predictable0058

just off of organization in their location on the periodic table.0069

The patterns that we are dealing with have to do with both the physical and the chemical properties of atoms.0075

For today this presentation, we are just going to look at a series of these periodic trends.0082

The first one we are going to look at is what is called electron configuration.0089

Electron configuration of an atom is going to answer the following questions about an atom's electrons.0095

Number one, what type of atomic orbitals are involved--s, p, d, or f?0102

Number two, how many electrons are in each of these orbitals?0111

Electron configuration typically has the following format.0117

It is related to quantum numbers actually.0122

It is always the principal quantum number n, followed by the letter of the atomic orbital.0126

That is going to be s, p, d, or f.0139

There is always going to be a superscript.0143

This superscript is going to be also positive whole numbers--1, 2, 3, etc.0147

Basically this is equal to the number of electrons in that orbital; number of electrons in specific orbital.0156

What we are going to do now, we are going to use the periodic table to help us to determine an atom's electron configuration.0175

Basically when you look at the periodic table, at first it looks like one big mess, but it is incredibly organized.0182

There is so much information we can gather from this very important tool.0189

This first two columns, group 1 and group 2, this is what we call the s block; s block.0196

After the s block comes this whole area in green.0209

It is basically what we call the p block.0212

All of the transition metals are in yellow.0216

That is what we call the d block.0220

The lanthanides and actinides are at the bottom.0224

That is what we call the f block.0229

To read this, the n level, remember n comes first.0233

Basically n is basically just the row that the element is in.0237

n is equal to row number.0241

For example, when you look at hydrogen, you see 1s.0245

When you look at lithium, 2s.0249

When you look at sodium, it is 3s for example.0251

Another part, that s is actually helium.0255

You want to count helium as being right adjacent to hydrogen, just for electron configuration purposes.0259

The d block, this is where n changes.0266

Even though the d block doesn't start until the fourth row of the periodic table,0270

the first n value for the d block is going to be always one less than the row.0274

n is equal to one less than actual row.0279

In the fourth row, the first n value is actually 3.0285

Same thing for the f block; same thing for the f block.0292

It is going to be a little different.0295

You notice here that the f block, we don't start until n equal to 4.0296

But the f block actually starts in the sixth row.0302

The n is going to be two less than the actual row.0308

Basically now what you want to do is you want to count blocks you are in.0320

For example, we can go ahead and let's just start off with hydrogen.0326

Hydrogen has just the one electron; hydrogen is going to be 1s1.0332

When we move over to helium right here, that is going to be a total of two electrons.0343

Helium is going to be 1s2.0350

Lithium, lithium has a total of three electrons.0356

But remember we always start from the beginning; we always start from hydrogen.0360

Starting at the beginning, you get hydrogen 1s1; helium was 1s2.0366

Now lithium coming down through the next row is going to be 2s and just the one block.0372

Remember one block is equal to one electron.0378

When you see something like this, 1s1 1s2, you only count the maximum.0382

You only count the last one when they both have the same n value and the same letter.0387

Therefore I only count the last one; it is just 1s2 2s1 for lithium.0393

I know this may seem a little weird at first.0401

But let's go ahead and look at it more in detail now.0405

How do we fill orbitals?--the filling orbitals is...0407

We are going to base it off of what we call the Aufbau principle; the Aufbau principle.0411

Lithium is what we just looked at; lithium was again 1s2 2s1.0417

Let's go ahead and look at oxygen; oxygen is going to be 1s2 2s2.0427

Now oxygen is in the p block; it is four blocks over.0436

That is going to be 2p4; there is something you always want to check.0441

The superscripts remember tell me the number of electrons.0447

The sum of the superscripts should equal to the number of electrons in that atom.0453

That is how you should always check.0464

When we look at lithium, lithium is element number 3.0466

Indeed, 2 plus 1 gives us 3; that is good to go.0469

Remember how we only count the last one, using Aufbau principle.0473

1s1 1s2 was what we initially had; and then 2s1.0479

If I add those up, I get 4; we know that is not right.0487

That is why you only count the last one; that is oxygen.0491

Let's go ahead and do another example as chlorine.0498

You should follow along on the periodic table.0500

Chlorine, this is going to be... again you always start from the beginning.0502

Again only count the last one; 1s2 completes the first row.0506

2s2 is going to be lithium and beryllium.0511

2p6 completes the p block in the second row.0516

We go down to the next line.0520

It is just like reading sentences in a paragraph.0521

You go from top to bottom, left to right; top to bottom, left to right.0525

After 2p6 comes 3s2; finally now we are into the block for chlorine.0530

Chlorine is in the 3p block five blocks over, 3p5.0536

Again double check, do all of the superscripts add up to the atomic number of the element?0542

Indeed it does.0548

We are going to do electron configuration, two formats; the first way we just did.0551

The second way is to do what we call a noble gas notation.0557

This is going to save us time.0561

Basically what you want to look at is the preceding noble gas; look at preceding noble gas of the element.0563

If I look at lithium, the preceding noble gas for lithium is going to be helium.0577

Helium's configuration is 1s2.0584

If I look at oxygen, the preceding noble gas is also helium.0588

Helium is also represented by 1s2.0595

Finally for chlorine, chlorine's noble gas is going to be argon.0598

Argon is going to be represented... excuse me, not argon... neon.0605

Chlorine's preceding noble gas is neon; that is represented by this entire thing, 1s2 2s2 2p6.0613

The shortcut notation is just to put the noble gas in brackets like this.0623

Followed by all the other electrons to the immediate right; helium in brackets, 2s1.0629

Oxygen is going to be helium in brackets, 2s2 2p4.0636

Finally chlorine, chlorine is going to be neon in brackets and then 3s2 3p5.0644

We want to make a note then; what do these electrons mean?0655

Everything in brackets that is represented by the noble gas, think about this.0662

These are the lowest n values.0666

That means these are the innermost electrons, closest to the nucleus.0672

This is what we call the core electrons.0678

Everything else that is to the right of the bracket has the highest n value, which we just underlined right now.0681

We are going to give these electrons a very important name which is going to come into play later on.0690

These are what we call valence electrons.0695

Valence electrons are the electrons farthest from the nucleus.0699

They are the outermost electrons; outermost electrons.0702

A very important finding is that you can always tell the number of0710

valence electrons for main group elements just by looking at the periodic table.0715

For main group elements, the number of valence electrons equals simply to the column number.0721

Let's go ahead and think about this.0736

For lithium, the valence electron is 2s1.0738

There is only one of them; lithium is in column 1.0742

For oxygen, you have valence electrons of 2s22p4.0746

That is a grand total of six valence electrons; oxygen is in column 6.0751

Finally chlorine has valence electrons of 3s2 3p5; that is a grand total of seven.0756

Look where chlorine is--in column 7 of the periodic table.0763

There is another way of representing the electrons in an atom.0770

This is what we call orbital box diagrams.0779

Basically orbital box diagrams are just schematic depiction to show filling of atomic orbitals.0783

Two rules are essentially followed; very important rules.0792

Pauli exclusion principle, the Pauli exclusion principle states that an atomic orbital can hold a maximum of two electrons which must be opposite spin.0796

Once again the Pauli exclusion principles states that an atomic orbital can hold a maximum of two electrons which must be opposite spin.0807

Hund's rule states that a degenerate set of orbitals are to be filled singly with parallel spin.0818

What degenerate means is that they are going to have the same n value, the same principal energy level.0825

After each orbital is singly filled, you can then go ahead and insert a second electron but with opposite spin.0831

That is a lot to take in; let's go ahead and look at chlorine.0840

Chlorine again is 1s2 2s2 2p6 3s2 3p5.0843

What we are going to do is we are going to use a line or a box to represent each of these orbitals.0855

Remember an s orbital from the quantum mechanics presentation, there is only one per energy level.0860

A p orbital, there are three per energy level.0871

A d orbital means five per energy level.0877

Finally f was seven per energy; again this is only when applicable.0882

For example, a d orbital is not going to appear until n equal to 3.0890

F orbitals won't appear until n equal to 4.0898

That is what I mean by when applicable.0900

Again the periodic table is your best friend in this chapter, in this presentation.0902

You have to use the periodic table and you can see this easily.0906

We are going to go ahead and draw the orbital box diagram for chlorine.0910

That is it; one line represents the 1s orbital; 2s, also represented by one line.0915

2p, remember three per n; one, two, three; this is 1s2s; this is 2p.0923

3s is one line; finally 3p, three lines.0932

Again depending on what textbook you use, sometimes the lines will be instead drawn as boxes.0938

They are the same.0946

What we are going to do right now, we are going to fill electrons using Pauli exclusion principle and Hund's rule.0947

1s2, if I pair electrons, two electrons per orbital, they have to be of opposite spin.0954

That is what Hund's rule tells me.0965

If I put one as spin up, the second electron is going to be spin down.0968

We are representing electrons basically as arrows.0971

2s2, same thing; one spin up, one spin down.0975

The 2p orbital, you notice that there are three of them.0981

That is what we call a degenerate set; they have the same n value.0986

Hund's rule tells me I am going to fill these p orbitals singly first before doubling up.0991

I am going to do this with parallel spin.0996

For example, this is one electron of the six.0999

When I go to the next one and to the next one, that is 2p3 so far.1006

Again I have filled the atomic orbitals singly; now I can double up.1011

They are going to be all of opposite spin--2p4 2p5 and finally 2p6 in that order.1016

Yes it does matter how you write it; very important.1025

Go ahead now onto 3s2; spin up, spin down.1028

Now onto 3p5, once again one electron at a time per orbital before doubling up with opposite spin.1033

One, two, three, now four and five.1041

Once again we have applied Pauli exclusion principle and Hund's rule to what are called orbital box diagrams.1047

The Aufbau principle generally works.1059

There are of course a couple of exceptions that we need to account for.1062

Again you should ask your instructor which one you need to know.1066

But the two most common exemptions are in the first d row.1069

That is going to be copper and chromium.1072

Chromium we would expect to be [argon] 4s2 3d6; expected.1075

But actual configuration is going to be [argon] 4s1 3d5.1086

Copper, copper we would expect [argon] 4s2 3d9; expected.1096

But actual is going to be [argon] 4s1 3d10.1106

We are going to save the reasoning for perhaps a higher level course.1117

But for now, what you want to remember is these exceptions arise from experimental data.1126

The simple explanation is that the configuration is more stable due to what is called a half-filled or a completely filled orbital.1133

You see that this s orbital can hold two electrons.1144

It is holding only one right now; that is half-filled.1149

Here you have five of the d orbitals.1152

Each have two electrons, giving you a maximum of ten.1154

That is going to be a completely filled orbital.1157

This is half-filled; this is completely filled; half-filled and completely filled.1159

Something else to point out, this is going to be true not just for copper and chromium, but for transition metals in general.1172

It is to count the valence electrons.1179

Remember for main group elements, we go by the column number; that is it.1182

But for transition metals, they are not main group elements.1186

What rule do we go by for counting valence electrons?1189

Basically it is a simple rule.1192

You count all of the electrons that come after the noble gas for transition metals.1195

Chromium is going to have a total of six valence electrons.1201

Copper is going to have a total of eleven valence electrons.1207

Once again for transition metals, all electrons coming after the noble gas in brackets will be counted as valence.1212

Let's now take a look at the electron configuration for monatomic ions.1224

We are going to examine the configurations in orbital box diagrams.1230

Aluminum, aluminum is 1s2 2s2 2p6 3s2 and 3p1.1235

What does it mean to be aluminum 3+?1248

3+, that means that we remove three electrons.1252

Is it easier to remove electrons closest or farthest from the nucleus?1257

Electrons that are close to the nucleus, they are being drawn toward the nucleus because of that positive charge.1263

Electrons farther out from the nucleus are not held as tightly so they are easier to remove.1271

When you form cations, you remove valence electrons first.1279

In other words, remove the outermost electrons; 1s2, for aluminum 3+, 2s2 2p6.1288

The first electron I am going to remove is the 3p1 because that is farthest out, followed by 3s2.1304

Aluminum 3+'s configuration is 1s2 2s2 2p6.1312

Let's go ahead and repeat this now for chlorine.1317

Chlorine is 1s2 2s2 2p6 3s2 3p5.1321

Chlorine 1-, that 1- means we have added an additional electron.1330

Once again we are going to add it to the outermost orbital because that is most easily accessible.1335

You add electrons also to valence orbitals.1340

Chlorine 1- is 1s2 2s2 2p6 3s2; then 3p5, you add one more, becomes 3p6.1350

The reason why I wanted to bring this up is because there is something in common that these electron configurations of the ions share.1363

They all end in 2p6 which means they have a completely filled p block.1375

What elements are in that six p block column?--it is the noble gases.1381

This is all noble gas configuration; this is what we call noble gas configuration; noble gas configuration.1388

You can go ahead and check with the periodic table that all of them end in p6; all of the noble gases.1401

The valence electrons for a noble gas, two and six, two and six, which means eight valence electrons now.1409

This number 8 is going to come into play in the next presentation.1423

It is going to become very important.1427

But what I want you to take away from this section is the following.1429

That elements react to form ions and covalent bonds in order to achieve eight electrons.1437

Elements react to achieve eight valence electrons; again just take my word for now.1452

This number eight is going to become hugely important later on.1466

One application of knowing the electron configuration for an atom is we are going to be able to predict an atom's magnetic behavior.1475

Paramagnetism arises when the atom has at least one unpaired electron; one unpaired electron.1483

Why that is important is because paramagnetic species are going to be attracted to an external magnetic field.1491

In other words, they will stick to a magnet if you will.1497

Diamagnetism, the prefix di, you could probably tell what this means already.1501

Diamagnetism arises when an atom or ion contains no unpaired electrons.1505

In other words, all electrons are paired; all electrons are paired.1510

Diamagnetic species are going to be repelled by an external magnetic field.1517

An example of a paramagnetic species, we have a lot of them, maybe like 1s2 2s1.1524

Here we have the one unpaired electron; this will be for lithium.1535

Then let's go ahead and look at an example of a diamagnetic species.1540

Diamagnetic species will be just 1s and then 2s.1544

1s2 2s2, that is going to be for beryllium.1549

Zero unpaired electrons, that is what we mean.1556

Electron configuration was the first big periodic trend.1560

We saw how useful the periodic table came into play to help us predict configuration.1563

Let's now take a look at other periodic trends.1568

Atomic radius, atomic radius literally is the size of the atom.1573

It is the distance an atom's valence electrons are from its nucleus.1577

It is also known as atomic radius or atomic size; it turns out the following.1582

That in a column of the periodic table, atomic size increases top to bottom.1588

The reason is because of the following.1593

As you go down a column, you are increasing the n value.1595

As n increases, so does the size of the atomic orbital, remember the higher the energy.1600

In a row of the periodic table, atomic size decreases left to right.1606

Let's go over the reasoning for that.1613

Within the same row of the periodic table, you are pretty much keeping the same n value.1615

n does not change at all; what does change is the following.1620

That is going to be the number of protons inside the nucleus.1625

As the element gains more protons, its nuclear charge is going to increase1629

which essentially pulls the electrons in to a greater degree.1636

This results in contraction; once again atomic size is going to decrease left to right.1642

Ionic radii, when we look at ionic radii, remember that electrons are added to or taken away from the valence shell.1651

The outermost shell that is; that is very important.1659

For anions, because we are adding electrons to the outermost shell, anions typically are larger than the parent anion.1663

For example, O2- will be greater than just oxygen by itself.1672

Cations, we are removing electrons from the outermost shell.1677

Cations are typically going to be smaller than the parent ion.1682

For example, aluminum 3+ is going to be smaller than aluminum 0.1686

How about for what we call isoelectronic species?1694

Isoelectronic species means we have the same number of electrons.1697

For example, N3- versus O2-, you can look at their electron configurations.1700

They are going to have the same electron configurations, the same number of electrons.1709

The rule for isoelectronic species is just to go by nuclear charge.1712

The element basically with the greater nuclear charge, the greater number of protons,1718

is going to have a greater pull on the electrons, decreasing the size.1724

Therefore N3- which has fewer protons than oxygen is going to be greater in size than O2-.1730

Another periodic trend is what we call ionization energy; let's start off with the definition.1743

Ionization energy is the energy required to remove a valence electron in the gas phase.1750

Again you have to put in energy to overcome the attraction the electron has with the nucleus.1756

Removal of an electron is an endothermic process.1765

Typically then ionization energies are going to be positive values.1770

Think about ΔH in units of kilojoules per mole.1775

Ionization energy is completely opposite to atomic size.1782

Therefore smaller atoms have very large ionization energies; just think about that.1785

When we have smaller atoms, we have the electrons much closer, the valence electrons much much closer to the nucleus.1791

That attraction with the nucleus is incredibly strong.1799

You have to put in more energy to overcome that attraction.1802

For our purposes, fluorine is going to be the smallest of our atoms and therefore the largest ionization energy.1807

There is some additional terminology to get across.1825

An atom's first ionization energy refers to removing a valence electron from the ground state.1828

For example, oxygen goes on to form... excuse me.1834

For example, calcium can go on to form Ca1+ plus an electron when energy is applied.1849

We are just removing the first valence electron.1862

This is what is called the first ionization energy.1865

The second, third, etc, ionization energies then refer to subsequent removal of the electrons from a cation.1868

Let's go ahead and now take Ca1+.1876

Ca1+ now has its turn of being ionized to form Ca2+ plus an electron.1878

Once again this is what is called the first ionization energy.1887

This will be calcium's second ionization energy; very good.1891

The next periodic trend is what we call electron affinity.1899

Electron affinity refers to the energy that is released when an electron is added to an atom in the gas phase.1902

Which means addition of electrons is an exothermic process.1910

We will represent that as for example oxygen plus an electron goes on to form O1-.1915

We can also build on that; O1- plus an electron goes on to form O2-, etc.1922

Again this is what we call electron affinity.1929

Again it is going to be typically an exothermic process which means these are1932

going to be reported in less than 0, negative kilojoules per mole.1937

Electron affinity therefore parallels ionization energy.1945

We are going to find that nonmetals are going to be more apt to1951

gain electrons to form anions which will have the largest electron affinities.1955

What we want to do, we want summarize all of the periodic trends looking at the periodic table.1961

As you go left to right upward, this is a direction of increasing ionization energy and electron affinity.1969

However both of these are opposite to atomic size which increases this way.1982

This is increasing atomic size.1990

Really we use our start point and end point, fluorine being at one corner and then perhaps francium being at another corner.1998

Again fluorine being one of the smallest atoms is going to have the largest ionization energies, a very large electron affinity, etc.2010

These are our periodic trends; of course they come from consequences of quantum mechanics.2024

The periodic table we saw is very organized and informative concerning an element's physical and chemical properties.2033

An atom's electron configuration we learned, we do the filling of the orbitals based off of the Aufbau principle.2040

Finally orbital box diagrams are schematic representations of electrons following two important rules--the Pauli exclusion principle and Hund's rule.2049

Let's go ahead and tackle some sample problems.2063

For each of the following, give the ground state configuration and orbital box diagram.2066

Ground state simply refers to the neutral atom.2071

It is called ground state because it refers to the configuration that represents the lowest energy for the atom; of lowest energy.2078

Iron, the preceding noble gas for iron is going to be argon.2090

Then that is going to be 4s2 3d6.2098

Let's go ahead and draw it out.2102

I am just going to do the valence, the valence electrons; that is 4s.2105

Remember for these, how many d orbitals are there per energy level?2109

There is five of them; you have to draw them always as a degenerate set; degenerate set.2112

Let's go ahead and fill them; 4s2 first; now 3d6.2118

Remember one at a time with parallel spin before pairing up.2122

When you pair up, opposite spin.2127

Finally a phosphorus, phosphorus, you can go ahead and look at the preceding noble gas.2132

Preceding noble gas for phosphorus is going to be neon.2140

After neon is going to be 3s2... it is going to be 3s2 3p3.2143

Let's go ahead and draw that out, 3s.2154

Remember there is three degenerate orbitals for p's; let's go ahead and fill the electrons.2157

3s2 first; now 3p3; one, two, and three; let's just go ahead and complete them.2163

Iron is going to have a grand total of eight valence electrons.2174

Phosphorus is going to have a grand total of five valence electrons then.2179

We expect both of these, because they do have unpaired electrons, we expect both iron and phosphorus to be paramagnetic.2186

Let's go on to now sample problem number two.2200

For each of the following, circle the element and ion that has the larger indicated trait.2203

Basically we are going to use the periodic table all the way.2207

Ionization energy, when we look at sodium and lithium, they are in the same column.2210

Lithium is above sodium; sodium is going to have the smaller ionization energy.2215

Lithium should have the larger one; again the reason is because lithium is smaller.2222

It is going to require energy to remove its valence electron than in sodium.2233

Atomic size for oxygen or nitrogen, oxygen and nitrogen are in the same row.2245

Atomic size decreases left to right; oxygen is going to have the smaller size.2251

Nitrogen is going to have the larger size; remember why?2256

Because nitrogen is going to have fewer protons in its nucleus; so a smaller nuclear charge.2260

Because it is a smaller nuclear charge, electrons are going to be farther out from the nucleus.2270

Last one is atomic size, this time for O2- and N3-.2281

Again anytime you see ions, try to think isoelectronic species.2286

Again for isoelectronic species, you want to just go by the atomic number again; the number of protons.2292

Once again oxygen is going to have more protons.2300

Nitrogen is going to have fewer protons.2306

For the same reasoning as the previous question, N3- should be larger because of the fewer protons.2310

The attraction is less with the electrons; and they are farther out.2319

It was good to see everyone on again; thank you for your attention.2325