Trends in the Periodic Table

Summary

This video explores key trends in the periodic table, focusing on atomic radius, ionization energy, and electronegativity. Atomic radius increases top to bottom within a group and decreases left to right in a period. Ionization energy, the energy required to remove an electron, decreases down a group and increases across a period. Electronegativity, or an atom's desire for electrons, increases left to right in a period and decreases top to bottom. The video explains how these trends relate to atomic size and electron attraction, highlighting the importance of understanding why they occur.

Takeaways

• 🔍 A trend in the periodic table refers to a pattern in the properties of elements as you move either down a group or across a period.
• 📏 Atomic radius is the size of the atom, measured from the nucleus to the furthest electron in the electron cloud.
• ⬆️ Atomic radius increases as you move down a group because more electron shells are added, making the atoms larger.
• ➡️ Atomic radius decreases as you move from left to right across a period, due to an increased attraction between protons and electrons, pulling the electron cloud closer to the nucleus.
• ⚡ Ionization energy refers to the energy needed to remove an electron from an atom, with lower ionization energy indicating it's easier to remove an electron.
• ⬇️ Ionization energy decreases from top to bottom in a group because larger atoms have electrons further from the nucleus, making them easier to remove.
• ➡️ Ionization energy increases as you move from left to right across a period, due to smaller atomic size and a stronger hold on the electrons.
• 🧲 Electronegativity measures how much an atom wants an electron, with a higher value indicating a stronger attraction for electrons.
• ⬇️ Electronegativity decreases from top to bottom in a group because larger atoms have valence electrons farther from the nucleus, reducing their ability to attract electrons.
• ➡️ Electronegativity increases from left to right across a period as smaller atoms, like fluorine, have a stronger pull on electrons to complete their valence shell.

Q & A

What is a trend in the periodic table?

A trend in the periodic table refers to a pattern observed in certain properties of elements as we move either from top to bottom within a group or from left to right within a period.

What is atomic radius?

Atomic radius is the size of an atom, specifically the distance from the nucleus to the furthest electron in the electron cloud.

How does atomic radius change as you move from top to bottom in a group on the periodic table?

As you move from top to bottom in a group, the atomic radius increases due to the addition of more electron shells.

Why does the atomic radius decrease from left to right across a period?

The atomic radius decreases from left to right across a period because, as more protons are added to the nucleus, the increased nuclear charge pulls the electron cloud closer to the nucleus, shrinking the atom.

What is ionization energy?

Ionization energy is the amount of energy required to remove an electron from an atom.

How does ionization energy relate to atomic radius?

Ionization energy increases as atomic radius decreases, because smaller atoms hold their electrons more tightly due to the stronger attraction between the nucleus and the electrons.

How does ionization energy change across a period?

Ionization energy increases from left to right across a period because atoms become smaller, making it harder to remove an electron.

What is electronegativity?

Electronegativity is a measure of how much an atom attracts electrons. It is based on a scale from 0 to 4, where 4 indicates a strong desire for electrons.

Which element has the highest electronegativity and why?

Fluorine has the highest electronegativity because it is small, and its nucleus exerts a strong pull on neighboring electrons, making it highly attractive to electrons.

What is the general trend for electronegativity as you move down a group?

Electronegativity decreases as you move down a group because the atomic radius increases, and the nucleus is farther from the valence electrons, reducing the atom's ability to attract electrons.

Outlines

00:00

🔍 Introduction to Periodic Trends

The video introduces trends in the periodic table, which are patterns observed when examining properties of elements. Trends can be identified either from top to bottom within a group or from left to right within a period. Three key properties are discussed: atomic radius, ionization energy, and electronegativity. Atomic radius, the size of the atom, is explored first, with its pattern explained through electron shells and the snowman analogy for easier recall.

05:02

📏 Atomic Radius: Top to Bottom Trend

Atomic radius refers to the distance from the nucleus to the outermost electron. In Group 1 of the periodic table, from hydrogen to cesium, the atomic radius increases as more electron shells are added with each successive element. As we move from top to bottom in a group, more electron shells are needed to house additional electrons, which increases the atom's size. The snowman analogy is used to illustrate how atoms grow bigger down the group.

➡️ Atomic Radius: Left to Right Trend

Despite the addition of electrons across a period (left to right), atomic radius decreases. This happens because as we add more protons and electrons, the stronger attraction between them pulls the electron cloud closer to the nucleus. For example, in period two, lithium is the largest atom and neon is the smallest, as the electron cloud is drawn closer to the nucleus with increasing proton count.

⚡ Ionization Energy: Definition and Concept

Ionization energy is the amount of energy required to remove an electron from an atom. Atoms with a low ionization energy are easier to strip of electrons, while those with high ionization energy hold their electrons tightly. Smaller atoms, with electrons closer to the nucleus, have higher ionization energies. Larger atoms, with more electron shielding, have lower ionization energies since their outermost electrons are less attracted to the nucleus.

📈 Ionization Energy: Periodic Trends

Ionization energy increases from left to right across a period and decreases from top to bottom in a group. The reason for this is tied to atomic size—smaller atoms like neon hold onto their electrons more tightly, requiring more energy to remove one, while larger atoms like lithium have loosely bound electrons, making them easier to steal. In summary, the bigger the atom, the lower its ionization energy, and vice versa.

🔗 Electronegativity: Definition and Trends

Electronegativity measures how strongly an atom attracts electrons, with a scale ranging from 0 to 4. Fluorine has the highest electronegativity, while elements like lithium have much lower values. This difference is because fluorine, being small and just one electron short of a full valence shell, strongly attracts electrons. On the other hand, elements like lithium prefer to lose electrons to achieve a more stable configuration. Electronegativity increases from left to right across a period and decreases from top to bottom in a group.

📊 Final Thoughts on Periodic Trends

The video concludes by emphasizing the importance of understanding why periodic trends occur. All three properties—atomic radius, ionization energy, and electronegativity—are closely tied to the size of atoms and the distance between the nucleus and valence electrons. Understanding atomic size helps to explain the variations in these trends across the periodic table, making it easier to predict the behavior of elements.

Keywords

💡Atomic radius

Atomic radius refers to the size of an atom, measured as the distance from the nucleus to the outermost electron. In the video, it is discussed as a key trend in the periodic table. The atomic radius increases from top to bottom within a group due to the addition of electron shells and decreases from left to right across a period due to the stronger attraction between protons and electrons.

💡Ionization energy

Ionization energy is the amount of energy required to remove an electron from an atom. The video explains that smaller atoms have higher ionization energy because their electrons are closer to the nucleus and more tightly bound. The trend shows that ionization energy increases from left to right in a period and decreases from top to bottom in a group, mirroring the trend of atomic radius.

💡Electronegativity

Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. The video explains that smaller atoms with a higher nuclear charge, such as fluorine, have higher electronegativity. The trend shows that electronegativity increases from left to right across a period and decreases from top to bottom within a group, similar to ionization energy.

💡Periodic table trends

Trends in the periodic table refer to patterns observed in properties of elements as you move across periods (rows) or down groups (columns). The video highlights three major trends: atomic radius, ionization energy, and electronegativity. These trends are driven by the number of electron shells and the effective nuclear charge.

💡Electron shells

Electron shells are layers around the nucleus where electrons reside. The video explains that as you move from top to bottom in a group, the number of electron shells increases, resulting in a larger atomic radius. This plays a crucial role in explaining why atoms become larger as more electron shells are added.

💡Group

A group in the periodic table is a vertical column where elements share similar properties. The video discusses how trends such as increasing atomic radius, decreasing ionization energy, and decreasing electronegativity are observed as you move down a group, due to the addition of electron shells and increased distance between the nucleus and valence electrons.

💡Period

A period in the periodic table is a horizontal row of elements. The video highlights that as you move from left to right across a period, the atomic radius decreases, ionization energy increases, and electronegativity increases. These changes are due to the increasing number of protons, which pull the electron cloud closer to the nucleus.

💡Valence electrons

Valence electrons are the outermost electrons of an atom that are involved in chemical bonding. The video uses valence electrons to explain trends in ionization energy and electronegativity, noting that atoms with fewer valence electrons, like lithium, have lower electronegativity and ionization energy compared to elements like fluorine.

💡Shielding

Shielding refers to the effect where inner electron shells block the attraction between the nucleus and the outermost (valence) electrons. The video explains that in larger atoms, increased shielding reduces the pull of the nucleus on valence electrons, leading to lower ionization energy and electronegativity as you move down a group.

💡Nucleus

The nucleus is the central core of an atom, containing protons and neutrons. The video frequently references the nucleus when discussing atomic radius, ionization energy, and electronegativity. The attraction between the nucleus and electrons determines the size of the atom and how strongly the atom holds onto or attracts electrons.

Highlights

A trend is a pattern in the periodic table, observed from top to bottom within a group or from left to right within a period.

Atomic radius is the size of the atom, measured from the nucleus to the outermost electron.

As you move from top to bottom in a group, atomic radius increases due to more electron shells being added.

As you move from left to right in a period, atomic radius decreases because the attraction between the nucleus and electrons increases.

The atomic radius trend: largest atoms are found in the bottom-left of the periodic table (Francium), while the smallest are in the top-right (Helium).

Ionization energy is the energy needed to remove an electron from an atom, with smaller atoms requiring more energy.

Ionization energy increases from left to right across a period because smaller atoms hold onto their electrons more tightly.

Ionization energy decreases from top to bottom in a group, as larger atoms have electrons further from the nucleus, making them easier to remove.

Electronegativity measures how much an atom wants an electron, with a scale from 0 to 4.

Fluorine has the highest electronegativity, as it strongly desires one more electron to complete its valence shell.

Elements on the left side of the periodic table, like lithium, have lower electronegativity because they prefer to lose electrons.

Electronegativity increases from left to right across a period, as atoms on the right strongly attract electrons to fill their valence shells.

Electronegativity decreases from top to bottom within a group, as larger atoms find it harder to attract electrons due to their size.

Atomic radius, ionization energy, and electronegativity are interconnected, often depending on the size of the atom and distance between the nucleus and valence electrons.

Understanding these trends allows prediction of element behaviors across the periodic table, including how easily atoms lose or gain electrons.

Transcripts

hi in this video we're looking at trends

in the periodic table and a trend is

just a pattern so if we look in the

periodic table from top to bottom within

a group we might see a pattern for a

certain property similarly if we look

left to right within a period in the

periodic table we might also see a trend

in a certain property and so we're gonna

look at three properties the first

property is something called atomic

radius

the second is ionization energy and the

third is electronegativity so let's

start with just what is atomic radius

atomic radius is the size of the atom

it's actually physically the radius of

the atom itself the distance from the

nucleus to the furthest out electron in

the electron cloud and so if I just look

at Group one here from hydrogen to

cesium and I look at how many electron

shells are occupied by electrons in

hydrogen I really only see that there's

one electron shell because there's only

one electron so we only need one shell

to house that one electron if I look at

lithium though lithium has three total

electrons that means two of those

electrons can fit in the first shell but

the third will have to go into the

second shell so lithium has two shells

and as you kind of go from top to bottom

in this group we see that we just have

an increasing number of electron shells

because the electron amount is going up

and we need more room to house those

electrons the way that we house those

electrons in an increasing amount is by

just bringing in more electron shells to

hold them and what you can see is that

as we go top to bottom there is a bigger

radius because we have more and more

electron shells the way to remember this

just quickly is to think of a snowman if

you look at these bottom three here you

see a snowman that was always my trick

for remembering this and I'm sure it's

plenty of other people's too it can be

yours but the reason that that's

happening is because there are more and

more electron shells as we go from one

period to the next down a group let me

lay out the atomic radius for all of the

elements on the periodic table what we

see is that top to bottom trend of the

atomic radius increasing but what you'll

also see is that as we go from left to

right within a period the atomic radius

actually decreases now

may seem strange to you because if we're

adding electrons as we go top to bottom

and that's causing our radius to get

bigger shouldn't our radius get bigger

as we go from left to right in a period

because in that situation we're also

adding electrons the explanation for

this one is a little more complicated so

let's just take a look at the elements

in period two lithium to neon lithium

has three protons and three electrons as

I slide over from lithium to neon I'm

adding one more proton but I'm also

adding one more electron and as that

happens the attraction of the electrons

to the nucleus becomes greater and

greater and that causes the entire

electron cloud to just kind of scoot you

in closer to the nucleus so you can see

of this grouping here neon is actually

the smallest atom and lithium lithium

lithium is the biggest atom in period

two so now we have our full trend here

top to bottom in a in a group the atomic

radius is going to increase think of the

snowman effect there the radius actually

decreases as we go from left to right

within a period so the overall trend if

you think of it this way this may help

you helium is the smallest atom and

francium or a caesium that would be on

this slide here but francium is the

largest atom and so knowing this trend

you can actually get a lot of the

information for the other two properties

we're looking at ionization energy as

the next property we'll investigate

ionization energy has a fancy name but

it actually is not all that fancy when

you know what it is ionization energy is

the amount of energy needed to remove an

electron from an atom you can think of

this almost like it's the cost of taking

an electron from an atom a low

ionization energy means it's really easy

to take an electron from that atom a

high ionization energy means it's more

difficult now here's the overall idea

the smaller and atom is the harder it is

to remove an electron from it if I just

look at an extreme case where I have a

small atom the valence electrons are

really close to the nucleus in this

scenario but if I have a larger atom the

valence electrons are much

farther away from the nucleus and

therefore they're much more willing to

kind of be donated to somebody else

another effect that's going on here is

the fact that we have all of these rings

that contain electrons already and so

this is called shielding shielding is

where the interior electrons kind of

mitigate the attractive force of the

nucleus the outside valence electrons so

said differently the signal is kind of

being blocked by the middle electrons

and so the valence electrons are not as

loyal to the atom and so it's much

easier to steal from a bigger atom as

compared to a smaller one now if we know

the trend in atomic radius in the

periodic table and we know that the

bigger the atom is the easier it is to

steal from it and therefore the lower

lianas ation and energy is required then

we can figure out what the trend in

ionization energy would be just based on

knowing the size of the atoms across the

periodic table the ionization energy is

going to increase from left to right

because if you just compare the size of

lithium for example to the size of neon

this is much more difficult so higher

energy over here because the size is

smaller those electrons are being held

really tightly to that nucleus lithium

not as much as neon so that means lower

energy over this way and remember the

lower the energy requirement the easier

it is to take an electron from that

particular element if we know that from

left to right across a period and it's

really just based off of the size then

we can also figure out what the pattern

is from top to bottom and a group for

ionization energy it's at the ionization

energy energy decreases so again the

bigger the atom the easier it is to

steal from it and that means the lower

the ionization energy will be okay so

the last piece is going to be

electronegativity electronegativity is

going to come back when we talk about

bonds and bond types so this is

something that is certainly important to

know it's going to help you when we when

we talk about bonds but in an

electronegativity

what we're talking about is how much an

atom wants an electron so it's a measure

of an atoms

action two electrons and it's based on a

scale from zero to four where four is

the element that really really wants

electrons and zero is an element that

really really does not want electrons so

let's look back at the periodic table we

fill in all the atomic radii for each of

the elements here I'd actually just like

to look at kind of the two extreme

scenarios here let's let's compare

lithium to fluorine fluorine has seven

valence electrons

if fluorine can just gain one more

valence electron it'd be a lot like neon

and that would fill its valence shell it

would satisfy the octet rule and

fluorine would absolutely love to gain

an electron there four fluorines

electronegativity is going to be really

high in fact fluorine has the highest

electronegativity of all the elements on

the periodic table now contrast that

against lithium they're in the same

period but lithium is on the other side

of the periodic table the left side

lithium has one valence electron and

does not want to gain any more valence

electrons in fact lithium would actually

prefer to lose that one valence electron

to become a lot like helium and so the

electronegativity for lithium and for

all the elements on sort of the left

side of the periodic table is going to

be far lowers let's take a look at it

just a special periodic table that only

shows the electronegativity values if

I'm yellow I really don't want electrons

in this periodic table and if I'm red

that means I really do again let's look

at fluorine three point nine eight

compared to lithium's point nine eight

there it is

and so if I just kind of look from left

to right within a period I'm gonna see

that the electronegativity increases if

I go top to bottom in a group the

electronegativity actually decreases and

this again unsurprisingly

has a lot to do with the radius of the

atom if I just kind of look kind of down

this way here all of these halogens

would like to gain one electron to be

like their noble gas neighbors it is

much easier for fluorine to attract an

electron that isn't originally its own

because fluorine is so small its nucleus

is located very close to its valence

shell as you go from top to bottom the

nucleus

gets further and further tucked in the

center the valence electrons get kind of

further and further out so the

electronegativity actually is a little

bit lower as we go top to bottom not

because they don't want that one

electron but because they're pull on

neighboring electrons is a little lower

because they're just bigger atoms and so

that's what's going on top to bottom in

a group on the periodic table for

electronegativity so that's it those are

the three trends in the periodic table

it's important that you know what the

trends are for sure but I believe it is

way more important that you know why

these trends exist and if you think

about it long enough it really all boils

down to the size of the atoms in many

cases it's about the distance between

the nucleus and the valence electrons

thank you

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