3.1 Atomic Theory and Atomic Structure | High School Chemistry

Summary

This chemistry lesson delves into atomic theory and structure, highlighting the pioneering work of John Dalton and others. It explains the four tenets of modern atomic theory, the discovery of the electron, and the atomic nucleus. The script introduces subatomic particles, isotopes, and atomic weights, emphasizing the concept of isotopes as different forms of an element with varying neutron counts. It also clarifies the distinction between atomic mass numbers and atomic weights, illustrating how atomic weights are calculated as weighted averages of naturally occurring isotopes.

Takeaways

• 🌟 Atomic theory has evolved through various stages, with modern atomic theory having four key tenets: all matter is composed of atoms, elements are defined by their atomic number, atoms of the same element are identical, and atoms of different elements are distinct.
• 👨‍🔬 John Dalton is recognized as the father of modern atomic theory, pioneering the foundational concepts of atomic structure.
• 🔬 J.J. Thomson discovered the electron and determined the mass-to-charge ratio, while Robert Millikan's oil drop experiment established the charge of an electron and indirectly its mass.
• 💥 Ernest Rutherford's gold foil experiment led to the discovery of the atomic nucleus, showing that atoms are mostly empty space with concentrated areas of mass and charge.
• ⚛️ Atoms consist of three subatomic particles: protons, neutrons, and electrons, with protons and electrons being charged and neutrons being neutral.
• 🔢 The atomic number defines the element and the number of protons in the nucleus, while the mass number represents the total number of protons and neutrons.
• ⚖️ Atomic mass units (amu) are used to express the mass of subatomic particles, with protons and neutrons each roughly weighing 1 amu and electrons much less.
• 🌐 The nucleus of an atom contains the majority of its mass, with electrons orbiting at a significant distance, creating a large electron cloud.
• 📊 Isotopes are variants of an element with different numbers of neutrons, resulting in different mass numbers but the same atomic number.
• 🔝 The atomic weight or mass listed on the periodic table is an average of the masses of all naturally occurring isotopes of an element.
• 📚 Memorizing the first 30 elements of the periodic table and their symbols is recommended, as it provides a solid foundation for further study in chemistry.

Q & A

What is the main topic of this lesson?

The main topic of this lesson is atomic theory and atomic structure, which is part of a high school chemistry playlist.

What are the four tenets of modern atomic theory?

The four tenets of modern atomic theory are: 1) All matter is composed of atoms, 2) All atoms of an element are the same, and different elements have different atoms, 3) Atoms of one element cannot be converted into atoms of a different element through normal chemical reactions, and 4) Atoms of different elements combine in whole number ratios to form compounds.

Who is credited with pioneering modern atomic theory?

John Dalton is credited with pioneering modern atomic theory.

What did J.J. Thomson discover that contributed to the understanding of atomic structure?

J.J. Thomson discovered the electron and the cathode ray, determining the mass-to-charge ratio of electrons.

What experiment did Robert Millikan conduct to determine the charge of an electron?

Robert Millikan conducted the oil drop experiment, which helped him determine the charge of an electron by observing the behavior of charged oil droplets.

What did Ernest Rutherford discover through his gold foil experiment?

Ernest Rutherford discovered the nucleus of an atom by observing the scattering of alpha particles through a thin gold foil.

What are the three subatomic particles and their charges?

The three subatomic particles are protons (positively charged), neutrons (no charge), and electrons (negatively charged).

What is the significance of the atomic mass unit (amu)?

The atomic mass unit (amu) is a convenient unit of mass for expressing the mass of atoms and their constituents, as their actual masses in grams or kilograms are extremely small.

How can the mass of an atom be compared to the mass of its nucleus?

The mass of an atom is mostly concentrated in its nucleus, which contains protons and neutrons. Electrons have a much smaller mass compared to protons and neutrons, so they contribute very little to the overall mass of an atom.

What is an isotope symbol and what does it represent?

An isotope symbol represents a specific isotope of an element, showing the element symbol, the mass number (total number of protons and neutrons), and the atomic number (number of protons).

What is the difference between an atom and an ion?

An atom is a neutral species with an equal number of protons and electrons. An ion is a charged species that has gained or lost electrons, resulting in an imbalance between the number of protons and electrons.

Why are the atomic weights listed on the periodic table averages?

The atomic weights on the periodic table are averages because they represent the weighted average of all naturally occurring isotopes of an element, taking into account their relative abundances.

How are atomic weights calculated for elements with multiple isotopes?

Atomic weights are calculated using a weighted average, where the mass of each isotope is multiplied by its natural abundance (expressed as a decimal), and the results are summed to get the average atomic mass.

Outlines

00:00

🔬 Introduction to Atomic Theory and Structure

This paragraph introduces the topic of atomic theory and structure, which is the focus of the lesson. It is part of a high school chemistry playlist, with the intention to cover atoms, molecules, and ions. The instructor emphasizes the development of atomic theory over time, highlighting John Dalton as a key figure. The modern atomic theory is outlined with four main tenets: 1) All matter is composed of atoms, 2) Atoms of the same element are identical, and atoms of different elements are distinct, 3) Atoms of one element cannot be converted into atoms of another through normal chemical reactions, and 4) Atoms combine in whole number ratios to form compounds. The lesson promises to delve into the history of atomic theory with notable scientists like J.J. Thomson and Robert Millikan, and to explore the structure of atoms, including subatomic particles such as protons, neutrons, and electrons.

05:02

🌌 Understanding Atomic Structure and the Concept of Empty Space

This paragraph delves into the atomic structure, explaining the roles of subatomic particles: protons, neutrons, and electrons. It describes the nucleus, composed of protons and neutrons, as the center of an atom where most of the mass is concentrated, while electrons orbit the nucleus, forming an electron cloud. The instructor uses the analogy of a marble-sized nucleus with an electron cloud spanning half a mile in diameter to illustrate the vast empty space within an atom. The concept that most of the universe, including atoms, is composed of empty space is emphasized, challenging the common perception of solidity in matter.

10:05

📊 Isotope Symbols and Atomic Numbers

The focus shifts to isotope symbols, using oxygen as an example to explain atomic numbers and mass numbers. The atomic number, which is the number of protons in an atom's nucleus, determines the element. The mass number represents the total number of protons and neutrons. The instructor clarifies that the atomic number can be omitted from an isotope symbol since it is inherent in the element's identity. The concept of isotopes, which are different forms of the same element with varying numbers of neutrons, is introduced, with oxygen-16 and oxygen-18 as examples.

15:05

⚛️ Atomic Weights and the Existence of Ions

This paragraph discusses the concept of atomic weights, which are averages of the masses of all naturally occurring isotopes of an element. It explains that atoms can gain or lose electrons, resulting in ions with either a positive or negative charge. The atomic weight is derived from the weighted average of the masses of the isotopes based on their natural abundance. The instructor provides a step-by-step example using chlorine isotopes to demonstrate how atomic weights are calculated.

20:08

📐 Conclusion on Atomic Theory and Structure

The final paragraph wraps up the lesson on atomic theory and structure, summarizing the key points covered in the video. It includes an invitation for viewers to like, share, and support the channel, as well as a mention of additional resources available on chatsprep.com for further study and practice materials. The instructor reiterates the importance of understanding atomic theory and structure as foundational knowledge in chemistry.

Keywords

💡Atomic Theory

Atomic theory refers to the concept that matter is composed of atoms, which are the smallest units of matter that retain their properties. In the video, atomic theory is foundational to understanding the structure of matter and is attributed to John Dalton, who is considered the father of modern atomic theory. The script discusses the four tenets of modern atomic theory, emphasizing that atoms of different elements combine in whole number ratios to form compounds.

💡Subatomic Particles

Subatomic particles are smaller components that make up an atom. The video identifies three main subatomic particles: protons, neutrons, and electrons. Protons carry a positive charge, neutrons are neutral, and electrons have a negative charge. These particles are crucial for understanding atomic structure and are central to the script's explanation of how atoms are composed.

💡Isotopes

Isotopes are variants of a particular chemical element which differ in neutron number, and consequently in nucleon number. The video explains isotopes by discussing oxygen-16 and oxygen-18, highlighting that isotopes of the same element have the same number of protons but different numbers of neutrons, resulting in different mass numbers.

💡Periodic Table

The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configuration, and recurring chemical properties. The script mentions the periodic table in the context of organizing elements by atomic number and introduces the concept of element symbols, such as 'O' for oxygen.

💡Electron

Electrons are subatomic particles found outside the nucleus of an atom, carrying a negative electric charge. The video describes electrons as being much lighter than protons or neutrons and as crucial for forming chemical bonds due to their attraction to the positively charged protons in the nucleus.

💡Proton

Protons are subatomic particles with a positive electric charge found in the nucleus of an atom. The video explains that the number of protons in an atom's nucleus determines the element's identity, and that protons and neutrons have approximately equal mass.

💡Neutron

Neutrons are subatomic particles with no electric charge, also located in the nucleus of an atom. The video clarifies that the mass of a neutron is similar to that of a proton and that the number of neutrons, in addition to protons, determines the isotope of an element.

💡Atomic Mass Unit (AMU)

The atomic mass unit is a unit of mass that is used to express the mass of atoms and molecules. The video introduces AMU as a convenient unit for discussing the mass of subatomic particles, explaining that one proton or neutron is roughly equivalent to one AMU, while an electron is only a fraction of an AMU.

💡Nucleus

The nucleus is the central part of an atom, containing protons and neutrons. The video describes the nucleus as a region of concentrated mass and positive charge, which is surrounded by a cloud of electrons, emphasizing that most of an atom's mass is found in the nucleus.

💡Ion

An ion is an atom or molecule with a net electric charge due to the loss or gain of one or more electrons. The video explains that ions are formed when atoms gain or lose electrons, resulting in a charge. For example, an oxygen atom with a negative two charge would be an ion with ten electrons, as opposed to the eight electrons found in a neutral oxygen atom.

💡Atomic Weight

Atomic weight, also known as atomic mass, is a measure of how much mass is in a single atom of an element, on average. The video discusses how atomic weights are calculated as weighted averages of the masses of an element's naturally occurring isotopes, using chlorine as an example to illustrate the calculation process.

Highlights

Introduction to atomic theory and atomic structure in the context of high school chemistry.

Atomic theory has evolved through several iterations, with modern atomic theory being the focus of this lesson.

John Dalton is credited with pioneering modern atomic theory.

Four tenets of modern atomic theory are discussed, including the indivisibility of atoms and the uniformity of atoms within an element.

Atoms of different elements cannot be converted into one another through normal chemical reactions.

Atoms combine in whole number ratios to form compounds, as exemplified by the formula for water, H2O.

Historical figures like J.J. Thomson and Robert Millikan contributed to the understanding of atomic structure, including the discovery of the electron.

Ernest Rutherford's gold foil experiment led to the discovery of the atomic nucleus.

Subatomic particles—protons, neutrons, and electrons—are introduced as the building blocks of atoms.

The atomic nucleus, composed of protons and neutrons, is contrasted with the electron cloud surrounding it.

The atomic mass unit (amu) is explained as a convenient unit for expressing the mass of subatomic particles.

Electrons are much lighter than protons and neutrons, contributing to the atom's mass being primarily in the nucleus.

The concept of isotopes is introduced, with oxygen 16 and oxygen 18 as examples.

Isotope symbols and their components, such as atomic number and mass number, are explained.

The difference between neutral atoms and ions is clarified, with the gain or loss of electrons as the distinguishing factor.

Atomic weight or atomic mass is discussed as an average of the masses of all naturally occurring isotopes of an element.

The calculation of atomic weights using weighted averages is demonstrated with the example of chlorine.

The lesson concludes with a summary of key points and an invitation to access additional study materials.

Transcripts

atomic theory and atomic structure that

is going to be the

topic of interest here in this lesson uh

this is the third

chapter in my high school chemistry

playlist uh on atoms molecules and ions

we'll follow up this lesson

with an introduction to the periodic

table and then finish this chapter off

with

nomenclature which is just a fancy way

of saying naming

compounds now this is my brand new high

school chemistry playlist i'll be

releasing these lessons

weekly throughout the 2020-21 school

year so if you don't want to miss one

subscribe to the channel click the bell

notifications

you'll be notified every time i put one

of these up

so atomic theory is the first thing i'm

going to tackle here and atomic theory

uh

went through several iterations to kind

of get where it was so sometimes you'll

hear people talk about

modern atomic theory uh because some of

the earlier things had

slightly different uh views and stuff

like that it kind of went through a

process to get where it is and

john dalton's responsible for pioneering

this he's one of the

the big people in history you're going

to want to know for this chapter so

but they're really four tenets to modern

atomic theory and you kind of got to

know them all

and the first part is just that all of

matter is composed of these little atoms

that's the first part

so cool second part deals with elements

on the periodic table and then

all of the atoms of an element are the

same and then all of the atoms of

all the other elements then are

different than the one you're looking at

so so if you look at all the different

elements on the periodic table that's

the number of different types of

atoms that exist so carbon atoms are the

same as other carbon atoms but carbon

atoms are different than

oxygen atoms essentially now third tenet

of atomic theory says that you can't

convert the atoms of

one element into atoms of a different

type of element

at least not my normal chemical

reactions so it turns out there are

nuclear decay processes like nuclear

radiation

so where some of this actually does

actually happen but my normal chemical

means it's not possible there used to

be alchemists back uh several hundred

years who used to try and convert

lead into gold and things of his sort uh

and by normal chemical means it was

never going to happen all right finally

the fourth tenet

of atomic theory is that the the atoms

of the different elements are going to

combine

in different whole number ratios to form

compounds that's where compounds come

from

so if we look at like the formula for

water water

is h2o so and what this means is that

so to form water you're going to take

two hydrogen atoms for every one oxygen

atom so

whether you have a small sample of water

or a big sample of water

if you actually split apart all the

atoms and counted them up you would have

twice as many hydrogen atoms as

oxygen atoms they mix in a two to one

ratio cool so that's atomic theory and

as i said before this again was

pioneered by john dalton and i don't

care if

you really know their first names but

probably they're four guys really and

possibly a fifth but we'll go with four

uh names you should probably know for

the history kind of pioneering

atomic structure here so when john

dalton's the first and you should just

know that yep

he's kind of like the father of modern

atomic theory he pioneered atomic theory

came up with his version of atomic

theory

good to know for him so next guy on the

list is thompson that's

j.j thompson and jj thompson discovered

the electron and

uh it turns out he discovered what's

called a cathode tube which is just a

beam of electrons and

they made these cathode rays but long

before they ever knew they were actually

a beam of

electrons so but he found out they were

a beam of charged

particles so and things of this sort but

he found out the

the mass to charge ratio of electrons

but he didn't know the mass or the

charge individually just that ratio

so but he's credited largely with

discovering the electrons that's what

you'd

kind of take away with with jj

thompson's name uh robert milliken is

the next on the list and

milliken took thompson's work a step

further and he actually figured out the

charge on an electron and

he's got his famous we called it the oil

drop experiment

all he did was take some some small

little micro droplets droplets of oil

and he charged them with electrons and

what he found is that

every single one of these droplets had a

multiple of a certain charge and he

figured

well that smallest charge these are all

multiples of that that are on these

little droplets

that must be the fundamental charge and

turns out that's the charge of an

electron

so it is the smallest you know unit of

charge that's

possible and when something's you know

either positively or negatively charged

it is

always a multiple of the charge on an

electron so

and because thompson had already figured

out the mass to charge ratio once he

figured out the actual charge on

electron

he was then able to use thompson's ratio

and figure out the mass of an electron

as well

so but take that with you that milliken

figured out therefore both the charge

and then by default or or indirectly the

mass of an electron

uh and then finally rutherford's the

next guy and rutherford took a

thin thin super thin piece of gold foil

and he shined alpha particles which is a

form of nuclear radiation at this gold

foil so and when he shined the alpha

particles

at this gold foil interesting thing

enough is most of them just

passed right through so and this was you

know kind of a

unique discovery but every once in a

while some of them would get scattered

in different directions or even bounce

back in some cases

and so what he presumed then is that

this gold foil was made up of mostly

empty space so with then concentrated

areas of both mass and charge

and we now know that those concentrated

areas of mass and charge

are the nucleus of an atom and so

but big thing you should know is that

ernest rutherford discovered the nucleus

with his shining alpha particles through

thin gold foil

so that's kind of the basis these four

guys you should know their

last name and associate them just a

little bit with what they did so i kind

of tried to give you the bare minimum

on the study guide there cool now i gave

you also a little model of the atom

there so because we want to talk about

the structure of an atom so now that we

know that all of matter is made of atoms

now we want to talk about what's the

atom made of

well there's really three what we call

subatomic particles and

subatomic meaning smaller than the atom

so and that's the proton the neutron and

the electron

so let's take a look at these for a

second so we've got the proton which we

symbolize with the letter p

the neutron with the letter n and then

the electron

with the letter e and oftentimes we'll

put a little e negative on it so

it turns out that two of these particles

are charged so it turns out that protons

are

positively charged electrons are

negatively charged and they're attracted

to each other

so and then neutrons are neutral they

have no charge which is

you know the source of all sorts of

stupid jokes out there right you know

neutron walks into the bar orders a

drink

says how much will it be and the

bartender says for you

no charge yeah it's terrible so i told

you in the first lesson that this

wouldn't be

that it wouldn't be the last lame joke

you heard and now i'm making good on

that

all right so if we take a look at the

protons neutrons and electrons and

what's really going on inside of an atom

so at the center of an atom so

this is where you get protons and

neutrons this

is your nucleus and then going around

this nucleus

are going to be some electrons so

and we call it an electron cloud and

they're just kind of moving around and

stuff like this and so

the protons are positive so and then

therefore these electrons they stay

associated with atoms just

based on the attraction to the protons

in the nucleus so they're attracted to

the nucleus

just that plus minus attraction cool

now it turns out that protons and

neutrons roughly weigh about the same

amount

and if we tried to weigh that in like

grams or kilograms it would be a

ridiculously

tiny number because these are really

really small so what they did is

invented a whole new unit they called it

the atomic

mass unit or amu for an abbreviation

uh just specifically for talking about

this stuff it's kind of like saying

hey chad how many miles tall are you

well that's not really a convenient unit

to get my height

so but if you said hey chad how many

inches tall are you we could go there we

could go like 68.

so things of his source so uh it's just

picking a more convenient units talk

about the mass of these guys because if

we talked about the

the mass and grams it'd be like you know

1.67 times 10 to the negative 24 grams

or 10 to the negative 27 kilograms or

it'd be stupid you know small numbers so

what they said we'll just make this easy

roughly one amu for the proton roughly

one

amu for the neutron so but it turns out

if you're really

one eighteen hundred and twenty second

of an amu for an electron so a super

small fraction here

so electrons are far tinier than the

protons and neutrons

and so as a result then almost all of

the mass of an atom is concentrated in

that nucleus right at the center

with these little teeny tiny electrons

floating around now to just kind of give

you some scope here some scale here

so think about this this is one eighteen

hundred and twenty seconds wait for the

electron

of the proton or neutron this is kind of

like comparing a small

car to a baseball these electrons are

way

tinier than the protons and neutrons and

here's where things are going to get a

little

crazy and mind-boggling if you kind of

look at you know how big the nucleus is

compared to how big the whole atom is

that's where you're going to see that

the scale drawing in your study guide

there

is way off so if you say let's say we

had the nucleus here which again is

super tiny but let's just say we had it

and

we built a scale model and we used a

marble for the nucleus

just a regular standard marble so what

you'd find is that the electron cloud

around that nucleus would probably be

about half a mile

in diameter so a little marble nucleus

where almost all the mass is

concentrated in the atom

and then half a mile for the electron

cloud

with these super teeny tiny electrons

going around

so pretty crazy and so this kind of goes

hand in hand with what rutherford

discovered

is that you know most of an atom is made

of empty space

and you know every once in a blue moon

you get some concentrated mass

and charge when you hit that nucleus but

most of an atom is made of empty space

now this is kind of mind-boggling

because you know things feel

solid you know but they're not they're

made of mostly empty space so

think about this for a second if you

leave the earth's atmosphere and go out

into

space now they call it space because

it's mostly

empty space now the truth is even in

empty space

it's not completely empty every once in

a blue moon you're going to like bump

into a hydrogen atom

but not very often so it's almost

completely empty space

but then even when you come you know the

surface of the earth

and you start touching matter and stuff

like that even that

matter is made of mostly empty space

most of the universe whether you're out

in space or on a planet is made of most

mostly empty space it is mind-boggling

so if i asked you what is most of this

you know what is most of the stuff in

between my ears made up of you should

say

empty space that's right cool so most of

the universe made of empty space which

you just learned

all right so uh cool

so now we know what a nucleus is made up

of we want to move on to talking about

isotope symbols and we're going to

specifically look at

oxygen here so in oxygen 16

8 we'll put a minus 2 up there we'll get

rid of these guys

all right so all those protons neutrons

electrons we need to account for them

and this is called an isotope symbol

right here and specifically

the o here stands for oxygen so we've

got a whole bunch of symbols on the

periodic table and we're going to

do a little more formal introduction of

the periodic table in a minute but o

stands for oxygen

so most of the time teachers are not

going to make you memorize all of the

elements of the periodic table but it

does on occasion happen

so however i will let you know in

college it almost never happens

so because we know that you're going to

have the internet and you're going to

have a periodic table on you at all

times you know

you'll have a periodic table on your

sock in some cases i've seen skirts i've

seen shoes

i've seen ties i've seen tattoos

it's crazy there's pair of tables

everywhere it's on the front cover of

your book

you know so it's on the wall in every

chemistry room at most universities so

you're always going to have access to a

periodic table and so memorizing

you know the entire periodic table is

pretty much a waste of time but

it is a requirement on students every

once in a while and if that's the case

for you and your class i apologize

so but what i do recommend is this

probably

just memorize the first 30. and when i

say memorize the first 30 i don't even

memorize like their atomic numbers or

anything like that

i just want you to know that zn stands

for zinc know what the symbol what

element it actually stands for

so o is for auction f is for fluorine b

is for boron

that kind of thing up to the first 30.

just you can work with them and most of

the time it's not so bad right

c is for carbon usually the letter

corresponds to the element right

so there's a couple exceptions right you

know you go to fe right here and

fe is iron so and it turns out they use

the latin root fair

so for that instead and for some of

these they end up using a latin root i

apologize

so this would be a little more

challenging but most of the time knowing

that zn is inc

not so bad knowing that ni is nickel

again kind of makes sense

but cool associate the names of the

symbols for at least the first

30. any of the rest you need to know

you'll kind of get familiarized as you

go through the process

so but highly recommend there all right

so if we take a look at this

isotope symbol so we call this

bottom number here the atomic number

and it always gives you the number of

protons every single time

so when i see this i know that we've got

eight

protons now the truth is this though

they don't actually have to give you the

atomic number they can erase it

and there you go so the idea is that if

you look the entire periodic table is

organized by atomic number

and oxygen by definition that's the

number in blue here

is atomic number eight so the number of

protons determines what elements you

have

so however many new you know protons are

in the nucleus that

is going to identify the element for you

so if it's eight protons the nucleus

that's oxygen

period done and so in this case because

oxygen is atomic number eight is

associated with having eight protons

it doesn't actually have to be given

here so but i put it on there just so we

could actually talk about it

so but that's the one number they can

leave off of an isotope symbol

now the top one here we call the mass

number

so and that mass number is kind of like

a rounded

whole number and if you recall i erased

it

protons weighed about 1 amu neutrons

weighed about 1 amu

and electrons weigh a super duper duper

duper tiny tiny tiny fraction of an amu

and so what you really get here with

this mass number since uh protons and

neutrons each weigh

one amu this really and the electrons

weigh almost zero not exactly zero but

almost zero

so this number really just gets you the

total number of protons plus neutrons

and so in this case oxygen 16 here so

that mass number corresponds to having

16 total protons and neutrons combined

well if the atomic number already gave

us the fact that we have eight protons

then the remainder must all be neutrons

up to 16 total

and so that's going to tell us that we

must have eight neutrons

as well cool so

finally off on the other side here we've

got the charge

so and it turns out that atoms can often

lose and or gain

electrons from the electron cloud so

they typically aren't going to

lose or gain protons or neutrons from

the nucleus at least not by normal

chemical means

but they can lose or gain electrons from

the electron cloud

and so when you see that you've got a

negative charge here that means you've

got

two more negative charges than positive

charges

so in this case with oxygen having eight

protons which are positively charged

to have a negative two charge there must

be two additional

electrons associated with this atom and

so in this case

we can figure out that we've got

10 electrons with this oxygen ion

now one thing to note when we have a an

atom with no overall charge

it must be because the protons and

electrons are exactly the same number

so if this didn't actually have a charge

listed there if it just blank

that would imply that for the eight

protons in oxygen you also would have

eight electrons

not like the one we've got here and when

you've got a neutral species

that's when we call it a neutral atom

but the moment you get

a charge species rather than calling it

a charged atom

which is kind of technically true but

not the not the

the system of uh uh naming things we use

not the definitions we use

in this case we call this an ion so

neutral atoms are neutral atoms or atoms

but charge species we call ions

so first little piece of vocabulary here

cool uh a couple other things you should

know so

it turns out this is the most common

version of oxygen

but it's not the only one so there's

another one out there that exists

so called oxygen 18. in this case it's

got a mass number of 18

and the reason it's got a mass number of

18 must be because

let's work it out uh in this case

because it's oxygen it still has an

atomic number of eight

even if it's not listed and must have

eight protons

but with a mass number of 18 now

that means again this is protons and

neutrons combined and there's eight

protons so there must be ten

neutrons cool so it turns out that

different

uh uh there are different versions of

the same

element some heavier some lighter so

with a heavier mass number

lighter mass number and we call these

different

thinking of the definition as i write

this isotopes so

iso means the same so but you've got

different types of the same

element so and it's just because there's

a different number of neutrons

leading to a different mass number

that's what isotopes

are they're different types of the same

element different mass due to different

number of neutrons in the nucleus

okay last topic in this lesson we've got

to talk about atomic weight

also called atomic mass all right

so if you look at a typical periodic

table here so we've talked about having

an atomic number above the same atomic

number we saw in the isotope symbol i

just erased

but what you also have down below is an

atomic

mass or atomic weight so if like for

oxygen here it's 15.9994

so you get a nice little decimal number

and stuff like this and for carbon it's

12.011

things of this sort so what's weird is

students don't often get this is that

these numbers here these atomic masses

or atomic weights

are averages so we just saw that there

are different isotopes of oxygen that

actually exist

so it turns out one of the common ones

was oxygen 16 and then probably the next

most common the oxygen 18

so but it turns out there's not a whole

lot of oxygen 18 out there

in natural abundance and that's why when

you take the average it comes out almost

exactly to 16

because pretty close to 100 of the

naturally occurring oxygen atoms weigh

16

and so when you take the average it

comes out really close to 16. same thing

with carbon

so 98.9 of naturally occurring carbon

atoms weigh 12 and so when you take the

average of all of them

it comes out really close to 12. and for

most elements this is how it works

most of them have one major isotope and

when you take an average of

all the weights of all the naturally

occurring isotopes it comes really close

to one whole number because that's most

of it

but there's a couple elements where

that's totally not the case and one of

them here is

chlorine so if you look on your hand out

there i gave you a table there are two

major isotopes for chlorine not just one

and there's chlorine 35 and chlorine 37.

so and i put their exact masses

to two decimal places anyways so 34.97

amu's versus 36

0.97 amuse and then i also gave their

natural abundances so here it's it's

roughly a three to one ratio so roughly

75

to 25 percent but if we give it exactly

let's go 75.8 percent versus 24

point two percent now it turns out

there's

trace amounts of chlorine 36 that exist

but when i say trace amounts i mean like

point zero zero zero zero zero zero zero

zero zero zero something percent it's

super tiny that we can pretty much for

all practical purposes ignore it

but for chlorine 35 and 37 we've got a

fair amount

of both now i can see that we have

significantly more

chlorine 35 though and so with 75

and close to 76 percent and when we take

the average here then

the average should have come out closer

to 35 than it should to 37.

and that's what we see on the predict

table it's 35.4527

so a lot of periodic tables you'll see

round to the two decimal places and just

say 35.45

cool our students look at something like

that they're like so i know it's got 17

protons

because atomic number 17. so does it

have like 18 and a half neutrons or

something

well you can't have half a neutron it

doesn't work that way and the key is

that this

number isn't the same thing as a mass

number that we saw on the isotope

symbols

that isotope symbol that mass number was

a specific mass for a specific isotope

whereas this number right here is the

average mass

for all the naturally occurring isotopes

cool and you've got another process for

how we calculate these

based on giving just this kind of

information so and we're going to do

what's called a weighted average

and you're going to take the percentage

here and turn that into its decimal form

and

notice 75.8 percent means 75.8 over 100

which means move the decimal black back

two places here so we'll have

0.758 and you'll multiply that by the

corresponding mass

of that isotope so multiply that by 34.9

so cool we'll do the same thing with the

other isotope

so we've got 24.2 percent means per 100

so 24.2 over 100

which again just means move the decimal

back two places so that's

0.242 times its corresponding mass 36.97

cool and then you'll just add them

together and that will get you this

weighted average

and let me pull out my trusty calculator

here

all right so .758

times 34.97 plus

point two four two times thirty six

point nine seven and we get

thirty five point four five four

and if we look 35.454

is pretty close so now

i rounded these to two decimal places

and i rounded these to one

decimal place on the percents and had i

not rounded them we would have got

even closer to this number you see

published on the periodic table

cool but big takeaway here again so one

the process for just calculating these

atomic masses

and again when we say atomic weight or

atomic mass that already implies an

average of the naturally occurring

isotopes

cool and that's what's published on the

periodic table here

cool that sums up this lesson on atomic

theory and atomic structure in a chapter

on atoms molecules and ions

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