3.1 Atomic Theory and Atomic Structure | High School Chemistry
Summary
TLDRThis 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
🔬 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.
🌌 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.
📊 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.
⚛️ 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.
📐 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.
Mindmap
Keywords
💡Atomic Theory
💡Subatomic Particles
💡Isotopes
💡Periodic Table
💡Electron
💡Proton
💡Neutron
💡Atomic Mass Unit (AMU)
💡Nucleus
💡Ion
💡Atomic Weight
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
so if you feel like you got something
out of this lesson please consider
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