Download Chapter 2 - Chemistry

Chapter 2
Atoms and Elements
Dr. Peter Warburton
[email protected]
http://www.chem.mun.ca/zcourses/1010.php
Law of conservation of mass
In a chemical reaction, matter is neither
created nor destroyed, but changed
from one form to another.
The total mass of chemicals in your
reaction mixture will not change as long as
you do not add any more chemicals to the
mixture, or allow any chemicals to escape.
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Law of conservation of mass
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Law of definite proportions
All samples of a given compound,
regardless of where they came from or
how they were made, will have the
same proportions of their constituent
elements.
The make-up of a compound is a
fundamental aspect of it’s identity. No
matter it’s source, a chemical must be
made of the same pieces…
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Law of definite proportions
Water ALWAYS contains 16.0 g of
oxygen atom per 2.0 g of hydrogen.
Water out of your bathroom tap,
Quidi Vidi Lake, the Atlantic Ocean,
or on Mars ALWAYS has an
8:1 ratio of O:H by mass
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Problem
Two samples of carbon monoxide are
decomposed into their constituent
elements. One sample produces 17.2 g of
oxygen and 12.9 g of carbon, while the
other sample produces 10.5 g of oxygen
and 7.88 g of carbon. Show these results
are consistent with the law of definite
proportions.
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Law of multiple proportions
When two elements (A and B) form two
different compounds, the masses of B
that combine with 1 g of A can be
expressed as a ratio of small whole
numbers.
This points to something fundamental
about how elements combine to form
compounds, in that elements often
combine in discrete integer combinations.
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Law of multiple proportions
Per gram of carbon:
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑂 𝑖𝑛 𝐶𝑂2 2.67 𝑔 2
=
=
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑂 𝑖𝑛 𝐶𝑂
1.33 𝑔 1
We can see in the formulas (and models shown) that CO2 has 2
oxygen atoms per carbon atom while CO has 1 oxygen atom per
carbon atom. This is reflected in the mass ratio!
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Dalton’s atomic theory (1808)
Each element is made of tiny indestructible
particles called atoms.
2. All atoms of a given element have the same
mass and physical/chemical properties that
distinguish them from atoms of other elements.
3. Atoms combine in simple whole number ratios
to form compounds.
4. Atoms of one element cannot change into
atoms of another element. In chemistry atoms
can only change how they are bound to other
atoms.
1.
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Atomic structure
What are these tiny, indestructible,
unchangeable particles called atoms and
what is there about them that makes them
different from each other that we call the
elements with different physical and
chemical properties?
It turns out atoms have several distinct
pieces!
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The electron
J.J. Thompson explored the
properties of cathode rays in his
cathode ray tube.
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The electron
The cathode rays always travelled in straight
lines from the negative end to the positive end,
regardless of what the cathode was made of…
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The electron
He deduced that cathode rays must carry a
negative electrical charge, and that all elements
must contain them (since it didn’t matter what
the cathode was made of)…
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The electron
Through a delicate experiment Thompson showed
cathode rays have about 1/2000th the mass of a
hydrogen atom, the lightest element. Somehow the socalled indestructible atom is made of smaller pieces!
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The electron
Thompson realized
cathode rays are small
pieces of charged matter
which we now call
electrons.
Negative charges attract
positive charges and repel
other negative charges
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Charge of an electron
Millikan used charged oil droplets in an electric
field to show the charge of one electron is the
smallest possible charge you can have…
-1.60 x 10-19 C
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Charge of an electron
Just as importantly, he was able to deduce the
mass of one electron… it’s an incredibly small
9.10 x 10-28 g
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Plum pudding model
Thompson proposed
that the negatively
charged electrons
were distributed
throughout a
positively charged
sphere that held all
the electrons together
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Radioactivity
Henri Becquerel, Marie Curie and others
studied the emission of small energetic
particles from unstable atoms… we call
this radioactivity.
Alpha (a) particles are positively charged
Beta (b) particles are negatively charged
Gamma (g) rays are very high energy
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Rutherford’s gold foil experiment
Rutherford realized he could probe the
“plum pudding” with the positively charged
a particles. He
“shot them” at a
thin piece of gold
foil to see how the
negative electrons
would deflect them
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Rutherford’s gold foil experiment
If the plum pudding model was correct,
then the positively charged particle should
pass through the
evenly distributed
charge of the atom.
This didn’t happen!
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Rutherford’s gold foil experiment
A few of the alpha particles were deflected
through large angles. Rutherford found it “about
as credible as if you
had fired a 15-inch
shell at a piece of
tissue paper and
it came back and
hit you.”
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The nuclear atom
1. Most of the atom’s mass and its positive charge are in
a small core called the nucleus.
2. Most of an atom is actually empty space where the
tiny negatively charged electrons are dispersed.
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The nucleus
The positive charge of a nucleus actually
comes from the positive charges of one or
more protons.
The charge of a proton is of the same size
as the charge of an electron, but of
opposite sign 1.60 x 10-19 C.
However a proton has a mass 1836 times
bigger than that of an electron!
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Neutrons
If a nucleus contains more than one
proton, why don’t the positive charges
repel each other and the nucleus falls
apart?
Nuclei also contain neutrons, which have
no charge and are just slightly more
massive than the protons.
Neutrons act like “a glue” that hold protons
together.
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The particles of an atom
Relative charges ignore the magnitude of the charge and just
focus on the charge. We can do this since there has never
been seen a charge less than that of a proton or electron.
The atomic mass unit (amu) is defined as 1/12th of the mass
of a carbon atom with 6 protons and 6 neutrons.
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Atomic number Z
We saw that atoms of an element have the
same properties and mass as other atoms of
that element, but they are different than the
atoms of a different element.
The fundamental difference between
elements is measured by the atomic number
Z. This number tells us how many protons
(and also an equal number of electrons)
an atom of that element contains.
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How much glue do you need?
If neutrons are the
“glue of the nucleus”
then how many do
you need to “do the
job”? It turns out
that the answer can
vary, EVEN for
atoms of the same
element!
http://www.craftsy.com/blog/2014/05/how-to-glue-up-wood/
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How much glue do you need?
In other words,
Dalton’s definition of
the atom is partially
wrong. Not all atoms
of the same element
have the same mass
because they have
different numbers of
neutrons!
http://www.craftsy.com/blog/2014/05/how-to-glue-up-wood/
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Isotopes
We call atoms of the same element that
have differing numbers of neutrons the
isotopes of that element.
We can easily describe all the fundamental
information about an element and it’s
isotopes by using three pieces of
information: the chemical symbol, the
atomic number and the mass number.
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Isotopes
The chemical symbol X is a one- or two-letter symbol representing the
elements name.
For most elements the symbol connects easily to the English name:
C for carbon, Ne for neon, Cs for cesium, etc
For 10 of the elements the chemical symbol is actually based on the Latin
name for the element: Na for sodium (natrium), K for potassium (kalium), Fe
for iron (ferrum), Cu for copper (cuprum), Ag for silver (argentum), Sn for tin
(stannum), Sb for antimony (stibium), Au for gold (aurum), Hg for mercury
(hydragyrum) and Pb for lead (plumbum)
One element has a symbol based on it’s Swedish name : W for tungsten
(wolfram)
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Isotopes
As we saw before, the atomic number Z tells us how many
protons (or electrons) are in ANY atom of the element.
The mass number A tells us how many protons AND
neutrons are in the SPECIFIC ISOTOPE of the element.
Therefore if we calculate
A–Z
we can know how many neutrons are in the nucleus of that
given isotope.
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Isotopes
Since the atomic number Z and the chemical symbol ultimately tell us
exactly the same thing (what element is this?) we often refer to an
isotopes name just by the chemical symbol and it’s mass number.
i.e. Ne-20, Ne-21 or Ne-22
All of them contain 10 protons (Z = 10 for neon).
The first isotope has 10 neutrons, the second isotope has
11 neutrons and the third has 12 neutrons.
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Isotopes of neon
Neon has three naturally occurring isotopes.
The natural abundance tells us the relative percentage of each
type of isotope we would find in a sample of that element.
For example, if I had a sample of 10000 neon atoms, I would find:
9048 of them are Ne-20 with 10 neutrons
27 of them are Ne-21 with 11 neutrons
925 of them are Ne-22 with 12 neutrons.
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Problem
What are the atomic number, mass
number and symbol for the carbon isotope
with seven neutrons?
How many protons and neutrons are
present in an atom of
39
𝐾
19
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Ions
Sometimes atoms can gain or lose electrons to
form charged ions as compared to the neutral
atoms. Positively charged ions (loss of
electrons) are called cations while negatively
charged ions (gain of electrons) are called
anions.
Li  Li+ + 1 eF + 1 e-  F-
Mg  Mg2+ + 2 eO + 2 e-  O2-
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Atomic mass
The concept of isotopes leads to an
interesting conclusion. The chemical
properties of an element are determined
by the number of protons and electrons in
that element, and NOT by the mass of the
element. If mass were a factor, then the
chemical properties of different isotopes of
an element would be different for each
isotope. They’re not!
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Atomic mass
Because of different isotopes having
different masses and natural abundances,
it makes sense to assign an atomic mass
to each element that includes all isotopes.
The atomic mass is a weighted average.
Atomic mass
=
𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑛 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑛
𝑛
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Atomic mass
Chlorine has two isotopes: Cl-35 and Cl-37
Cl-35 makes up 75.77% of all chlorine atoms
Cl-37 makes up 24.23% of all chlorine atoms
0.7577 34.97 𝑎𝑚𝑢 + 0.2423 36.97 𝑎𝑚𝑢
= 35.45 amu
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I need some volunteers…
You only need to stand up in front of the
class and be willing to say how tall you
are.
Any takers?
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Problem
Titanium has five naturally occurring
isotopes. Using the info in the table below,
calculate the atomic mass of titanium:
Isotope Natural abundance (%) Isotope mass (amu)
Ti-46
8.25
45.9526
Ti-47
7.44
46.9518
Ti-48
73.72
47.9479
Ti-49
5.41
48.9479
Ti-50
5.18
49.9448
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Measuring isotope masses
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Moles and molar mass
If the mass of an isotope does not affect
it’s chemical properties, then we are likely
more interested in how many atoms of
that element we have. The problem is how
do we count atoms? It turns out we can do
that by mass (since weighing is easy to
do). But we do need to know the
connection between mass and number of
atoms!
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Mole
We saw earlier that the SI unit of amount
of substance is the mole (mol), which we
characterized as the “chemist’s dozen”
1 mol = 6.02214 x 1023 particles
This value is Avagadro’s number.
Notice the use of the word “particles”,
since we can talk about a mole of atoms,
electrons, ions, or molecules.
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Mole
Since the mole is a SI unit it actually has a
VERY specific definition:
“The value of the mole is equal to the
number of atoms in exactly 12 grams
of carbon-12”
Notice this connects a number of
atoms to the total mass they have!
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Problems
How many atoms are there in 3.92 mol of
phosphorous?
If I have 3.69 x 1018 fluoride ions (F-), then
how many moles of fluoride ions do I
have?
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Molar mass
We’ve seen the definition of the mole is
connected to a specific mass of carbon-12.
If we want to convert from a mass of a
substance to the number of moles of that
substance (or vice versa) then we would
need equivalent mass/moles relationships
for other substances. We call this the
molar mass (M)
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Molar mass
The molar mass of any element, in
grams per mol (g mol-1 or g/mol), is
numerically equal to the atomic mass of
the element in atomic mass units.
Earlier we saw that the atomic mass of
chlorine was 35.45 amu.
This means the molar mass of chlorine is
35.45 g
-1
mol
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Molar mass
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Problems
How many carbon atoms are there in a
1.3-carat diamond (which is pure carbon) if
1 carat = 0.20 g)?
Calculate the mass of 2.25 x 1022 tungsten
atoms.
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The periodic table
Dmitiri Mendeleev organized the known
elements (in 1869) according to his
periodic law that the elements showed
repeating patterns in terms of what
proportions they reacted with each other
and trends in their physical properties.
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Mendeleev’s periodic table
Ordered by increasing mass
Grouped by formulas of hydrides and oxides
Blanks left for elements yet to be discovered
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Modern periodic table
Now grouped by atomic number
Missing elements are here!
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Additional detail in periodic table
Elements
are
classified
as
Metals
Metalloids
Nonmetals
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Modern periodic table
Down the periods (columns)
the elements of the group
show similarity of chemistry
and trends in properties
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Group 18 – Noble gases
The elements of the last column (Group
18) are the noble gases. These elements
are remarkably unreactive (thus the
name).
He
Ne
http://en.wikipedia.org/wiki/Neon#/media/File:NeTube.jpg
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Group 1 – Alkali metals
The first column of
the periodic table
(Group 1) are the
very reactive alkali
metals which react
vigorously with
water, for example.
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Group 17 - Halogens
Next to the noble
gases are the Group
17 halogens which
are very reactive
non-metals. They
will react quite
vigorously with alkali
metals, for example.
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Ions and the periodic table
Main group metals tend to lose electrons to form cations with the same
number of electrons as the noble gas in the previous row.
Main group non-metals tend to gain electrons to form anions with the same
number of electrons as the noble gas in the same row.
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Ions and the periodic table
These tendencies explain why the alkali
metals and halogens are very reactive,
especially with each other!
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