Download PowerPoint for Ch 2 Part 2 - Dr. Samples` Chemistry Classes

Chem 400
Chapter 2
Going Further: The Structure of Atoms
• Dalton thought that atoms were the smallest
particle of matter, but through a series of
experiments starting in the late 1800’s, this was
proved to be incorrect.
• In a series of experiments by various scientists,
the existence of electrons, protons, and
neutrons were deduced and verified.
Going Further: The Structure of Atoms
• Electrons were discovered in 1897 by JJ
Thomson, and they were found to have a
negative electric charge.
• Protons were hypothesized in 1911 by Ernest
Rutherford, and were verified in 1919 by
Rutherford. They have a positive electric
charge.
• Neutrons were discovered in 1932 by James
Chadwick. They have no electric charge.
Going Further: The Structure of Atoms
• There are 2 important experiments that you
should be aware of:
– Robert Millikan’s Oil-Drop Experiment of 1909,
which enabled him to measure the magnitude of the
electric charge on electrons, and calculate their mass
in grams.
– Ernest Rutherford’s 1911 Gold-Foil Experiment, also
called the Alpha-Scattering Experiment. This
experiment enabled him to hypothesize the existence
of protons in the nucleus of atoms.
Oil-Drop Experiment
The Structure of Atoms
• From this experiment, Millikan obtained the actual
charge on an electron, -1.60x10-19 C.
• And from this charge and Thomson’s charge/mass
ratio, the exact mass of an electron was calculated to be
9.10x10-28 g.
• So from these experiments, scientists deduced that
atoms were made up of even smaller subatomic
particles, one of which was the electron.
• Since electrons have a negative charge, while atoms are
neutral, scientists also realized that there had to be at
least 1 more subatomic particle with a positive charge.
The Structure of Atoms
• Where was this atomic particle & what
did it look like?
Gold-Foil Experiment
The Structure of Atoms
• How could something in the atom cause such
huge deflections in a massive positively charged
particle like an alpha particle?
• So Rutherford proposed that atoms were
composed of mostly empty space (where the
electrons moved in circular orbits) with a very
small, very massive, very dense center called
the nucleus.
• The nucleus has a positive charge. This was
Rutherford’s Atomic Model.
The Structure of Atoms
The Structure of Atoms
• Rutherford then proved the existence of
protons in 1919, and neutrons were
discovered by James Chadwick in 1932.
• So what’s the overall picture of an atom,
and what are the sizes, masses, charges,
and densities of the particles and regions?
The Structure of Atoms
Particle
Mass (g
x10-24)
Mass
(amu)
Relative
Mass
Charge
(C x10-19)
Relative
Charge
Location
in Atom
proton,
p
1.673
1.0073
1
1.602
1
nucleus
electron, 0.000911 0.000549
e or e-
0
-1.602
-1
electron
cloud
neutron, 1.675
n
1
0
0
nucleus
1.0087
The Structure of Atoms
• The diameter of a typical atom is around 1x10-10 m or 1
Å.
• The diameter of a typical nucleus is only 0.0001 Å.
• You can see that most of the mass of the atom is
contained in a very small volume, so the nucleus is
incredibly dense.
• The density of a typical nucleus is 1x1013 to 1x1014
g/cm3, beyond our comprehension! If a matchbox had
this density, it would weigh 2.5 billion tons!
The Structure of Atoms
Atomic Number, Mass Number and Isotopes
• Dalton thought that atoms of different elements
differed mainly by mass, but we now know that
atoms of different elements differ by the number
of protons which they contain.
• The number of protons which an element
contains is called the Atomic Number, Z.
• The Atomic Number is found on the Periodic
Table above the elemental symbol.
Atomic Number, Mass Number and Isotopes
Atomic Number, Mass Number and Isotopes
• It is true that every atom of the same element
contains the same number of protons.
– So every H atom has 1 proton, and every C atom has
6 protons.
• So the number of protons defines the element.
• But it is NOT true that all atoms of the same
element are identical.
• What’s different? Well, what else is there?
Atomic Number, Mass Number and Isotopes
• Although all atoms of the same element have
the same number of protons, they DO NOT
have the same number of neutrons!
– and if it is a neutral atom, they all have the same
number of electrons
• Atoms of the same element which have
different numbers of neutrons are called
isotopes.
Atomic Number, Mass Number and Isotopes
• So isotopes of the same element differ by the
number of neutrons.
• And since neutrons are the same relative mass
as protons, isotopes also differ by mass.
• Although this is not shown on the Periodic
Table, every element has at least 2 isotopes
(except some of the newly synthesized elements
like Mt).
Atomic Number, Mass Number and Isotopes
• To show different isotopes, we have several different
isotopic notations or isotopic symbols.
• They all use the Mass Number, A, which is the sum of
the protons and neutrons in the nucleus.
• For example, H has 3 common isotopes, H-1, H-2, and
H-3.
• Carbon also has 3 common isotopes, C-12, C-13, and C14.
• The number after the symbol or the superscript left
number is the Mass Number.
Atomic Number, Mass Number and Isotopes
Mass Number
1
H
1
1
Mass Number
H H-1 Hydrogen-1
Atomic Number
• Practice with Isotopic Notation: How many electrons,
protons, and neutrons do the following isotopes have?
112
Ag Se-77
126
Te Tl-205
136
54
Xe
Ions and Ionic Isotopic Notation
• A neutral atom has equal numbers of protons
and electrons. Why?
• During chemical reactions, atoms may gain or
lose (or share) electrons, e-.
• If an atom gains or loses 1 or more e-, there is
an imbalance between protons and e-, and the
result is a charged particle called an ion.
• So an ion is formed when an atom gains or loses
e-.
Ions and Ionic Isotopic Notation
• If an atom loses 1 or more e-, then it has more
protons than e-, so the ion has a + charge. It is
called a cation.
• If an atom gains 1 or more e-, then it has less
protons than e-, so the ion has a - charge. It is
called an anion.
• Note that it is difficult to gain more than 3 e- or
lose more than 4 e-.
• How do we show ions?
Ions and Ionic Isotopic Notation
• If we have an isotope which is an ion, we can
show a complete isotopic notation for the ion.
• How many protons, e-, and neutrons do the
following ionic isotopes have?
Elements and the Periodic Table
• Elements are fundamental substances.
• They can’t be broken down into smaller
substances by chemical reactions.
• The Periodic Table arranges the known
elements (114 of them).
• 90 of these are naturally occurring, while
the rest have been synthesized in nuclear
reactions.
Elements and the Periodic Table
Elements and the Periodic Table
• Notice that the elements names have been given
shorthand notations (called symbols) of 1 or 2
letters.
• Unnamed elements actually have a 3 letter
designation until they are named.
• The first letter is ALWAYS capitalized, while
the second letter is ALWAYS lowercase.
• What elements do you have to memorize
(names and symbols)? 1-40; 42; 46-57; 76-90;
92; and 94.
Elements and the Periodic Table
• Although most of the symbols are obviously
related to the name, like N for nitrogen, others
seem to make no sense, like Pb for lead!
• This is because some of the symbols come from
old Latin names or other languages.
• Plumbum was an old Latin name for lead.
• W for tungsten comes from the German name
wolfram.
Elements and the Periodic Table
• Chemistry in some fashion has been
around for centuries.
• Some elements were known
thousands of years ago.
• But most elements were discovered
and identified in the last 250 years.
Elements and the Periodic Table
• In the early to mid 1800’s, chemists were
trying to organize the 60-some known
elements into some sort of pattern.
• Mendeleev designed a Periodic Table in 1869
which was based on the masses of the known
elements (atomic weights) and the compounds
they formed with hydrogen (hydrides) or
oxygen (oxides).
• Today’s Table is similar, but the elements are
arranged by atomic numbers (number of
protons) instead of by atomic weights.
Elements and the Periodic Table
• If you look at a Periodic Table, there are 18
columns called Groups or Families.
• They are called families as they share common
chemical properties or characteristics.
• The 7 rows are called Periods.
• The groups are numbered 2 ways on US Tables.
• The old US system uses numbers with A or B
sections, while the internationally approved
system simply numbers the groups from 1 to 18
going across from left to right.
Elements and the Periodic Table
• There are several basic regions on the Table:
–
–
–
–
–
–
–
–
Metals
Nonmetals
Semimetals (metalloids or semiconductors)
Main Block or Representative Elements
Transition Metals
Inner Transition Metals
Lanthanides
Actinides
Elements and the Periodic Table
• Several important Groups also have
names:
– Group 1, except hydrogen, are the
Alkali Metals
– Group 2 is the Alkaline Earth Metals
– Group 17 are the Halogens
– Group 18 are the Noble Gases
Elements and the Periodic Table
• Metals: lustrous, silvery, malleable, ductile,
generally hard, solids except Hg, conductors,
lose electrons to become cations, react with
nonmetals to form ionic salts.
• Nonmetals: nonconductors, react with metals,
gain electrons to form anions, brittle, most gases
(1 l, 5 s)
• Metalloids: B, Si, As, Te, At, Ge, Sb: in between
metals & nonmetals, semiconductors, solids
Elements and the Periodic Table
• Here are some shared characteristics in the
regions and groups:
• Alkali Metals: Very reactive metals, soft, not
found in nature as pure element
• Alkaline Earth Metals: same as Alkali metals
but less reactive
• Halogens: most reactive nonmetals, corrosive,
not found in nature as pure element
Elements and the Periodic Table
• Noble Gases: also called Inert gases as very
nonreactive, don’t form compounds except Xe
• Lanthanides: f-fillers, rare earth metals, inner
transition metals, reactive, silvery-grey
• Actinides: f-fillers, rare earth metals, inner
transition metals, reactive, silvery-grey,
radioactive, synthetic above 92
• Why is hydrogen placed in Group 1 if it is NOT
an Alkali Metal and is actually a nonmetal?
Elements and the Periodic Table
Elements and the Periodic
Table
Atomic Mass and Weighted Averages of Elements
• As atoms have a very tiny mass in grams,
scientists use another scale to state the masses
of atoms, the atomic mass unit, amu. You see
this in the table with the masses of p, e, and n
given earlier.
• The conversion factor between mass in g and
mass in amu is:
1 amu = 1.66054x10-24g OR
1 g = 6.02214x1023amu
Atomic Mass and Weighted Averages of Elements
• The average atomic mass of the elements is
shown beneath the elemental symbol on the
Periodic Table.
Atomic Mass and Weighted Averages of Elements
• But every element has different isotopes with
different masses!
• That’s why the atomic masses on the Table are
average masses: it is really the mass in amu of a
single “average” atom of an element.
• But what does an “average” atom of an element
look like?
• What does the “average” student look like?
Does “it” exist?
Atomic Mass and Weighted Averages of Elements
• For H, 99.985% of all H atoms are H-1, while 0.015%
are H-2 (there are basically 0% H-3). This is called the
natural abundance or %-abundance of an isotope.
• So shouldn’t the “average” H atom look a lot like H-1,
and shouldn’t the average atomic mass of H be very
close to the mass of the H-1 isotope?
• Because the different isotopes do not have equal
natural abundances, we calculate atomic masses of
elements using a weighted average of all the isotopes:
Atomic Mass and Weighted Averages of Elements
• So if lead has 4 common isotopes with the
following masses and %-abundance, what is
the atomic mass of lead?
Pb-204
Pb-206
Pb-207
Pb-208
203.973020 amu
205.974440 amu
206.975872 amu
207.976627 amu
1.40%
24.1%
22.1%
52.4%
Molar Mass of Atoms & Avogadro’s Number
• The atomic mass is the mass in amu of a single
“average” atom.
• Is this useful in the lab? Can we pick out and
weigh an individual atom?
• Chemists weigh in g, which is a HUGE number
of atoms.
• So we need a unit to express large numbers of
atoms or molecules without using scientific
notation.
• Chemists defined a counting unit to do this.
Molar Mass of Atoms & Avogadro’s Number
• They chose a unit called mole so that the atomic
mass on the Periodic Table is also used for
measuring grams.
• 1 mole = 1 mol = 6.02214x1023 things
• So if you have 1 mol of pennies, how many
pennies do you have? How many dollars is
this?
• The number 6.02214x1023 is Avogadro’s
number.
Molar Mass of Atoms & Avogadro’s Number
• Avogadro’s number is very important as it is a
conversion factor between number of things
and moles of things.
• EX: if you have 2.5 mol of aluminum, how
many atoms of aluminum do you have?
• Avogadro’s number is also very special as if we
have 1 mol of an element, we have the atomic
mass in g of the element.
Molar Mass of Atoms & Avogadro’s Number
• If we have 1 mol of Na atoms, prove that this is
22.99 g of Na:
Molar Mass of Atoms & Avogadro’s Number
• The molar mass is the mass in g of exactly 1 mol
of an element.
• The atomic mass and the molar mass are the
same number, they differ only by units!
Molar Mass of Atoms & Avogadro’s Number
• The molar mass is also a conversion factor: it
converts between mol of an element and g of an
element.
– Ex: If you have 25.7840 g of gold, how many mol of
gold is this?
• There are 6 types of simple calculations that we
can do using molar masses and Avogadro’s
number.