Download Lecture Notes Part 2 - Dr. Samples` Chemistry Classes

Chem 400 Chapter 2 Lecture Notes Part 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.
• Electrons were discovered in 1897 by JJ Thomson, and they were found to have a negative charge.
• Protons were hypothesized by Ernest Rutherford in 1911 and their existence was verified in 1919 by
Rutherford. Protons have a positive charge.
• Neutrons were discovered (although they had been deduced earlier) in 1932 by James Chadwick.
They are neutral.
• There are 2 important experiments that you need to be aware of: Robert Millikan’s Oil-Drop
Experiment on 1909, and Ernest Rutherford’s 1911 Gold-Foil Experiment, also called the alphascattering experiment.
•
In 1909, Robert Millikan conducted his famous Oil-Drop Experiment:
•
From this experiment, Millikan obtained the actual charge on an electron, 1.60x10-19 C (Don’t
memorize!).
From this charge and the charge/mass ratio of an electron which JJ Thomson had measured, the
exact mass of an electron was calculated to be 9.10x10-28 g (Don’t memorize!).
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.
But where was this particle in the atom and what did it look like?
In 1911, Ernest Rutherford conducted his famous -scattering experiment, or the Gold-Foil
Experiment:
•
•
•
•
•
Rutherford was shocked! 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, so very dense center called the nucleus.
The nucleus has a positive charge. This was Rutherford’s Atomic Model.
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?
Particle
Mass (g)
Mass (amu)
Relative
Mass
Charge (C)
Relative
Charge
Location in
Nucleus
proton,
p
electron,
eneutron,
n
•
•
•
•
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 atom is 1x1013 to 1x1014 g/cm3, beyond our comprehension! If a matchbox
had this density, it would weigh 2.5 billion tons!
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 that 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.
•
It is true that every atom of the same element contains the same number of protons. So every H atom
has 1 proton, 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.
•
Although they all have the same number of protons (and if it is a neutral atom, they all have the
same number of electrons), they DO NOT have the same number of neutrons!
•
Atoms of the same element that have different numbers of neutrons are called ___________.
•
So __________ differ by the number of neutrons. And since neutrons are the same relative mass as
protons, ___________ 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).
•
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 _________ and ________ in the nucleus.
•
For example, H has 3 common isotopes, _______________, while carbon also has 3 common
isotopes, ______________. The number after the symbol or the superscript left number is the Mass
Number.
•
Practice with Isotopic Notation: How many electrons, protons, and neutrons, do the following
isotopes have?
Ions: Atoms Gaining or Losing Electrons
• For a neutral atom, the number of protons must equal the number of electrons.
• During chemical reactions, atoms may gain or lose electrons (and share as well).
• If an atom gains or loses one or more electron, it now has an imbalance between protons and
electrons, so it gains a charge.
• A particle which has a charge due to gaining or losing 1 or more electrons is called an ion.
• Positively charged ions are called __________ and they have __________ 1 or more electrons.
• Negatively charged ions are called __________ and they have __________ 1 or more electrons.
• Note: it is difficult to gain more than 3 electrons, and it is difficult to lose more than 4 electrons.
• How do we show ions? (Al; Mg; S; F)
•
We can also have a complete isotopic notation for ions as well. How many protons, electrons, and
neutrons do the following have?
Elements, the Periodic Table, and their Properties
• Elements are fundamental substances. They can’t be broken down into smaller substances by
__________ reactions.
• The Periodic Table arranges the known elements (114 of them). ___ of these are naturally occurring,
while the rest have been synthesized in nuclear reactions.
• 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 _______________, while the second letter is ALWAYS ____________.
• 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, while W for tungsten comes from the German
name wolfram.
• What elements do you have to memorize (names and symbols)? 1-40; 42; 46-57; 76-90; 92; and 94.
• Although chemistry in some fashion has been around for centuries, and some elements were known
thousands of years ago, most elements were discovered and identified in the last 250 years.
• 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.
• If you look at a Periodic Table, there are 18 columns called ___________ or ______________.
• They are called ______________ as they share common chemical properties or characteristics.
• The 7 rows are called ____________.
• The groups are numbered in 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.
• There are several basic regions on the Table: metals, nonmetals, and semimetals (metalloids); and
there are Main Block or Representative Elements, Transition Metals, and Inner Transition Metals.
The Inner Transition Metals are further broken down into Lanthanides and Actinides.
• Several important Groups also have names: Group 1, except hydrogen, is the ______________;
Group 2 is the ______________ Metals; Group 17 is the __________; and Group 18 are the
___________.
• Here are some shared characteristics in the regions and groups:
Metals:
Nonmetals:
Metalloids:
Alkali Metals:
Alkaline Earth Metals:
Halogens:
Noble Gases:
Lanthanides and Actinides:
• Why is hydrogen placed in Group 1 if it is NOT an Alkali Metal and is actually a nonmetal?
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 above.
•
The conversion factor between the amu unit and the gram unit is:
•
The average atomic mass of the elements is shown beneath the elemental symbol on the Periodic
Table.
•
But remember, 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?
•
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:
Molar Mass of Atoms & Avogadro’s Number
• The number underneath an Elemental Symbol on the Periodic Table is its atomic mass (the mass of 1
“average” atom in amu).
• Is a formula mass very useful in the lab? Well, can we weigh out individual atoms or ions on a lab
balance? They are too tiny to weigh or pick out individually!
• Chemists weigh out compounds and elements in grams, which contain a huge quantity of atoms or
ions.
• Chemists needed a way to express large amounts of atoms easily without scientific notation.
• Therefore chemists defined a counting unit so that they could use the same atomic mass for
measuring grams as well.
• This counting unit (just like a dozen is a counting unit equal to 12) is called the mole (abbreviated
mol) where:
• This number is called Avogadro’s Number. Avogadro’s Number is a conversion factor between
atoms (or ions) and moles of a substance.
• Ex: If you have 2.5 mol of aluminum, how many atoms of aluminum do you have?
• Again, it’s really convenient as if we have 1 mol of Na atoms, we have 22.9898 g of Na.
• The mass in grams of exactly 1 mol of any substance is called the molar mass.
• So the atomic mass and the molar mass of a substance are the same number, but they have different
units:
• The molar mass is important for another reason: IT’S A CONVERSION FACTOR! It converts
between g of a substance and moles of a substance. You will learn to love it!
• Example: If you have 25.7840 g of gold, how many moles do you have?
• There are 6 types of simple calculations with moles
• Moles to mass in grams: How many g are in 35.7 moles of silver?
• Mass in grams to moles: How many mol are in 358.0 g of iron?
• Moles to atoms: How many atoms are in 1.04 mol of gold?
• Atoms to moles: If you have exactly 1 billion atoms of potassium, how many mol do you have?
• Mass in grams to atoms: How many atoms are in 12 g of silicon?
• Atoms to mass in grams: Convert 4.57x1025 atoms of silver to g silver.