Transcript Chp 4 slideshow notes - Lower Cape May Regional School District

Chemistry Unit I:
Chapter 4 “Atoms”
Chp 4.1 Atomic Structure,
p. 113-118



Our understanding of atoms took many
centuries.
The 4th century Greek philosopher,
Democritus, suggested that the
universe was made of invisible units
called “atoms”. The word atom means
“that which cannot be divided”.
The Greek idea of atoms had to wait
2000 years before it became accepted.
John Dalton


In 1808, an English schoolteacher named
John Dalton proposed his own atomic
theory. He used evidence such as the “law
of definite proportions”.
His theory proposed:



Every ELEMENT is made of tiny, unique
particles called atoms that cannot be subdivided.
Atoms of the same element are exactly alike.
Atoms of different elements can join to form
molecules.
J.J. Thomson



In 1897, British scientist, J.J. Thomson, found
that atoms contain negatively charged
particles. These particles were named
“electrons”.
Since atoms are neutral, Thomson reasoned
that atoms must contain some sort of positive
charge. (Cathode Ray Tube experiment)
His model, like a plum pudding, had negative
electrons (plums) embedded in a positive
sphere (pudding).
Ernest Rutherford



In 1911, Ernest Rutherford, discovered that
atoms contain a dense, positively charged
“nucleus” through his Gold Foil Experiment.
Later Rutherford name the positive particles
“protons”.
In his model, atoms were mostly empty
space with electrons moving around a small,
very dense, positively-charged nucleus in
the center of the atom.
One addition…the neutron


In 1932, British scientist, James
Chadwick, discovered another particle
in the nucleus of an atom. The particle
was nearly the same mass as a
proton.
This particle is called a “neutron”,
because it is electrically neutral.
Chapter 4.2 “The Structure of
Atoms,” p.119-127




Three main subatomic particles are
distinguished by mass, charge and
location in the atom.
Proton, charge +1, located in the nucleus,
mass = 1.67x10-27
Neutron, charge 0, located in the nucleus,
mass= 1.67x10-27
Electron, charge -1, located in orbitals,
mass = 9.11x10-31
Atoms




Each element has a unique number of protons.
Each atom of the element will have the same
number of protons.
Atoms are electrically “neutral” overall. Atoms
have the same number of protons as electrons.
These charges exactly cancel leaving no
charge.
Positive and negative charges attract each
other with a force known as “electric force”.
Ions




Atoms may undergo a process called
“ionization”, in which they lose or gain
electrons in order to fill their valence energy
level.
An “ion” is an atom (or group of atoms) that
has lost or gained one or more electrons,
therefore has a net charge.
Removing electrons, causes a positive ion
to form, aka cation.
Adding electrons, causes a negative ion to
form, aka anion.
Atomic Nucleus




The atomic number tells you how many
protons are in an atom.
Every element has a different atomic
number.
Atoms may or may not have the same
number of neutrons as protons.
The mass number tells you how many
protons plus neutrons are in the nucleus of
an atom all together.
Isotopes



Some atoms contain a different number of
neutrons thereby making them “heavier” or
“lighter” than the average atom. These
atoms are called isotopes.
Isotopes have the same number of protons
just a different number of neutrons.
To calculate the number of neutrons,
subtract the atomic number (Z) from the
mass number (A) like this:

A – Z = number of neutrons
Atomic Masses




The mass of a single atom is extremely small!
Atomic masses are usually expressed in unified
atomic mass units.
A unified atomic mass unit (u) is equal to onetwelfth of the mass of a carbon-12 atom. Aka
“amu”
The average atomic mass for an element is a
weighted average and appears as a decimal
number on the periodic table.
Moles




Scientists often deal with very large
numbers of small particles. They use a
counting unit called a “mole”.
A mole is 602213670000000000000000 !!!
This number is usually written as
6.02 x 10 23 /mole and is called Avogadro’s
constant. This is how many particles are in a
mole of any substance.
Like the word “dozen” means 12, a mole of
a substance contains 6.02 x 10 23 particles.
Molar Mass



The mass in grams of 1 mol of a
substance is called its molar mass.
Example, one mole of lithium has a
mass of 6.941 g.
Because mass and amount of a
substance are related, it is possible to
convert moles to grams and vice
versa.
Calculating with Moles




Using a conversion factor allows you to
change from mass to amount and back
again.
The numerator of the conversion factor
contains the unit you want to change to. The
denominator of the conversion factor
contains the unit you start with.
The numerator and denominator equal 1
since they are equal.
Example: you could use either


10 gumballs/ 21.4 g
21.4 g/ 10 gumballs
Calculating with moles
Converting between amount of an
element in moles and its mass in
grams you can use this chart:
Amount (mol) x molar mass (g)/ 1 mol of
element = Mass (g)
Mass (g) x 1 mol of element/ molar mass
of element (g) = Amount (mol)

Chp 4-3 Modern Atomic
Theory p.128-132

Modern Model of the Atom
 Electrons can be found only in
certain energy levels, not between
levels.
 Furthermore, the location of
electrons cannot be predicted
precisely.
Niels Bohr



In 1913, Niels Bohr, a Danish scientist,
revised the Rutherford’s model.
Bohr showed that electrons could only
have specific amounts of energy,
leading them to move in certain orbits.
This model resembled planets orbiting
the sun.
Cloud of Electrons



In the 1920’s, scientists again revised the model.
They found that electrons can be anywhere in a
cloudlike region around the nucleus. This region is
called an orbital.
Electrons behave more like waves on a vibrating
string than like particles.
The different energy levels contain different
numbers of electrons depending on the energy level
size and electron energy.
Electron Energy Levels





The number of energy levels that are filled in an
atom depends on the number of electrons.
The electron(s) in the outermost energy level of an
atom are called “valence electrons”.
The most valence electrons that an atom can have
is eight (8).
Atoms with a filled valence energy level are the
“stable” and don’t react chemically.
Electron dot diagrams show the element’s symbol
and valence electrons as dots.
Electrons and Orbitals

Electrons may occupy four different kinds of orbitals
within atoms:




“S” is the simplest orbital, a sphere
shape, lowest energy.
“p” is dumbbell-shaped, and can be
oriented 3 ways, and are the next higher
energy levels.
“d” and “f” orbitals are much more
complex. There are 5 possible “d” orbitals
and 7 possible “f” orbitals. These are the
highest energy levels.
These orbitals are all different shapes
and can hold a maximum of two electrons
each.
Electron Transition





Electrons jump between energy level when an
atom gains or loses energy.
Ground state is the lowest state of energy of an
electron.
If an electron gains energy, it moves to an
excited state.
It gains energy by absorbing a particle of light
called a photon.
It will release the photon when it moves back to
its ground state.
Stop for Now!!!! Chapter 4
content complete

(the following slides are from Chp
5…please skip for now)
Chapter 5.3 “Families of
Elements,” p 156-165




The two main groups on the periodic table
are metals and non-metals.
Most metals are shiny solids that can be
stretched (ductile) and shaped (malleable).
The are good conductors of heat and
electricity.
Non-metals can be solid, liquid or gas. They
are typically dull, brittle and are poor
conductors.
Between the two groups are
semiconductors, or metalloids, which share
characteristics of both groups.
Alkali Metals



Group 1 of the periodic table contains
elements with one valence electron, which
can be easily removed to form a positive
ion.
They are all highly reactive and are not
found in nature as elements.
Examples include sodium and potassium.
These easily combine with non-metals to
form “salts”.
Alkaline Earth Metals



Group 2 of the Periodic Table are the
alkaline earth metals.
They contain 2 valence electrons and
react (not as violently as group 1)
losing their 2 electrons to become a
positive ion then forming compounds.
Examples in calcium and magnesium.
Transition Metals



Groups 3 through 12 contain the
transition metals.
They are less reactive than groups 1
or 2 but some still combine with nonmetals to form compounds. Some
have many versions of cations.
Examples include gold, silver, copper
mercury etc.
Non-metals



Except for hydrogen, non-metals are found
on the right side of the periodic table
including some elements in groups 13 thru
16 and all elements in groups 17 and 18.
Non-metals tend to share or gain electrons
(anion) and are plentiful on Earth.
Examples include hydrogen, carbon,
oxygen, sulfur, nitrogen, etc.
Halogens


Group 17 contain non-metals called
halogens. These elements contain 7
valence electrons and are highly
reactive.
Examples include fluorine, chlorine,
iodine, etc.
Noble Gases



Group 18 of the periodic table contain
the noble gases. Because these atoms
have filled valence energy levels, they
do not react with other elements to
form compounds….hence the name
“noble”.
They all exist as single gas atoms.
Examples include helium, neon, argon,
etc.
Semiconductors




Located along the “staircase line”,
semiconductors or metalloids have
properties of both metals and non-metals.
Semiconductors are able to conduct heat
and electricity under certain conditions,
some are hard, some are shiny and so on.
The 6 elements are:
Boron, silicon, germanium, arsenic,
antimony, and tellurium. Sometimes
polonium and astatine are also included.