Download Atoms - Mrs. Carlyle`s Classroom

Chapter 3
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
The smallest particle of an element that retains the
chemical properties of that element.

Regions:

Subatomic Particles:
◦ Nucleus: very small region located at the center of atom
◦ Outside the nucleus: electrons most in the “electron cloud”
◦ Neutron, neutral found in nucleus (mass number subtracted
by number of protons indicates neutrons)
◦ Proton, positive found in nucleus (atomic number indicates
number of protons in the atom)
◦ Electron, negative found surrounding the nucleus (equal
number of protons and electrons)

Dalton’s atomic theory explains all three of these
laws.
◦ All elements were composed of atoms and that only whole
numbers of atoms can combine to form compounds. The
following statements sum up his theory:
 All matter is composed of extremely small particles called atoms.
 Atoms of an element are identical in size, mass, and other
properties.
 Atoms cannot be subdivided, created, or destroyed.
 Atoms of different elements combine in simple-whole number
ratios to form chemical compounds
 In chemical reactions, atoms are combined, separated, or
rearranged.

Several changes have been made to Dalton’s
theory…
◦ Dalton said…Atoms of an element are identical in size,
mass, and other properties.
◦ Modern Theory states: Atoms of an element have the
same number of protons (atomic number) but can have
different masses (because of the different number of
neutrons)
◦ Dalton said…Atoms cannot be subdivided, created, or
destroyed.
◦ Modern Theory states: Atoms cannot be subdivided,
created, or destroyed in ordinary chemical reactions.
However, these changes CAN occur in nuclear reactions!

Late 1800s, JJ Thomson, used an cathode ray
tube to determine the existence of negative
electrons by passing an electric current
passed through low pressure gases.
◦ Cathode-ray tube experiment
 Cathode (- charged metal disk side of the tube)
 When electricity was passed through the tube, the
directly opposite side of the cathode would glow.
 The assumed the glow was due to the particle stream
and called this glow a “cathode ray.”
 Particles that compose cathode rays are negatively
charged (since they were attracted to the + anode end)
 Ultimately, this is how an electron was found!

Ernest Rutherford’s Gold Foil Experiment
◦ Fired helium nuclei (alpha particles) at a thin sheet
of gold foil
◦ A screen record where they hit


Most of the particles passed right through
A few particles were deflected
•

Only about 1 in 8000 alpha particles bounced off the gold
foil.
VERY FEW were greatly deflected
◦ Conclusion:
 The nucleus is small
 The nucleus is dense
 The nucleus is positively charged

Protons
 Positive charge equal in magnitude to the negative charge of
an electron.
 1.673 x 10-27 kg

Neutrons
 Electrically neutral charge.
 1.675 x 10-27 kg
 Neutrons are slightly larger in mass than protons

The nuclei of atoms of different elements differ
in their number of protons, thus the number of
protons determines the atom’s identity.



Generally, particles that have the same charge
repel one another.
More than one proton in the nucleus you would
expect to be unstable. However, it’s the opposite.
When two or more protons are extremely close to
each other, they have a strong attraction.
Nuclear force is known as short-range protonneutron, proton-proton, and neutron-neutron
forces hold the nuclear particles together.

Developed the planetary model and proposed
that the electrons orbit around the nucleus in
a circular part at a constant speed.
◦ Incorrect!!

Quantum Mechanical Model
◦ The model we use today!!
◦ Also called the electron cloud model.
◦ The electrons don’t have a distinctive orbit, they
have a probability region where they can be found.
 Think about a fly buzzing around the room…very
similar to how an electron would act!
Charge
Mass
Location
Proton
+
1 amu
Nucleus
Neutron
None
1 amu
Nucleus
Electron
-
Almost
nothing…
Outside Nucleus
(in motion)
For any element:
# of Protons = Atomic #
# of Electrons = # of Protons = Atomic #
# of Neutrons = Mass # - Atomic #



Atoms of different elements have different
number of protons.
The atomic number of an element is the
number of protons of each atom of that
element.
Hydrogen, H, has atomic number 1.
◦ One proton in each atom of Hydrogen.

Atoms of the same element that have
different masses.
 Remember, protons and neutrons make up the
majority of the mass of an atom.
 All isotopes of hydrogen have the same number of
protons and electrons but vary in neutrons.


Most elements consist of isotope mixtures.
Tin has 10 stable isotopes, the most of any
element.


Just because every
atom of Hydrogen has
one proton…it doesn’t
mean they have the
same number of
neutrons.

Protium
◦ Accounts for 99.9885% of
hydrogen atoms on Earth
◦ No neutrons, 1 proton
◦ Total mass of 1

Deuterium
◦ Accounts for 0.0115% of
hydrogen atoms on Earth
◦ 1 neutron, 1 proton
◦ Total mass of 2
Actually, three types of
Hydrogen atoms exist!

Tritium
◦ Radioactive and is rarely
found on Earth only in
trace amounts.
◦ 2 neutrons, 1 proton
◦ Total mass of 3

Two Methods:
◦ Hyphen-Notation – where the mass number appears after
the element with a hyphen.
◦ Nuclear-Symbol – the composition that shows the mass
number and atomic number
 Nuclide – is a general term for a specific isotope of an element

Mass is neither
created nor
destroyed during
ordinary chemical
reactions or
physical changes.

A chemical compound contains the same
elements in exactly the same proportions
by mass regardless of the size of the
sample or source of the compound.

If two or more different compounds are
composed of the same two elements, then the
ratio of the masses of the second element
combined with a certain mass of the first
element is always a ratio of small whole
numbers.


Scientists use the relative atomic masses
because they are most convenient.
Standard for this measure is the carbon-12.



The weighted average of the atomic masses
of the naturally occurring isotopes of an
element.
Depends on mass and relative abundance
Calculated by multiplying the atomic mass of
each isotope by its relative abundance and
adding the results.

Suppose you have a box containing two sizes of marbles. If
25% of the marbles have masses of 2.00 g each and 75% have
masses of 3.00 g each, how is the weighted average
calculated? Suppose you have 100 marbles.
25% of 100 = 25 marbles
75% of 100 = 75 marbles
25 marbles x 2.00 g = 50 g
75 marbles x 3.00 g = 225 g
50 g + 225 g = 275 g ÷ 100 marbles = 2.75 g per marble

What chemical laws
can be explained by
Dalton’s theory?

Three compounds
containing potassium
and oxygen are
compared. Analysis
shows that for each
1.00 g of O, the
compounds have
1.22 g. 2.44 g, and
4.89 g of K,
respectively. Show
how these data
support the law of
multiple proportions.