Download Physical Science Chapter 16 Notes Section 1: Structure of the Atom

Physical Science
Chapter 16 Notes
Section 1: Structure of the Atom
1. Atom – the smallest piece of matter that still retains the property of the element
♦ All atoms are composed of protons, neutrons, and electrons
2. Quarks – small particles that bond together to form protons and neutrons
3. The Atomic Model
♦ Ancients Greeks thought matter was composed of tiny, indivisible particles called atoms
• Four different atoms: fire, water, earth, and air
♦ In 1803 John Dalton developed the first true atomic theory
• Believed that atoms may be distinguished by their mass
• Stated that all atoms of the same element have the same mass
• Atoms of different elements have different masses.
♦ In 1897 J.J Thompson proved the existence of negatively charged particles within atoms
• In a cathode ray tube particles would flow from one electrode to another
• Also, the rays would bend toward a positively charged plate, so the ray had to carry a
negative charge
• Thompson discovered the first subatomic particle, the negatively charged electron
♦ In 1911 Ernest Rutherford demonstrated the existence of protons
• Alpha particles – positively charged particles emitted by certain radioactive materials
• Rutherford fired alpha particles at a very thin sheet of gold foil that was surrounded
by a screen. The screen would glow whenever an alpha particle struck it. Rutherford
found that three things would happen: some of the alpha particles would pass
straight through the foil, some would be deflected by electrons, and some would
bounce back.
• Rutherford deduced that the alpha particles that bounced back were repelled by
subatomic, positively charged particles. These positively charged subatomic particles
are called protons.
♦ Even with the discovery of protons scientists knew that there was still some “missing”
mass in atoms
• Example: hydrogen has one proton and one electron, so helium should have just two
protons and two electrons. If that were true, the ratio of the mass of the helium atom
to the mass of a hydrogen atom would be 2:1, but experiments showed that the ratio
was actually 4:1.
• In 1932 James Chadwick conducted experiments similar to Rutherford’s using a
sheet of beryllium instead of gold.
• Unidentified high-energy radiation was given off. Chadwick was able to prove that the
radiation was made up of electrically neutral particles that weighed about as much as
protons.
• These electrically neutral subatomic particles are found in the nucleus of an atom,
and are called neutrons.
♦ In the early 1910s, Niels Bohr developed the idea that electrons circled the nucleus of an
atom in orbits called orbital.
♦
In 1926 Werner Heisenberg demonstrated that both the motion and the exact position of
an electron could never be known precisely at the same time. As a result, he proposed that
there are regions called energy levels where the electrons are most likely to be.
Section 2: The Mass of Atoms
1. The mass of an atom is measured using relative units called atomic mass units (amu).
♦ Protons and neutrons each have about the same mass—nearly 1 amu
♦
1 amu is equal to one-twelfth the mass of a carbon atom containing six protons and six
neutrons
2. Atomic mass – the sum of the relative masses of all of an atom’s protons and neutrons.
4. Atomic Number – the number of protons an atom has in its nucleus.
♦ The atomic number of each element is unique
3. Mass number – the atomic mass rounded to the nearest whole number. It indicates the
number of protons and neutrons in an atom.
♦ Because electrons have such a small mass they are ignored when calculating the mass
of an atom. It takes nearly 2000 electrons to equal the mass of one proton.
5. Although atoms of the same element always have the same number of protons, the number
of neutrons can vary.
♦ Isotopes – atoms of the same element that have a different number of neutrons
♦ Isotopes are indicated by writing the element’s name followed by the mass number
• Example: carbon-12 would have six protons and six neutrons, while carbon-13 would
have 6 protons and 7 neutrons
6. Dalton assumed that all the atoms of the same element would have the same mass, this is
not true. We know that atoms of the same elements can have different masses because they
can have a different number of neutrons.
♦ Average atomic mass – the average mass of all the isotopes for that element
Section 3: The Periodic Table
1. In 1869 Russian chemist Dmitri
Mendeleev proposed an arrangement of
the known elements based on their
atomic mass
atomic
number
state
8
symbol
G
O
Oxygen
16.00
2. The modern Periodic table differs from
Mendeleev’s.
♦ The current periodic table is
arranged in order of increasing
mass
atomic number
number
♦ In any square of the periodic table
you can find the following information: the symbol for the element, the atomic number,
and the average atomic mass. Also, some periodic tables will indicate the state of the
element as it is found in nature.
3. Elements are arranged in vertical columns known as groups or families.
4. The electron cloud structure
♦ In a neutral atom, the number electron equal the number of protons
♦ Electrons within the electron cloud have different amounts of energy
• The differences in energy determines the location of the electrons
• Energy levels nearer the nucleus have lower energy than those levels that are farther
away
• Energy levels are filled with electrons from the levels closest to the nucleus to the
levels farthest from the nucleus
• Elements in the same group have the same number of electrons in their outer energy
level
5. Periods – the horizontal rows of elements on the periodic table