1. Mass Spectrometer

2. The Mass Spectrometer and Atomic Mass
* The actual mass of atoms was not known until the invention of the mass spectrometer.
* In a mass spectrometer, ions of high mass deflect less than those of low mass.
* If the charge of the particle is known, then the mass of the particle can be calculated.

3. Atomic Terms
Atomic Mass Unit (amu): exactly equal to 1/12 of the mass of a carbon-12 atom. ( x 10-23g)
1 amu = x 10-24g)
Isotopes: Atoms of an element having different atomic masses.
Atomic Mass: The average relative mass of the isotopes of an element compared to the mass of carbon – 12.

5. Isotope
Isotopic mass (amu)
Abundance
(%)
6329Cu
69.09
6529Cu
30.91
Atomic mass is then calculated in the following way:
( )(0.6909) =
( )(0.3091) =
63.55

6. Particle
Symbol
Relative Electric
Charge
Relative Mass (amu)
Actual Mass
(g)
Electron e- -1
1/1837
9.110 E-28
Proton p +1 1
1.673 E-24
Neutron N
1.675 E-24

7. Composition of the nucleus
Atomic Masses
Atomic mass is the average of all the naturally isotopes of that element.
Carbon =
Isotope
Symbol
Composition of the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
7 neutrons
1.11%
Carbon-14
14C
8 neutrons
<0.01%

8. The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons)
Electrons are found outside of the nucleus (the electron cloud)
Most of the volume of the atom is empty space
“q” is a particle called a “quark”

9. About Quarks… Protons and neutrons are NOT fundamental particles.
Protons are made of two “up” quarks and one “down” quark.
Neutrons are made of one “up” quark and two “down” quarks.
Quarks are held
together by “gluons”