1. ATOMIC STRUCTURE
Chapter 4

2. Atom Review Atoms – the fundamental building blocks of matter.
Major Subatomic Particles
Particle
Charge
Location
Mass
proton +1
nucleus
1 u
neutron 1u
electron -1
electron cloud
1/1837 u

3. electrons in shells
nucleus

4. Nucleus – dense central region of atom; contains essentially all of the atom’s mass.
Atomic Number (Z) – the number of protons in the nucleus. Each element has a different atomic number.
The atomic number identifies the element.
Isotopes – atoms of the same element that have different masses. Isotopes have different masses because they have a different number of neutrons.

5. Mass Number – the number of protons AND neutrons in the nucleus.

6. The isotopes of a particular element have different mass numbers:
Atomic Number:
Mass Number:

7. U-238 Most common isotope of Uranium
You can write the mass numbers of isotopes 2 ways:
1) Write the name or symbol of the element followed by a hyphen and the mass number.
Hydrogen-1 or H-1
Carbon-12 or C-12
Carbon-13 or C-13
Hydrogen-2 or H-2
U-238 Most common isotope of Uranium
(99.3 %)

8. 235U Fissionable isotope of Uranium
2) Write the mass number as a superscript to the left of the chemical symbol:
1H 2H
12C 14C
235U Fissionable isotope of Uranium
Sometimes the atomic number is added as a subscript:
14C 6
235U 92

9. *number of electrons = number of protons in neutral atom
Helium-4 isotope
Shell
proton N + - + N -
neutron
electron
*number of electrons = number of protons in neutral atom

10. 2 electrons (electrons = number of protons in neutral atom)
ATOMIC STRUCTURE He
the number of protons and
neutrons in an atom 4
Mass number
the number of protons in an atom 2
Atomic number
2 electrons (electrons = number of protons in neutral atom)

11. Ions have a different number of electrons than the neutral atom.
Negative ions have additional electrons (equal to their charge).
EX: O has 8 electrons; O2- has 10 electrons
Positive ions have fewer electrons (equal to their charge).
EX: Al has 13 electrons; Al3+ has 10 electrons

12. Symbol: Sr P
Charge: +2
Atomic #: 20
protons: 16
electrons: 18
neutrons: 50
Mass #: 33 31

13. Atomic Mass
Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units (u).
Therefore, 1 u = 1/12 the mass of a carbon-12 atom.
1 u is approximately the mass of a single proton or neutron.

14. Information in the Periodic Table
The number at the bottom of each box is the average atomic mass of that element.
This number is the weighted average mass of all the naturally occurring isotopes of that element.

15. Average Atomic Mass Almost all elements occur as a mixture of isotopes
The percentage of each isotope is a constant
EX: Hydrogen is composed of three isotopes: Isotope Atomic Mass Natural Abundance
H %
H %
H trace

16. The atomic mass listed on the periodic table for each element is a weighted average of the masses of the isotopes for that element.
A weighted average takes into consideration the percentage of each isotope.
Calculate the average atomic mass of hydrogen using the previous given isotopic data.

17. Calculating Average Atomic Masses
EX: Chlorine is composed of two isotopes: Isotope Atomic Mass Natural Abundance
Cl %
Cl %
Calculate the average atomic mass of chlorine.

18. ***To calculate average atomic mass:***
a) Multiply the atomic mass of each isotope by
its percentage (don’t forget to divide percentage by 100)
b) Add results of part “a” together to get average
EX2: Copper is composed of two isotopes: Isotope Atomic Mass Natural Abundance
Cu %
Cu %
Calculate the average atomic mass of copper.

19. Calculate the average atomic mass of carbon.
EX3: Carbon consists of the following isotopes: Isotope Atomic Mass Natural Abundance
C %
C %
C (trace amount)
Calculate the average atomic mass of carbon.

20. Using the information below, calculate the average atomic mass of element X:
Isotope Percent Abundance Mass (u)
26X
24X
22X