1. Atomic Theory

2. Greek Philosophers: 4 Elements – earth, air, fire and water

3. Democritus: world made of two things – empty space and tiny particles called “atomos” = unable to cut
atoms are the smallest particles and each substance had its own type of atom
- wood atoms, air atoms, water atoms

4. Dalton: 1. all matter is made of tiny particles called atoms 2
Dalton: 1. all matter is made of tiny particles called atoms 2. atoms can’t be broken down further 3. atoms of different elements differ 4. atoms of the same element are identical 5. atoms combine to form compounds in specific ratios and can be rearranged to make new compounds

5. Atomic Composition: Nucleus surrounded by an electron cloud nucleus contains protons and neutrons Charges: Protons = positive (+) Electrons = negative (-) Neutrons = neutral (0) Mass: Protons = × kilograms or 1 amu Neutrons = × kilograms or 1 amu Electrons = × kilograms or 0 amu

6. Breaking Down the Nucleus: Protons and Neutrons Protons: - the number of protons in the atoms of an element is always the same - changing the protons will change the type of element The Number of Protons = The Atomic Number The Number of Protons and Neutrons = The Atomic Mass

7. At. Mass Element Symbol At. #
Element Symbols Symbol, Atomic Mass, Atomic Number 12 C 6 - information: element, # of protons (and electrons by implication), # of neutrons
At. Mass
Element Symbol
At. #

8. Differences Among Atoms of the Same Element IONS: in electrically neutral atoms, # protons = # of electrons loss or gain of electrons  ion Cation = (+) ion = loss of electrons Anion = (-) ion = gain of electrons

9. Ex: Sodium – atomic # = protons, 11 electrons - lose 1 electron = 11 p+ and 10 e- = (+1) charge Chlorine – atomic # = protons, 17 electrons - gains 1 electron = 17 p+ and 18 e- = (-1) charge

10. Calculating Atomic Mass and Charge
Atomic Number = # Protons
Atomic Mass = Protons + Neutrons
Atomic Charge = Protons + Electrons

11. EX:
An atom has an atomic number of 12, 13 neutrons and 12 electrons.
Identity:
Atomic Mass:
Charge:

12. EX:
An atom has an atomic mass of 35, 18 neutrons and 18 electrons.
Identity:
Protons:
Charge:

13. EX:
An atom has 20 protons, 21 neutrons and 18 electrons.
Identity:
Atomic #:
Atomic Mass:
Charge:

14. ISOTOPES: same number of protons, different numbers of neutrons changes the mass of the atom

15. ATOMIC MASS = # of protons and neutrons each has about the same mass which is designated as an atomic mass unit (1 amu) Determination of amu one element chosen as a standard – Carbon 12 6 protons, 6 neutrons = 12 amu therefore 1 amu = 1/12th Carbon atom

16. HOWEVER: Is the Atomic Mass of an Element a whole number. NOPE WHY
HOWEVER: Is the Atomic Mass of an Element a whole number? NOPE WHY? The Atomic Mass on the Periodic Table is the AVERAGE ATOMIC MASS OF ALL THE KNOWN ISOTOPES AND THEIR ABUNDANCE

17. Average Atomic Mass:
Mass Spectrometer:
Sample of gaseous element
Charged
Propelled by electromagnetic fields toward a photographic plate which records how much is present
Less massive fall shorter
More massive fall farther
Allows for determination of relative mass
Masses are then combined and averaged  decimal

18. Mass Spectrometer

19. Calculating Average Atomic Mass: [(# of atoms X mass of isotope A) + (# of atoms X mass of isotope B) ] divided by (total number of atoms of all isotopes combined) OR (Mass of Isotope A X Relative Abundance) + (Mass of Isotope B X Relative Abundance) = Average Atomic Mass

20. Boron has two naturally occurring isotopes: Boron-10 (abundance = 19
Boron has two naturally occurring isotopes: Boron-10 (abundance = 19.8%, mass = amu) and Boron-11 (abundance = 80.2%, mass = amu. Calculate the atomic mass of boron.

21. Calculate the atomic mass of magnesium
Calculate the atomic mass of magnesium. The three magnesium isotopes have atomic masses and relative abundances as follows:
amu (78.99%)
amu (10.00%)
amu (11.01%)

22. Particle Physics Part I Particle Physics Part II

23. Radioactive Decay Unstable Nuclei: Radioactive Atoms - Radioisotopes
atoms are unstable because they are high in energy
atoms give off the energy (radiation) to become more stable
Process of losing the energy is radioactive decay
Atoms can actually become other elements

24. Radioactive Decay
Alpha particles – positively charged Helium Nucleus Alpha particle decay: Unstable nucleus loses an alpha particle – result, atom loses two protons and two neutrons mass decreases by 4 amu atom becomes another element U  Th + He

25. Radioactive Decay
Beta particles – negatively charge particle (electron) Beta particle decay: results from the break down of a neutron into a proton and an electron - atom becomes another element n  p + e Th  Pa + e

26. Radioactive Decay
Gamma waves – high energy (no matter, no charge) Gamma Ray Emission: following Beta particle decay the nucleus still has energy so the nucleus releases it as gamma rays - both the atomic mass and number remain the same 230m 230 Th  Th + γ 90 90
Means lots of energy
Gamma symbol

27. Radioactive Decay
Positron Emission – release of a positive electron (positron) the nucleus does not lose mass but does decrease in atomic number – the mass of a positron is nearly zero N  e + C

28. Radioactive Decay
Electron Capture – rare instance where an electron runs into the nucleus – the electron combines with a proton to form a neutron – therefore the mass does not change but the atomic number does Ca + e  K

29. Calculating Nuclear Decay: Half-life
time required for half of the sample to decay to the products
Ex: Uranium-238 takes 4.47x109 years
If you had a 10.0g sample, it would take 4.47x109 years for it to decay to 5.0g
It would take 4.47x109 more years to degrade to 2.5 g
It would take 4.47x109 more years to degrade to 1.25 g…….
Ex: Substance X has a half life of 10 years.
Half Life 1 2 3 4
Time 10
10 more
20 total
30 total
40 total
Percent
Remaining
100%
50%
25%
12.5%
6.25%
Fraction
1/8
1/16

30. n = number of half lives that have passed
Equations:
Amount Remaining = (Initial amount)(1/2)n
n = number of half lives that have passed
Amount Remaining = (Initial amount)(1/2)t/T
t = elapsed time
T = duration of half-life

31. Ex: Radioactive iodine-131 has a half-life of 8
Ex: Radioactive iodine-131 has a half-life of 8.04 days 1) If you have 8.2 ug (micrograms) of this isotope, what mass remains after 32.2 days? 2) How long will it take for a sample of iodine-131 to decay to 1/8 of its activity?