1. Atomic Structure
Ch. 3

2. Early Philosophers Democritus – 400 BC Aristotle
Thought matter was made of individual, indivisible particles – “atomus”
Aristotle
Disagreed with Democritus
Thought matter was continuous
His idea was accepted until late 1700s

3. Beginnings of Modern Atomic Theory
1790s – Quantitative analysis begins with improved measuring equipment.
Chemists studied chemical reactions
Law of Conservation of Mass – Matter cannot be created or destroyed
Law of Multiple Proportions – different compounds with the same elements will have different ratios of those elements (CO and CO2)
Law of Definite Proportions – the same compound will have the same proportions by mass no matter what the sample size

4. Dalton’s Theory All matter is composed of atoms.
Atoms of the same element are identical.
Atoms cannot be divided, created or destroyed.
Atoms can combine in simple whole number ratios to form compounds.
Atoms are rearranged, separated, or combined in chemical reactions.

5. Modern Atomic Theory
Dalton helped turn Democritus’s idea into scientific theory.
All matter is composed of atoms.
Atoms of one element differ in properties from atoms of another element.

6. Discovery of the Electron
JJ Thomson used cathode ray tubes to study electricity and matter

7. Observations of the Cathode Ray Experiments
An object placed in the tube caused a shadow
A paddle wheel between the electrodes moved
Cathode rays were deflected with the negative end of a magnet and attracted by the positive end of a magnet
Video on Cathode Ray Tube

8. Conclusions of the Cathode Ray Experiments
The particles in the cathode ray have mass.
The particles are negatively charged.
The particles are very small.
Called Electrons

9. Discovery of the Nucleus
Rutherford used the Gold Foil Experiment to study the atom
Alpha particles are fired at the thin gold sheet.
Animation of the experiment

10. Observations of the Gold Foil Experiment
Some particles were deflected by the gold foil. Some at sizeable angles.
Most particles passed straight through the gold foil.

11. Conclusions of the Gold Foil Experiment
The nucleus is very small.
The nucleus is very dense.
The nucleus is positively charged.
Video of Rutherford’s experiment

12. Subatomic Particles Protons (p+) Neutrons (n0) Electrons (e-)
Positively charged, found in nucleus
Mass number of 1
Neutrons (n0)
Neutral, found in nucleus
Electrons (e-)
Negative, found in electron cloud surrounding nucleus, equals the number of protons in a neutral atom
Mass number of 0, Negligible mass

13. “Picture” of cesium and gallium atoms created using a scanning tunneling microscope

14. Subatomic Particles 3 quarks = 1 proton or 1 neutron Quarks He
component of protons & neutrons
6 types He
3 quarks = 1 proton or 1 neutron

15. 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”

16. Atomic Number Equals the number of protons Identifies the element
Found as the whole number on the Periodic Table
EX: Carbon has the atomic number of 6, so carbon has 6 protons in the nucleus.

17. Mass Number Equals the number of protons plus neutrons
NOT the mass found on the Periodic Table
Has to be provided to you by either
Hyphen Notation
Nuclear Notation

18. Nuclear Symbol Notation
Mass Number
(p+ + n0)
Atomic Number (p+ only, it is not necessary
to write the atomic number)
How many protons does this element have? 6
How many neutrons does this element have? 8
How many electrons does this element have? 6

19. Hyphen Notation Chlorine – 35 or Cl – 35 How many protons? 17
How many neutrons? 18
How many electrons? 17
Mass Number

20. Relative Atomic Mass Masses of atoms are very small
Ex: One O atom = x g
Carbon-12 has been assigned 12 atomic mass units
This is the standard for comparison
1 amu = 1/12th the mass of Carbon-12

21. C. Relative Atomic Mass atomic mass unit (amu)
12C atom = × g
atomic mass unit (amu)
1 amu = 1/12 the mass of a 12C atom
1 p = amu 1 n = amu 1 e- = amu
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22. Isotopes
Elements with the same number of protons but different number of neutrons.
Have different mass numbers.
Ex: Carbon-12 and Carbon-14
Both have 6 protons, but carbon-12 has 6 neutrons and carbon-14 has 8 neutrons.

23. Another Example of an Isotope
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

24. Average Atomic Mass Average mass of the isotopes found in nature
Mass found on Periodic table
Based on the percent abundance in nature
Example:
99.985% of hydrogen found in nature is hydrogen-1 and 0.015% is hydrogen-2.
This creates an average mass of amu.)

25. Example of Average Atomic Mass
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%

26. D. Average Atomic Mass Avg. Atomic Mass
weighted average of all isotopes
on the Periodic Table
round to 2 decimal places
Avg.
Atomic
Mass

27. D. Average Atomic Mass
EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.
Avg.
Atomic
Mass
16.00
amu

28. D. Average Atomic Mass
EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37.
Avg.
Atomic
Mass
35.40 amu

29. The Mole Chemist’s counting number SI unit for amount of a substance
Avogadro’s Number = x 1023
There are always x 1023 particles in one mole of any substance
Ex: 1 mole of Carbon atoms = x 1023 atoms
Chemists use a really big number because atoms are so small.

30. Molar Mass Mass of one mole of any substance
Relates the amount of atoms to their mass
Helps chemists know how many atoms are present – Remember atoms are too small to count!!!
Chemists can convert between mass (grams) and moles using the Periodic Table

31. Mass (Gram) to Mole Conversions
Molar mass is the conversion factor used to convert grams to moles
Found for every element on the Periodic Table – Round to two decimal places!!!!!
What is the molar mass of carbon?
12.01 g/ 1 mol
What is the molar mass of chlorine?
35.45 g/ 1 mol

32. Example of a Conversion
How many grams of lithium are in 3.50 moles of lithium?
Look up the molar
mass on the PT.
Write down given
3.50 mol Li
6.94 g Li
= g Li
45.1
1 mol Li
Multiply. Round
to correct sig figs.
Don’t forget the unit!
Determine which conversion factor is necessary.
Place the unit to cancel on the bottom.

33. Conversion Practice
Use dimensional analysis to solve these problems.
How many moles of carbon are present in 11.9 g of carbon?
What is the mass of 3.50 moles of copper?
How many atoms are present in 2.75 moles of lithium?
How many atoms are present in 95.5 grams of barium?