1. Unit Six: Atomic structure
How small is an atom?

2. Evidence of Atomic Structure
An understanding of how matter behaves and the understanding of atomic structure did not develop until scientists were able to measure Mass.
Use of a balance to measure the mass of materials and to monitor chemical reactions, lead to the three basic properties of matter:

3. The Law of Conservation of Mass
“ Matter is neither created or destroyed in a chemical reaction”.
Which means:
the mass of the substances produced (products) in a chemical reaction is always equal to the mass of the reacting substances (reactants).

4. Law of Constant Composition
In a pure substance, the elements are always present in the same proportion by mass.
For example:
Water will always be composed of two Hydrogen atoms : 1 oxygen atom.
That means: 12% of the mass is Hydrogen and 88% of the mass is oxygen….
It doesn’t matter how much you have a cup full or an ocean worth of water! H2O

5. Law of Multiple Proportions
When two elements, A and B, combine to form more than one compounds, the mass of A and B will combine in simple ratios.
Mass measurements would reflect the 1:2 ratio of carbon and oxygen in CO2 and the 1:1 ratio in CO.

6. Prove it: Law of Conservation of Mass
Objective:
Conduct an experiment to see if the law of conservation of mass is true.

7. Applying laws of matter to chemistry:
Balancing equations.
Naming compounds and writing formulas
Chemical analysis.

8. Historical Understanding of Atomic Structure Important Scientists and their contributions to atomic theory:
John Dalton:
Used evidence from many experiments (change in mass/balances)
Identified basic properties of elements and how atoms interact and form compounds.
All elements consist of atoms that cannot be divided.
Atoms of different elements have different masses.
Atoms of one element cannot be changed into atoms of another element by chemical reactions.
Compounds are formed when atoms of more than one element combine in specific ratios.

9. Important Scientists and their contributions to atomic theory:
JJ Thompson
Used a cathode ray tube (vacuum tube containing a gas)
Discovered that atoms contain electrons (negatively charged particle)
He thought that electrons were scattered throughout the atom along with positive charge particles, “plum pudding” model.

10. Thomson’s Atomic Model
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

11. Important Scientists and their contributions to atomic theory:
Ernest Rutherford:
Gold Foil Experiments
Discovered the positively charged nucleus of the atoms.
He suggested that the atom is mostly empty space but has a positive charge in the center.
Rutherford called the positively charged particles protons.

12. Rutherford’s Gold Foil Experiment
Alpha particles are helium nuclei
Particles were fired at a thin sheet of gold foil
Particle hits on the detecting screen (film) are recorded

13. Try it Yourself!
In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?

14. The Answers
Target #1
Target #2

15. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

16. Important Scientists and their contributions to atomic theory:
Neils Bohr:
Suggested that electrons are found in orbits around the nucleus.

17. Atomic Structure and Cloud Model:
More recent research suggests that electrons move rapidly within cloud-like regions called orbitals:
James Chadwick latter identified that the nucleus also contains neutrons which have no charge)

18. The Structure of the Atom
Subatomic particle
Location
Charge
Property
Proton
Nucleus +
Contributes mass
Little space
Neutron
Neutral
Electron
Orbital -
Little/no mass
Provides space

19. Understanding Nuclear Symbols
Atomic Number (Z)- the number of protons in the nucleus.
Atomic Mass (A)- the number of protons (Z) + neutrons (N) in the nucleus.

20. Practice: Name= Chlorine Atomic Number (Z) = 17 Atomic Mass = 35
# Protons = 17
# Neutrons = 18 (35-17 = 18)
# Electrons 17

21. Atomic Structure & Counting Atoms
Atomic number- defines the number of protons in the nucleus of an atom
Atomic mass - the mass (weight) of an atom
Mass Number- the sum of the number of protons and neutrons in the nucleus of an atom.
Element
Symbol
Atomic Number
Atomic mass
Mass Number
# of protons
# of electrons
# of neutrons
Aluminum Al 13
26.98 27 14

22. Changing Atoms
Ions: electrically charged atoms formed by the loss or gain of electron(s).
Isotopes: atoms of the same element that have a different atomic mass. (protons remain the same but the neutrons are different).

23. Changing Atoms: Ion Atom Isotope Add/remove electrons
Add/remove neutrons

24. Ions:
Cation: positively-charged (+) ion formed when an atom loses electrons
“I positively love cats”
Anion: negatively-charged (-) ion formed when an atom gains electrons

25. Counting Ions: Sodium Ion Protons = _____________
Neutron = ____________
Electrons = _____________

26. Practice: Ions Ca2+ Cl- K+ O2- Element/ Ion Atomic Number Mass Number
Protons
Neutrons
Electrons
Ca2+ 20 40 18
Cl- 17 35 K+ 19 39
O2- 8 16

27. Isotopes and Average Atomic Mass:
Isotopes are atoms that have the same atomic number but different mass number. This means that isotopes have the same number of protons but a different number of neutrons.
Remember that protons and neutrons are located in the nucleus of the atom.
The average atomic mass descries the average mass of all the naturally occurring isotopes for an element according to their natural abundance.

28. Calculating the Average Atomic Mass:
The average atomic mass is calculated by added the contributing mass of each isotope based on its natural abundance (% amt.)
Chlorine (Cl) has two naturally occurring isotopes:
35.0 amu 75.5%
37.0 amu 24.5%
Average = (35.0 x .755)+ (37.0 x .245) = 35.5amu

29. Practice: Ions & Isotopes
Element/Ion
Atomic Number
Atomic Mass
Mass number
Protons
Neutrons
Electrons H 1
1.01 H+
612C
37Li+
Cl-