1. Atomic Structure Objectives: History of an Atom Atomic Models
Modern Atomic Theory
Ions
Mass of an Atom
Mass Number
Isotopes
Atom Energy
Wave – Mechanical Model
Electrons

2. The History
Democritus – a Greek philosopher that lived in 450 B.C. Said that all matter is composed (made) of tiny parts
He developed a definition of an atom – smallest particle of an element that retains the chemical properties of that element

3. History Continued Antoine Lavoisier – lived during the
1700’s and developed the law
of conservation of matter
*There is NO mass lost
during a chemical reaction

4. History Continued
Joseph Louis Proust – lived in 1799 and developed the Law of Constant Composition. A given compound always has the same elements in the same proportion by mass.

5. History Continued
John Dalton – Lived in 1803 and developed the first Atomic Theory of Matter.
1. Each element is made of very small parts called atoms
2. All atoms of a given element are identical but they differ from any other element
3. Atoms are not created nor destroyed in a chemical reaction
4. A given compound always has the same proportion and same type of atoms

6. Atomic Models
Michael Faraday – Lived in 1839 and showed how atoms contain particles that have an electrical charge.
*Rule of Charges
- Likes repel likes
- Opposites attract

7. Atomic Models
J. J. Thompson – Lived in 1897 and developed the “Plum Pudding Model”.
1. Atoms themselves are neutral
2. Atoms have electrons that have a
negative charge
3. Atoms have particles with a
positive charge

8. History Continued Rutherford – Lived in 1909 and
wondered about the positive and
negative charges. He used
the Gold Foil
Experiment to study how
alpha particles
interact with thin metal foil.

9. Gold Foil Experiment
Rutherford aimed a beam of high speed alpha particles at a piece of really thin (Au) Gold Foil.
He found that almost all particles passed through with no detection! Only 1 in 8000 did!
He then proposed that all atoms positive charge as well as the mass is concentrated in a small center core at the center of the atom. He called the center a nucleus and considered most of an atom empty space.

10. Apparatus used: Experiment Apparatus T – evacuation Tube
R- alpha particles from (Ra)
Radium
F – Foil
D – Diaphragm
M – Microscope
S – Screen

11. Modern Atomic Theory Atoms are made of three parts: Protons Neutrons
Electrons
These are subatomic particles

12. Modern Atomic Theory
Nucleus is the central core and contains protons and neutrons.
Protons have a positive charge (+)
Neutrons do NOT have a charge
(Zero charge)

13. Modern Atomic Theory Electrons: Move in space around the the nucleus
Electrons have a negative charge (-)
The outer most electrons are called valence electrons
*Note: the number of electrons in an atom can change, which changes the chemical properties of an atom.

14. Modern Atomic Theory
Atoms are neutral unless a change in electrons takes place
The number of protons = the number of electrons
Protons and electrons have equal but opposite charges
Protons have significantly more mass than electrons

15. Modern Atomic Theory Atomic Number
Atoms contain a unique positive charge located in the nucleus
The number of protons makes each element unique
Each element has a set number of protons and is used in element identification
Atomic number = the number of protons
The number of protons = the number of electrons in a neutral atom.
Example: Nitrogen (N) has 7e- and 7p+

16. Modern Atomic Theory
Gaining, losing or sharing electrons happens during chemical reactions and while dissolving compounds
When an atom loses or gains electrons, the atom becomes positive or negatively charged and is called an ION.

17. Ions An atom loses or gains an electron to become an ion.
1. Anion – an atom gains electrons and becomes negatively charged
2. Cation- an atom loses electrons and becomes positively charged
To find the net charge:
Protons – Electrons = Charge

18. Writing an Element Symbol with a Charge
Write the symbol and then show the charge as a superscript
Ex: Magnesium (Mg) lost 2 electrons and becomes an ion with 10 electrons and 12 protons.
12(p+) – 10(e-) = 2(p+)
The charge is +2. Mg then becomes Mg2+

19. Mass of an Atom Mass is measured in Atomic Mass Units (AMU).
AMU = # of Protons + # of Neutrons
Note: The average mass of an element is the weighted average of the masses of its naturally occurring isotopes

20. Mass Number The Mass Number identifies an isotope
Isotopes are represented by the chemical symbol and numbers of subscripts and superscripts
1. Mass number upper left
2. Atomic Number lower left

21. Isotopes
NOT all atoms of the same elements contain the same number of neutrons
Isotopes are atoms with the same number of protons but different number of neutrons but have the same chemical properties
Mass differs between isotopes, the more neutrons the more mass

22. Bohr Model
Electrons are arranged in shells (rings). Each shell represents a different energy levels and is assigned a number (n).
The lowest energy level is called the ground state.
***This is only a basic first model.

23. Energy Levels First energy level is n=1. This is the energy level
closest to the nucleus.
When an electron absorbs
energy it becomes excited
and jumps to higher energy
levels, n=2, n=3, n=4, etc.
This is called the
“excited state”.

24. Spectral Lines
Radiation (energy) is absorbed by an atom an electron jumps from the ground state to the excited state
When the electron returns to the ground state, radiation (energy) is emitted and gives off light. This can be used to identify an element.
The Bohr model is only good for hydrogen (H). It is a basic model but helped to form the current model for electron structure.

25. Energy Levels Explained
Notice how the electron (blue) jumps from the first shell (n=1), to the second shell (n=2)

26. Heisenberg’s Uncertainty Principle (1927)
Position and momentum of
moving objects cannot be
measured or known exactly
-We don’t know where
electrons actually are

27. Wave – Mechanical Model
Shows the relative location of electrons
1. It is represented by a cloud of negative charges around the nucleus
2. Based on high probability of finding an electron
3. The cloud mass is most dense where there is a high probability of finding an electron
these are the shells

28. Model of an Electron Cloud
Electrons move around the nucleus in a cloud.

29. Wave Mechanical Model
Uses Orbital to find relative location of electrons.
1. Probability is described by orbitals.
2. An atomic orbital is a region around the nucleus where an electron with a given energy is likely to be found.
3. Orbitals have certain shapes, sizes and energies.
4. Orbitals do NOT show how electrons move
5. The amount of energy an electron has determines the kind of orbital it occupies.

30. Orbitals and Energy
Energy and orbitals are assigned a certain value (quantized)
Principle energy levels are designated by the quantum number
Energy of an electron increases as the number (n) increases. N=1, n=2, n=3, etc.

31. Atom Encounters
When atoms come in contact with each other their electrons are what come in contact.
The interactions between the atoms are the basic principles of bonding. This only occurs in the valence shell. Valence shell is the outermost shell.

32. Practical Example
Cl has 17 protons (we know from the atomic number). Protons = Electrons, so it has 17 electrons. Where are they?
Hint: 1st shell can hold 2 electrons
2nd shell can hold 8 electrons
3rd shell can hold 18 electrons
4th shell can hold 32 electrons