1. Unit 8: Introduction to Chemistry
Atomic Theory, Atomic Calculations, and Electron Configuration

2. Atomic structure discovered
Ancient Greeks
Democritus ( BC) - indivisible particles called atomos
Prevailing argument (Plato and Aristotle) - matter is continuously and infinitely divisible
John Dalton (early 1800’s) –
reintroduced atomic theory to explain
chemical reactions

3. Dalton’s Atomic Theory (1800’s)
Every element is made up of tiny particles called atoms that cannot be subdivided.
Atoms of the same element are exactly alike.
Atoms of different elements can join in simple whole number ratios to form molecules.
Reactions occur when atoms are separated, joined, or rearranged.

4. Subatomic Particles
Electrons: negatively charged particles found outside of the nucleus
“Electron Cloud”
Protons: positively charged particles found inside of the nucleus
Neutrons: particles with no charge found inside of the nucleus

5. Discovery of the electron
J. J. Thomson (late 1800’s)
Performed cathode ray experiments
Discovered negatively charged electron
Identified electron as a fundamental particle

6. J.J. Thomson (1897)

7. Electron charge and mass
Robert Millikan (~1906)
Studied charged oil droplets in an electric field
Charge on droplets = multiples of electron charge
Charge + Thomson’s result gave electron mass

8. Early models of the atom
Dalton - atoms indivisible
Thomson and Millikan experiments
Electron mass very small, no measurable volume
What is the nature of an atom’s positive charge?
Thomson’s “Plum pudding” model
Electrons embedded in blob of positively charged matter like “raisins in plum pudding”

9. Thomson’s model

10. The nucleus Ernest Rutherford (1907)
Scattered alpha particles off gold foil
Alpha particles passed through without significant deflection
A few scattered at large angles
Conclusion: an atom’s positive charge resides in a small, positive nucleus
Later: positive charges = protons
James Chadwick (1932)
Neutral neutrons in the nucleus

11. Rutherford’s Gold-Foil Experiment (1911)

12. Distinguishing Among Atoms
Atomic number (Z) is the number of protons in the nucleus of an atom
Mass number (A) is the total number of protons and neutrons in an atom.
# of neutrons = A – Z
Atoms are identified using atomic number and mass number.
Example: 19779Au or gold-197

13. Isotopes Isotopes are different versions of the same atom.
They have the same number of protons and a different number of neutrons.

14. Atomic Mass The atomic mass is measured in atomic mass units (amu).
One amu = 1/12 of the mass of the carbon-12 isotope.
The average atomic mass is a weighted average of the masses of all the naturally occurring isotopes of each element.

15. Examples
98.93% of all of the Carbon earth is Carbon-12 with a mass of amu, and the remaining 1.07% is Carbon-13 with a mass of amu. Calculate the average atomic mass.

16. Examples
69.15% of all of the copper on earth is copper-63 with a mass of amu. The remaining 30.85% is copper -65 with a mass of amu. What is the average atomic mass of copper?

17. Mass of a Mole of an Element
Molar mass is the mass in grams of one mole of a representative particle
The atomic mass in grams is the mass of one mole of an element
For a compound, add the atomic masses in grams of each of the elements together. This is the mass in grams of one mole of a compound.

18. Examples Write the molar mass of the following:Na C
H2O
PCl3
H2O2
What is the mass of 2.00 moles of HCl?

19. Mole-Mass Relationship
Find the mass in grams of the following:
9.45 mol of Al2O3
4.52 x 10-3 mol C20H42
2.50 mol of iron (II) hydroxide

20. Mole-Mass Relationship
Find the amount in moles of the following:
92.2 g of Fe2O3
3.70 x 10-1 g of boron
75.0 g of dinitrogen trioxide

21. The Bohr Model (1913)
1913 – A Danish scientist Niels Bohr suggested that the electrons revolve around the nucleus like the planets revolve around the sun.
In Bohr’s model each electron has a certain amount of energy that is defined by its path around the nucleus. These defined orbits are called energy levels.

22. Bohr’s theory Planetary Model: Three rules:
Electrons orbit the nucleus like planets orbit the sun
Electrons can “jump” between orbits by absorbing or emitting photons
Theory explained the line spectra of H
Three rules:
Electrons only exist in certain allowed orbits
Within an orbit, the electron does not radiate
Radiation (light) is emitted or absorbed when changing orbits

25. Atomic Spectra
When electrons absorb energy they move to higher energy levels, and when electrons lose energy by emitting light they return to their ground state.

26. Hydrogen Line Spectrum

27. Quantum theory of the atom
Lowest energy state = “ground state” n1
In hydrogen, n1= eV
Higher states = “excited states”
n2, n3, n4…
Energy of orbit calculated by
Photon energy equals difference in state energies
1 eV=1.6E-19 J

28. The Energy of a Photon
The energy of a photon is proportional to the frequency of light and can be calculated using Plank’s constant
E = hf, where h = 6.63 x Js
How much energy would red light have with a frequency of 4.57 x 1014 Hz?

29. How does this tie in with Bohr’s Model
The light emitted from an electron moving from a higher to a lower energy level has a frequency proportional () to the energy change of the electron moving between energy levels

30. An electron in a hydrogen atom jumps from the exited energy level n = 4 to n = 2. What is the frequency and wavelength of the emitted photon?

31. Electron Filling Order
As you move down the periodic table each atom contains more electrons.
There are two ways of writing electron configuration.
The first is spectroscopic notation.
Example: Hydrogen atomic number 1
1s1
Energy level
Orbital
Number of electrons

32. Atomic Orbitals

33. Electron Filling Order
The second is orbital box diagrams.
Example: Hydrogen atomic number 1
___ 1s
Energy Level (n)
Orbital (l)
Electron spin (+1/2 or –1/2)

34. Writing electron configurations
Electrons fill available orbitals in order of increasing energy
Shell capacities
s = 2
p = 6
d = 10
f = 14

35. Writing electron configurations
Rules:
Determine how many electrons element has
Fill in according to Aufbaugh chart
Shorthand:
Find the previous Noble gas
Put symbol in brackets
Complete e- configuration from that point
Example: Strontium
[Kr] 5s2

36. Examples
Use the orbital box diagrams and draw out the electron configuration for carbon in its ground state.
Also write out carbons electron configuration using spectroscopic notation.

37. Examples
Use the orbital box diagrams and draw out the electron configuration for vanadium in its ground state.
Also write out vanadium electron configuration using spectroscopic notation.

38. Periodic Table
Elements in the same group have similar properties, because they have the same number of valence electrons.
The columns are also arranged on the periodic table according to the electron orbitals of the different atoms.