1. Atomic Structure and Periodicity

2. Chemistry Timeline #1  2000 years of Alchemy B.C.
400 B.C. Demokritos and Leucippos use the term "atomos”
 years of Alchemy
1500's
Georg Bauer: systematic metallurgy
Paracelsus: medicinal application of minerals
1600's
Robert Boyle:The Skeptical Chemist. Quantitative experimentation, identification of
elements
1700s'
Georg Stahl: Phlogiston Theory
Joseph Priestly: Discovery of oxygen
Antoine Lavoisier: The role of oxygen in combustion, law of conservation of
mass, first modern chemistry textbook

3. Chemistry Timeline #2
1800's
Joseph Proust: The law of definite proportion (composition)
John Dalton: The Atomic Theory, The law of multiple proportions
Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules
Amadeo Avogadro: Molar volumes of gases
Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements
Dmitri Mendeleyev: The periodic table
J.J. Thomson: discovery of the electron
Henri Becquerel: Discovery of radioactivity
1900's
Robert Millikan: Charge and mass of the electron
Ernest Rutherford: Existence of the nucleus, and its relative size
Meitner & Fermi: Sustained nuclear fission
Ernest Lawrence: The cyclotron and trans-uranium elements

4. Dalton’s Atomic Theory (1808)
All matter is composed of extremely small particles called atoms
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
John Dalton
Atoms cannot be subdivided, created, or destroyed
Atoms of different elements combine in simple whole-number ratios to form chemical compounds
In chemical reactions, atoms are combined, separated, or rearranged

5. Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties
Modern theory states:
Atoms of an element have a characteristic average mass which is unique to that element.

6. Modern Atomic Theory #2 Dalton said:
Atoms cannot be subdivided, created, or destroyed
Modern theory states:
Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!

7. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

8. 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.

9. Mass of the Electron
1909 – Robert Millikan determines the mass of the electron.
The oil drop apparatus
Mass of the electron is
9.109 x kg

10. Conclusions from the Study of the Electron
Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.
Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other particles that account for most of the mass

11. 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

12. 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?

13. The Answers
Target #1
Target #2
Rutherford Experiment Shockwave

14. 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

15. Atomic Particles Particle Charge Mass (amu) Mass (kg) Location
Electron -1
1/1000 = 0
9.109 x 10-31
Electron cloud
Proton +1 1
1.673 x 10-27
Nucleus
Neutron
1.675 x 10-27

16. The Atomic Scale
Most of the mass of the atom is in the nucleus (protons and neutrons)
Electrons are found outside of the nucleus (the electron cloud)
Most of the volume of the atom is empty space
“q” is a particle called a “quark”

17. 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”

18. Isotopes
Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Nucleus
Hydrogen–1
(protium) 1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium) 2

19. Atomic Number
Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.
Element
# of protons (A)
Atomic # (A)
Carbon 6
Phosphorus 15
Gold 79

20. Mass Number Nuclide p+ n0 e- Mass # (Z)
Mass number is the number of protons and neutrons in the nucleus of an isotope.
Mass # (z)= p+ + n0
Nuclide p+ n0 e-
Mass # (Z)
Oxygen - 10 - 33 42
- 31 15 18 8 8 18
Arsenic 75 33 75
Phosphorus 16 15 31

21. Format for writing isotopes of atomsA X Z
element symbol (X)
atomic number (Z) = number of protons
mass number (A) - atomic number (Z) = number of neutrons.

22. Bohr Model
Presented on overhead

23. Periodic Table with Group Names

24. The Properties of a Group: the Alkali Metals
  Easily lose valence electron
(Reducing agents)
  React violently with water
  Large hydration energy
  React with halogens to form salts

25. Predicting Ionic Charges
Group 1:
Lose 1 electron to form 1+ ions H+
Li+
Na+ K+

26. Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+

27. Predicting Ionic Charges
Loses 3
electrons to form
3+ ions
Group 13:
B3+
Al3+
Ga3+

28. Predicting Ionic Charges
Group 14:
Caution! C22- and C4- are both called carbide
Loses 4
electrons or gains
4 electrons

29. Predicting Ionic Charges
Nitride
Gains 3
electrons to form
3- ions
Group 15:
P3-
Phosphide
As3-
Arsenide

30. Predicting Ionic Charges
Oxide
Gains 2
electrons to form
2- ions
Group 16:
S2-
Sulfide
Se2-
Selenide

31. Predicting Ionic Charges
F1-
Fluoride
Br1-
Bromide
Gains 1
electron to form
1- ions
Group 17:
Cl1-
Chloride
I1-
Iodide

32. Predicting Ionic Charges
Stable Noble gases do not form ions!
Group 18:

33. Predicting Ionic Charges
Many transition elements
have more than one possible oxidation state.
Groups :
Iron(II) = Fe2+
Iron(III) = Fe3+

34. Predicting Ionic Charges
Some transition elements
have only one possible oxidation state.
Groups :
Zinc = Zn2+
Silver = Ag+

35. Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation and anion, including CHARGES!
Ba(NO3)2
2. Check to see if charges are balanced.
Ba2+
( )
NO3- 2
Not balanced!
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
4. Do not leave the charges in the formula.

36. Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation and anion, including CHARGES!
(NH4)2SO4
( )
NH4+
SO42-
2. Check to see if charges are balanced. 2
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!
4. Do not leave the charges in the formula.

37. Writing Ionic Compound Formulas
Example: Iron(III) chloride
1. Write the formulas for the cation and anion, including CHARGES!
FeCl3
Fe3+
Cl-
2. Check to see if charges are balanced. 3
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!
4. Do not leave the charges in the formula.

38. Writing Ionic Compound Formulas
Example: Aluminum sulfide
1. Write the formulas for the cation and anion, including CHARGES!
Al2S3
2. Check to see if charges are balanced.
Al3+
S2- 2 3
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!
4. Do not leave the charges in the formula.

39. Writing Ionic Compound Formulas
Example: Zinc hydroxide
1. Write the formulas for the cation and anion, including CHARGES!
Zn(OH)2
( )
2. Check to see if charges are balanced.
Zn2+
OH- 2
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!
4. Do not leave the charges in the formula.

40. Writing Ionic Compound Formulas
Example: Aluminum phosphate
1. Write the formulas for the cation and anion, including CHARGES!
AlPO4
2. Check to see if charges are balanced.
Al3+
PO43-
They ARE balanced!
3. Do not leave the charges in the formula.

41. Naming Ionic Compounds CaCl2 = calcium chloride
1. Cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride

42. Naming Ionic Compounds (continued) PbCl2 = lead(II) chloride
Metals with multiple oxidation states
some metal forms more than one cation
use Roman numeral in name
PbCl2
Pb2+ is the lead(II) cation
PbCl2 = lead(II) chloride

43. Naming Binary Compounds
Compounds between two nonmetals
First element in the formula is named first.
Second element is named as if it were an anion.
Use prefixes
Only use mono on second element -
P2O5 =
diphosphorus pentoxide
CO2 =
carbon dioxide
CO =
carbon monoxide
N2O =
dinitrogen monoxide