Download Ch. 2. Atomic Structure and Periodic Table

Physical Science
Class Notes
Ch. 2. Atomic Structure and Periodic Table
Greek Model: Matter could not be divided into smaller and smaller units.
 Created by Democritus (Greek Philosopher)
 Greek word “atomos” means indivisible.
John Dalton (English chemist)
He worked with invisible gases and electricity.
Atomic Theory:
1. All elements are composed of atoms
2. Atoms are indestructible particles
3. Atoms of the same element are exactly alike
4. Atoms of different elements are different
5. Compounds are formed by the joining of atoms of two or more elements*
J. J. Thomson (English Scientist)
 Discovered electrons, but could not explain how the negative charge was balanced.
Ernest Rutherford (English physicist)
 Fired positively charged particles through a thin (2000 atoms thick) sheet of gold.
 Some particles passed through while others bounced back.
 Concluded the atom has a nucleus, which contains all the positive charges.
 http://micro.magnet.fsu.edu/electromag/java/rutherford/
Neils Bohr (Danish Scientist)
 Proposed electrons were located in specific areas and described these areas as orbits, similar to the orbits of
planets.
Wave Mechanics Model (Current)
 Electrons do not move around in specific pathways like orbits. Therefore, probabilities (not calculus) must be
used to study atoms.
Ex: Pauli Exclusion Principle: No two particles can occupy the same space.
Electron Cloud/Wave Mechanics Model
Nucleus:
 The center of the atom contains 99 % of the mass.
 100,000 times smaller than the entire atom.
 “Like a bee in a football stadium”
 Contains Protons: Positively charged particles in the nucleus.
 Contains Neutrons: Neutrally charged particles in the nucleus.
 Surrounded by Electrons: Negatively charged particles outside the nucleus.
Key Terms:
Atomic number: The number of protons in the nucleus of an atom.
Isotopes: Atoms of the same element with different numbers of neutrons.
*Isotopes are identified by including the mass number with the element name.
Ex: Carbon 14, Carbon 12
Average Atomic Mass: An estimate of the mass of an element’s atoms based on all the naturally occurring isotopes.
Atomic Mass Units (amu): The mass of a proton = 1 amu.
Mass Number: The sum of the protons and neutrons in the nucleus obtained by rounding the atomic mass.
Calculating the Number of Neutrons
*Determined by subtracting the number of protons from the mass number.
Ex: Uranium’s atomic number is 92. Subtract it from its mass number 238. Neutrons = 146
As an “Exit Ticket”, calculate the number of neutrons in the following atoms, plus 3 that you choose randomly.
C, N, O, H, He, S, and Pb. To obtain credit for the exit ticket, you must determine the number of neutrons in the nucleus
of the atoms of 10 elements.
Show following video pertaining to the newest proton accelerator: https://www.youtube.com/watch?v=sxAxV7g3yf8
Periodic Tables
Dmitri Mendeleev: Russian Chemist (1834-1907)
He classified 63 known elements according atomic mass, density, melting point, and valence.
Valence #: The number of electrons in the outer energy level that can be lost, gained, or shared during a chemical
reaction.
When he arranged the cards in order of increasing atomic mass and in 7 rows, he noticed an almost perfect
pattern related to valence # (1,2,3,4,5,6,7,8) and oxidation # (1,2,3,4,-3,-2,-1,0)
He shifted the pattern to make the elements align and proposed the missing pieces of the puzzle were places
for elements that had not been discovered.
Ex: Germanium (Ge): He predicted it would have an atomic mass of 72. The fact that its mass is 72.6
supported his version of the periodic table.
*As more discoveries were made, problems with Mendeleev’s table increased.
Henry Moseley: British Physicist (1887-1915)
Discovered the importance of the atomic number.
When the elements are arranged according to atomic number, rather than increasing atomic mass, the pattern
identified by Mendeleev for valence is consistent.
Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers.
Periodic Table Interpretation
Rows = Periods, Columns = Families/Groups
*The position of elements within periods and groups and colors convey lots of information.
Online Resources
Chemicool.com
Periodictable.com
Webelements.com
Periodic Table Challenge
Ptable.com
Periodic.lanl.gov/default.htm
Thatquiz.com
http://www.ilpi.com/genchem/periodicquiz.html
Exit Ticket: What are the valence number of O, N, He, Ne, C and B
Periodic Table Interpretation
Metals
Physical Properties:
 Have luster: Shininess
 Conduct heat and electricity
 High densities
 High melting points
 Ductile: Able to be drawn out into thin wire.
 Malleable: Able to be hammered out into thin sheets.
Chemical Properties: They tend to lose outer electrons. Therefore, they react with water and other chemicals in the
atmosphere through the process of oxidation/corrosion/tarnishing.
Alkali metals (Fam. 1)
 Include the elements of family 1 with the exception of hydrogen (H).
 They have 1 valence electron. Therefore they are very reactive.
 Rarely found as free elements
 Soft, silvery white and shiny
 High conductivity
 Samples must be stored in oil to prevent them from reacting with oxygen and water in the air.
 Reactions can be violent and due to the production of heat and hydrogen gas.
Alkaline Earth Metals (Fam. 2)
 Rarely found as free elements.
 Contain 2 valence electrons, so they are very reactive.
 Magnesium (Mg) is commonly combined with aluminum (Al) to make light-weight metal parts.
Transitional Metals (Fam. 3-12)





Highly conductive
Usually brightly colored and are commonly used in the production of paint.
Includes mercury (Hg) which has been used in thermometers.
Have valence numbers of 1 or 2.
Because they tend to share electrons from their second most outer shell, they easily bond with other
elements.
Rare-Earth Metals
Two rows arranged separately due to their valence electrons.
Lanthanide Series
Soft and white
High luster
High conductivity
Malleable
Use to make alloys and high quality glass.
Actinide Series
Radioactive: When changes occur in the nucleus, particles and energy are given off.
Non-Metals
Physical Properties:
 No luster
 Low conductivity
 Not ductile
 Not malleable
 Relatively low density
 Relatively low melting point
 Nonmetals can be noticeably different from each other.
Chemical Properties:
They tend to share or gain electrons during chemical reactions, or not react.
Boron Family (Fam. 13)
 Aluminum is the most abundant metal in earth’s crust, conducts electricity, and lightweight.
Carbon Family (Fam. 14)
 Carbon is the most abundant element and has a valence number of 4.
 Carbon is referred to as the basis of life due to the ~ 5 million compounds it forms. For this reason, we
are carbon based life forms.
 Organic chemistry is a branch of science that deals with carbon compounds.
Nitrogen Family (Fam. 15)
 A variable group used primarily for fertilizers.
Oxygen Family (Fam. 16)
 Oxygen is most abundant element in Earth’s atmosphere.
Halogen (Salt Former) Family (Fam. 17)
 Halogen bulbs contain bromine or iodine
 Can be identified by their colors
Noble Gases (Fam. 18)
 Have full outer energy levels and do not naturally form compounds
 Used for producing neon signs and laser beams.
Metalloids
Physical Properties
 Located along the dividing line between metals and nonmetals.
 All are solids.
 May or may not have luster.
 They conduct heat and electricity better than nonmetals, but not as well as metals.
 Ductile
 Malleable
Chemical Properties
Variable. They may gain or lose electrons during chemical reactions.
Exit Ticket: What are the valence number of O, N, He, Ne, C and B
Chemical Formulas: A shorthand way of representing compounds.
Ex:
NH3 = Ammonia
(1 atom of Nitrogen and 3 atoms of Hydrogen)
CO2 = Carbon Dioxide
(1 atom of Carbon and 2 atoms of Oxygen)
H2O = Water
(2 atoms of Hydrogen and 1 Atom of Oxygen)
C3H7OH = Rubbing alcohol
(3 atoms of Carbon, 7 atoms of Hydrogen, 1 atom of Oxygen and 1 more atom of hydrogen).
Chemical Equations: Combinations of chemical formulas used to describe chemical changes.
Ex: Burning charcoal
Heat
C + O2  CO2
Ex: Formation of water
H2 + O  H2O
Ex: Burning of methane
2CH4 + 4O2  2CO2 + 4H2O
2CH4 + 3O2  2CO + 4H2O
2CH4 + 2O2  2C + 4H2O
Law of Conservation of Mass
The mass of all substances present before a chemical change equals the mass of all the substances reaming after the
change.