Ch 3 - The Evolution of Atomic Theory Flashcards
Leucippus and Democritus ~450 BCE
Defined the atom as the point at which matter can no longer be subdivided. Matter was made if tiny, discrete indivisible particles called atoms.
Aristotle ~384 BCE
Proposed that all matter is composed of four elements and all matter is a continuous (infinitely divisible), not atomistic
Law of conservation of mass
When a reaction takes place, matter is neither created or destroyed
Law of constant proportion
Multiple samples of any pure compound always contain the same percent by mass of each element making up the compound
Percent by mass equation
% mass = mass of element x 100
——————–
total mass of compound
Law of multiple proportions
Elements may combine in more than one set of proportions, with each set corresponding to a different compound
Dalton’s atomic theory
1) All matter is made up of atoms
2) Atoms can neither be created not destroyed
3) Atoms of a particular element are alike
4) Atoms of different elements are different
5) A chemical reaction involves the union or separation of individual atoms
Modern modifications of Dalton’s atomic theory
1) Dalton assumed atoms to be indivisible. Not quite true (electrons, neutrons, protons).
2) Dalton assumed that all the atoms of a given element were identical in all aspects, but we know ions are an exception to this rule.
3) The numbers of each kind of atom in simple compounds usually form a simple ratio.
4) Unmodified for chemical reactions. Atoms are broken apart in nuclear reactions.
Ball and hook model
Different size balls represent different atom types; different types have different numbers of hooks representing bonding
Discovery of the electron
1897: J. J. Thomson determined the charge using the mass ratio if cathode ratio
Electron was very small and lightweight (1/1836 mass of H atom)
Discovery of the proton
1907: J. J. Thomson and E. Goldstein
Much heavier than the electron (~same mass as a H atom)
Discovery of the neutron
1932: James Chadwick
Electron mass
9.1 x 10^-28 g
Plum-pudding model
Positive “pudding” base with free-floating negative electrons in it
Rutherford’s gold foil experiment
Fired alpha particles at metal foils and a glass substrate covered in zinc sulfide to monitor the alpha particles. Surprised by the large angles of deflection by some if the particles (disproved the plum-pudding model)
Characteristics of a nucleus
1) Atom contains dense center called nucleus
2) Nucleus is essential to the entire Aron’s mass
3) Nucleus is positively charged
4) Electrons are disbursed in the empty space if the atom surrounding the nucleus
Proton mass
1.67262 x 10^-24 g
Proton charge
+1.60 x 10^-19 C
Electron charge
-1.60 x 10^-19 C
Atomic number (Z)
The number of protons in the nucleus; determines the identity of the atom
Mass number
Protons + neutrons
Isotopes
Atoms with the same number of protons (same element), but different numbers if neutrons (different mass numbers).
Exhibit identical chemical properties.
Isotope symbol
Shows the mass number and the atomic number (often omitted)
Atomic mass
The actual mass of any atom
Atomic mass units (amu)
1 amu = (1/12) the mass of a carbon-12 atom
1 amu = 1.66054 x 10^-24 g
Relative atomic mass
Measures how massive an atom is in comparison to a carbon-12 atom
Weighted average of atomic mass
Weighted average of the isotope masses
Law of Octaves
Elements that are eight elements apart by mass react in similar manners.
Also known as chemical periodicity or periodic behavior.
Law of Mendeleev
Properties of the elements recur in regular cycles (periodically) when elements are arranged in order of increasing atomic mass.
Periods
Horizontal rows on the periodic table if elements
Groups
Vertical columns of the periodic table of elements
Main-group elements
Groups 1,2, 13-18
Noble gases
Group 18
Transition metals
Groups 3-12
Lanthanides
Top row if the lower section of the table (not numbered)
Actinides
Bottom row if the lower section of the table (not numbered)
Metals
Metallic luster, conduct heat and electricity, malleable, and ductile
(Ex: sodium and copper)
Non metals
Dull luster, good insulator, nonconductors, and brittle in the solid state.
(Ex: sulfur and bromine)
Metalloids
Demonstrate properties of both metals and nonmetals; used as semiconductors.
(Ex: silicon and arsenic)
Group names
Group 1 (1A): alkali metals Group 2 (2A): alkaline earth metals Group 16 (6A): chalcogens Group 17 (7A): halogens Group 18 (8A): noble gases
Alkali metals
Lithium (Li) Sodium (Na) Potassium (K) Rubidium (Rb) Cesium (Cs) Francium (Fr)
Alkaline earth metals
Beryllium (Be) Magnesium (Mg) Calcium (Ca) Strontium (Sr) Barium (Ba) Radium (Ra)
Halogens
Fluorine (F) Chlorine (Cl) Bromine (Br) Iodine (I) Astatine (At)
Chalcogens
Oxygen (O) Sulfur (S) Selenium (Se) Tellurium (Te) Polonium (Po)
Noble gases
Helium (He) Neon (Ne) Argon (Ar) Krypton (Kr) Xenon (Xe) Radon (Rn)
Law of octaves adjustments
Works well only when considering the main-group elements
Ions
Charges atoms or groups if atoms
Anions
Negatively charges ions (gain electrons)
Cations
Positively charged ions (loses electrons)
First ionization energy
Minimum amount of energy it takes to completely remove an electron forming a 1+ ion.
Exhibits periodicity similar to chemical properties. Metals exhibit lower values than nonmetals
