Chapter 2 Atoms, Molecules, Ions. HW: 4 11 23 25 31 35 39 45 50 53 55 59 61 63 65 67 69 71 102. 2.1 – Atomic Theory. Dalton’s Atomic Theory Elements are made of extremely small particles called atoms Atoms of an element are identical Atoms are not created or destroyed
HW: 4 11 23 25 31 35 39 45 50 53 55 59 61 63 65 67 69 71 102
Dalton’s Atomic Theory
Law of Definite Proportions (ConstantComposition)
-A compound always has the same proportion of its elements
Law of Conservation of Mass (Matter) = Matter cannot be created or destroyed
Law of Multiple Proportions
-Two or more compounds with the same elements must have different proportions of the elements
-Example: Water vs. Hydrogen peroxide
Atom = Basic unit of an element that can enter into a chemical reaction
Subatomic Particle = Particles that make up an atom (protons, neutrons, electrons are as small as we will study)
-Cathode (Ray) Tube / Crooke’s Tube / = Glass tube w/2 metal plates. Connected to high voltage source. Emits ray.
Crooke – Determined that the ray was made of negative particles
JJ Thomson = 1897 – Credited w/ finding electrons Determined charge to mass ratio to be -1.76 108 coulombs/g.
Millikan = 1909 - Performed the Oil Drop Experiment. Found the charge of an electron (-1.60 x 10-19C). Could then calculate mass of the electron
Radioactivity = Spontaneous emission of radiation
a = + charge = Helium nucleus = High mass
b = - charge (high speed electrons) = Low mass
g = No charge. High energy = No mass
Protons = + charge particles in the nucleus
Rutherford = 1919- Gold Foil Experiment
-Most alpha particles when through, some deflected
This proved that Thomson’s model was incorrect
Results of Gold Foil Experiment:
Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.
-Most of the volume of the atom is empty space.
-We will only discuss protons, neutrons and electrons because they are the only ones that affect chemical behavior
-Charge on proton = +1.602 x 10-19 C. Assigned a +1 charge.
-Charge on an electron = -1.602 x 10-19 C. Assigned a -1 charge.
-Neutrons – Electrically neutral. Similar mass to protons. Discovered by Chadwick.
Mass of electron so small it is ignored
amu = atomic mass unit. 1 amu = 1.66054 x10-24 g. Based on 12 amu = mass of one atom of Carbon-12. 1 amu is approximately equal to the mass of a proton/neutron.
Angstrom (Å) = Unit used to measure atom size.
1 Angstrom = 1 x 10-10 m
2. Carbon - 12
All atoms of the same element have the same number of protons: The atomic number (Z)
The total number of protons and neutrons in the atom.
Average Atomic Masses – Based on abundance of isotopes and mass of each isotope. =Atomic Weight
Example 2.4 – Page 47
Periodic Table – Arranged by atomic number with elements with similar properties in same column.
Mendeleev – Organized first accepted Periodic Table
Metals are on the left side of the chart.
-Good conductors of heat and electricity
-Malleable, Ductile, Luster
-Readily lose electrons
Nonmetals are on the right side of the periodic table (with the exception of H).
-Readily gain electrons
Metalloids border the stair-step line (with the exception of Al and Po).
-Properties are intermediate between metals and nonmetals.
These five groups are known by their names.
Chemical Formula = Shows the atoms present in a substance
Molecule – Two or more atoms in a definite arrangement held by covalent bonds. Can be the same or different atoms.
ex – H2, H2O, C6H12O6
Diatomic Molecule = Molecule made of only 2 atoms total
Molecular Compound – Must have DIFFERENT elements
ex- H2O is a compound, but H2 is NOT
-Nonpolar Covalent – Equal sharing of electrons
-No partial charges
-Electronegativity of two atoms is close
-Difference of electronegativities is <0.3
-Polar Covalent – Unequal sharing of electrons
-Partial charges exist on atoms
-Electronegativity difference is 0.3-1.7
Molecular Formula = Gives actual number and type of atoms in a molecule
Empirical Formula = Simplest whole number ration of atoms
Example = Glucose
Structural Formula – Similar to a Lewis Structure. Shows bonds, but not shape.
Ionic compounds (such as NaCl) are formed by ionic bonds. Contain both + and – ions.
Ionic bonds are generally formed between metals (cations) and nonmetals (anions).
Monatomic – one atom only with a charge
Polyatomic – more than one atom with an overall charge
Names and Formulas of Ionic Compounds:
Zn+2 = Zinc ion
Ca+2 = Calcium ion
Cu+1 = Copper (I)
Cu+2 = Copper (II)
-Old system uses –ous / -ic
Names and Formulas of Ionic Compounds:
Example – bicarbonate = hydrogen carbonate
Ionic Compounds – Cation + Anion
-Binary – Compound of 2 “atoms”
-ex – NaCl, Al2O3
-Ternary – Compound of 3 “atoms”
-ex – Mg(NO3)2
Molecular Compounds: Discrete molecular units. Contain covalent bonds. Use prefixes.