Chemistry Review-Unit I-Atomic Structure
Important Reference Tables: Table O, Centerfold
History
Dalton: atoms are indivisible particles
Thompson: Atoms have + and – particles spread through them (plum pudding)
Rutherford: the atom is mostly empty space, with a small, positive nucleus
Bohr: Electrons orbit the nucleus in energy levels
Modern: cloud/probabity wave model for electrons (orbitals and sublevels)
Nucleus
Protons: positive (+) charge, mass = 1 amu
Neutrons: neutral (0) charge, mass = 1 amu
Most of the mass of an atom is in the nucleus
Rutherford’s Experiment: Shot alpha particles through gold foil.
Proved the nucleus is extremely small-an atom is mostly empty space
Atomic Number: = protons (and electrons if the atom is neutral)
Mass Number: = protons + neutrons
Charge (net charge): = protons – electrons
(assume net charge = 0 unless given different information)
Nuclear Charge: = number of protons in the nucleus
Isotopes: atoms of the same element with different mass numbers
Same protons, different neutrons
Example: 126C and 146C…one has 6p and 6n, the other has 6p and 8n
Nuclear Symbols: The bottom number is protons (charge)
The top number is mass (protons = neutrons)
Calculate neutrons by subtracting
Average Atomic Mass: Different from mass number. Read on centerfold
The weighted average of all naturally occurring isotopes of the element
Shells
Bohr Model: seven shells. Electrons fill the lowest available shells first
Orbitals: the pathways for a electrons
Shell Capacity: 2n2 n= shell number
Ground State: all electrons are at their lowest available shells
Excited State: at least one electron is higher than necessary
Electrons gain energy to jump up, lose energy to drop down
Electron Configuration: centerfold.
Valence Electrons: Outer shell electrons
Chemistry Review-Unit II-The Periodic Table
Important reference tables: Centerfold, table S
Mendeleev: arranged the elements by increasing mass, with similar properties in the same group
Modern: arranged by increasing protons, with similar elements in the same group
Groups: column, family, elements with similar properties
Periods: Rows, elements with the same number of occupied shells
Staircase: Divides metals (left) from the nonmetals (right)
Metalloids touch the staircase (except Al, Po, and At)
Noble gases are NOT nonmetals
Metals vs. Nonmetals
Valence: metals have one or two, nonmetals have five or six
Electronegativity: metals are low, nonmetals are high
Ionization energy: metals are low, nonmetals are high
Radius: Metals are large, nonmetals are small
Reactivity: Metals lose electrons and become small positive ions
Nonmetals gain electrons and become large negative ions
Francium is the best metal
Fluorine is the best nonmetal
Specific Group Properties
Alkali Metals: group 1. Most active metals. Never exist in nature as metals
Usually exist in salts
Alkaline Earth Metals: group 2. Also active metals. Never exist in nature as metals
Usually found in minerals
Transition Metals: groups 3-11, Less reactive metals. Multiple oxidation states.
Form colored compounds and solutions.
Carbon: bonds four times. Important for organic chemistry
Nitrogen: Forms a triple bond with itself.
A Stable element that forms unstable compounds
Oxygen: forms a double bond with itself. Ozone = O3
Halogens: Group 17. The best nonmetals.
Noble Gases: Group 18. Nonreactive (inert) elements. All monatomic Molecules
Kr and Xe can react with Fluorine under extreme conditions
Miscelaneous
Liquid Elements: Mercury and Bromine
Unstable (radioactive) elements: Bottom of the chart. More than 83 protons
Unit III: Chemical Formulae and Equations
Important reference tables: centerfold, table E, table T
Formulae
Molecular Formula: subscripts tell the actual number of atoms in the molecule
Empirical formula:simplest formula. Atoms are reduced to their lowest termsC2H6=CH3
Oxidation Numbers: must add up to zero. Sodium Oxide = Na+1 O-2 = Na2O
Names
Binary Compounds: Two elements only. Name metal first. Change suffix to “ide”
Most compounds ending with “ide” are binary compounds.
(sodium chloride, potassium bromide, carbon dioxide)
Nonbinary compounds: Must contain a polyatomic ion (table E)
(Sodium sulfate, ammonium hydroxide)
Roman Numerals: tell the charge of the first element (always positive)
Iron (III) chloride = FeCl3, Nitrogen (II) oxide = NO
Molecular Mass (Formula Mass): add the atomic masses, CO2=12.0+16.0+16.0=44.0
Gram Molecular mass = the same, except using grams for units, CO2 = 44.0 grams
Moles: = grams/molecular mass
Stoichiometry
Balancing equations: Conservation of mass. Matter cannot be created or destroyed.
The number of atoms on each side of the arrow must be the same.
Using the balanced equation: The coefficients tell the ratio of moles (or molecules) between substances in the reaction.
N2 + 3 H2 à 2 NH3
According to the equation, one mole of nitrogen reacts with three moles of hydrogen,
Also…two moles of nitrogen would react with six moles of oxygen, etc…
Four types of reactions
Synthesis (composition): putting two or more substances together
2 Mg + O2 à 2 MgO
Decomposition (analysis): breaking up one substance in to two or more substances
2 HgO à 2 Hg + O2
Single Replacement: An active metal (table J) replaces a less active metal in a compound Fe + CuCl2 à Cu + FeCl2
Double replacement (ion exchange): two ionic compounds trade partners (usually by forming a precipitate) AgNO3 (aq) + KOH (aq) à KNO3 (aq) + AgOH (s)
Unit IV: Bonding
Important tables = centerfold, table S
Valence Electrons: outer shell electrons
Full outer shell: eight valence electrons (a stable octet) makes an atom stable
Noble gases: elements with a full outer shell (note: helium has only 2 valence)
Bonding: what other elements do to become stable
Electron Dot Diagram: Dots represent valence electrons
Ionic Bonding
Ionic Bond: transfer of electrons from a metal to a nonmetal
(ionic bonds always have both a metal and a nonmetal)
metals become small (+) ions, nonmetals become large (-) ions
Ionization Energy: table S, energy needed to remove one electron from an element.
Metals generally have low ionization energies, nonmetals high, noble gases highest
Lewis Structure: for ionic compounds must show separate positive and negative ions
(brackets are often used)
Covalent Bonding
Covalent Bond: Sharing electrons, between two nonmetals, or hydrogen with a nometal
Single bond = 2 shared electrons, double bond = 4 shared, triple = 6 shared
Lewis Structure: for covalent bonds must show shared electrons, no brackets.
Electronegativity: (table S) Attraction for electrons in a covalent bond.
Polar Bond: One atom in the bond has a stronger electronegativity. Unequal sharing.
Subtract electronegativities to determine how polar (or ionic) a bond is
Nonpolar Bond: Atoms have nearly the same electronegativity. Equal Sharing.
Polar Molecule: (dipole) has polar bonds and is unsymmetrical. H2O, HCl, NH3
Nonpolar Molecule: symmetrical. O2, CO2, CH4
Attractions between particles in crystals
Ionic Crystals: Salts, Metal with a nonmetal
Salts have high melting points due to (+)(-) attraction of ions.
Salts only carry electricity when melted or mixed with water
(free floating charges are needed to carry electricity)
Molecular Crystals: covalent molecules
Lower melting points, never carry electricity (sugar, wax, etc)
Van Der Waals Attraction: weak attraction between nonpolar molecules
Dipole Dipole Attraction: attraction between polar molecules (dipoles)
Hydrogen Bonding: extra strong attraction between molecules of H2O, HF, NH3
Accounts for the unusually high boiling point of water
Metallic Crystals: A lattice of positive ions immersed in a sea of free floating electrons
Free floating electrons allow metals to conduct electricity as a solid or a liquid
Network Covalent Solids: Continuous covalent bonds throughout the crystals
(diamonds)
