GOOD LUCK EVERYONE!!!!
UNIT 1: Atomic Structure and Properties
1.1 Moles and Molar Mass
Avogadro’s number = 6.02 * 10^23 (1 mol of everything has this many)
Atomic mass(amu) = molar mass(g/mol)
One unit → another unit; use conversion factors
1.2 Mass Spectra of Elements
Mass spectrum = graph of isotopes + abundance
o Each bar represents an isotope
To calculate atomic mass from spectrum: ∑ (isotope * relative abundance in decimal)
Mass number = protons + neutrons
1.3 Elemental Composition of Pure Substances
To find percent composition: 1) find mass of each element 2) find mass of compound 3)
(mass of each element / mass of compound) *100
Empirical formula = lowest whole number ratio of atoms vs Chemical/Molecular formula
= actual number of atoms
o Same empirical formula = same % composition
To find empirical formula from experimental data: 1) convert all units to moles 2) divide
everything by lowest number of moles 3) if needed multiply to make whole numbers
o If given in percentages, use the percentage as the mass in grams
For hydrate analysis problems: moles of water removed / moles of anhydrous substance
1.4 Composition of Mixtures
Elemental analysis: finding mass of elements in a mixture to see how pure the mixture is
1.5 Atomic Structure and Electron Configuration
Coulomb's law: F = (q1 * q2) / r^2
o f → force/attraction; q → charges; r → distance
1.6 Photoelectron Spectroscopy
Aufbau principle: electrons fill lower energy levels first
Photoelectron spectroscopy: each peak represents different sublevel (ex: 1s^2, 2s^2),
height of peaks represent # of electrons
1.7 Periodic Trends
Atomic radius, ionization energy, electronegativity, electron affinity
1.8 Valence Electrons and Ionic Compounds
UNIT 2: Compound Structure and Properties
2.1 Types of Chemical Bonds
Ionization energy is nucleus of one atom and electron of the same atom; chemical bond is
the nucleus of one atom and electron of another
Ionic (metal + nonmetal), covalent (nonmetals), metallic (metals, delocalized electrons)
2.2 Intramolecular Force and Potential Energy
Energy curves show stable arrangement for atoms; most stable finds a balance between
repulsive and attractive forces
o Positive values = unstable; negative values = stable
Lattice energy: energy required to separate ions in an ionic bond
, 2.3 Structure of Ionic Solids
To conduct electricity a substance must have charged particles and be able to move freely
2.4 Structure of Metals and Alloys
Properties of metals: conduct electricity, malleable, ductile
Alloys: combines 2+ metals
o Substitutional: atoms with similar radius; substitute in alloy
o Interstitial: atoms with different radius; smaller atom fill sup space between larger
ones; stronger than metal
2.5 Lewis Diagrams
Only valence electrons
Bond energy increases as number of bonds increase
Central atoms are the least electronegative
Elements in period 3 and below can have expanded octet
To draw lewis structure for molecules
o Count electrons
o Draw at least one bond between each element first
o Fill in remaining electrons
o If octet not met, draw double, then triple bonds
2.6 Resonance and Formal Charge
Length of resonance bonds are equal; either 1.5 or 1 ⅓ instead of double bond or single
bond
Resonance shown by double-sided arrow
Formal charge = num of valence electrons - num of assigned electrons
o Dominant structure = least number of nonzero charges; negative charges assigned
to electronegative elements
2.7 VSEPR and Hybridization
Valence Shell Electron Pair Repulsion
o Bonds and lone pairs will arrange themselves as far apart as possible
Geometry depends on how many electron domains (bonds / lone pairs) around central
atom
o Steric number
o Lone pairs are more repulsive
VSEPR CHART AND BOND ANGLES
Bond polarity
o In a molecule, if the bonds are polar, but equal in magnitude and opposite in
direction, the molecule is NOT polar
o Symmetrical shapes (cannot be polar): linear, trigonal planar, tetrahedral, trigonal
bipyramidal, square planar, octahedral
Atomic orbitals (s,p,d,f) are only in unbonded atoms
o When atoms bond, they form hybrid orbitals
o 4 electron domains = sp^3; 3 ED = sp^2; 2 ED = sp
o single bond = 1 sigma; double bond = 1 pi 1 sigma; triple bond = 1 sigma 2 pi
UNIT 3: Properties of Substances and Mixtures
3.1 Intermolecular and Interparticle Forces
Intra- in the molecule (covalent bonds); inter- attraction between two different molecules
UNIT 1: Atomic Structure and Properties
1.1 Moles and Molar Mass
Avogadro’s number = 6.02 * 10^23 (1 mol of everything has this many)
Atomic mass(amu) = molar mass(g/mol)
One unit → another unit; use conversion factors
1.2 Mass Spectra of Elements
Mass spectrum = graph of isotopes + abundance
o Each bar represents an isotope
To calculate atomic mass from spectrum: ∑ (isotope * relative abundance in decimal)
Mass number = protons + neutrons
1.3 Elemental Composition of Pure Substances
To find percent composition: 1) find mass of each element 2) find mass of compound 3)
(mass of each element / mass of compound) *100
Empirical formula = lowest whole number ratio of atoms vs Chemical/Molecular formula
= actual number of atoms
o Same empirical formula = same % composition
To find empirical formula from experimental data: 1) convert all units to moles 2) divide
everything by lowest number of moles 3) if needed multiply to make whole numbers
o If given in percentages, use the percentage as the mass in grams
For hydrate analysis problems: moles of water removed / moles of anhydrous substance
1.4 Composition of Mixtures
Elemental analysis: finding mass of elements in a mixture to see how pure the mixture is
1.5 Atomic Structure and Electron Configuration
Coulomb's law: F = (q1 * q2) / r^2
o f → force/attraction; q → charges; r → distance
1.6 Photoelectron Spectroscopy
Aufbau principle: electrons fill lower energy levels first
Photoelectron spectroscopy: each peak represents different sublevel (ex: 1s^2, 2s^2),
height of peaks represent # of electrons
1.7 Periodic Trends
Atomic radius, ionization energy, electronegativity, electron affinity
1.8 Valence Electrons and Ionic Compounds
UNIT 2: Compound Structure and Properties
2.1 Types of Chemical Bonds
Ionization energy is nucleus of one atom and electron of the same atom; chemical bond is
the nucleus of one atom and electron of another
Ionic (metal + nonmetal), covalent (nonmetals), metallic (metals, delocalized electrons)
2.2 Intramolecular Force and Potential Energy
Energy curves show stable arrangement for atoms; most stable finds a balance between
repulsive and attractive forces
o Positive values = unstable; negative values = stable
Lattice energy: energy required to separate ions in an ionic bond
, 2.3 Structure of Ionic Solids
To conduct electricity a substance must have charged particles and be able to move freely
2.4 Structure of Metals and Alloys
Properties of metals: conduct electricity, malleable, ductile
Alloys: combines 2+ metals
o Substitutional: atoms with similar radius; substitute in alloy
o Interstitial: atoms with different radius; smaller atom fill sup space between larger
ones; stronger than metal
2.5 Lewis Diagrams
Only valence electrons
Bond energy increases as number of bonds increase
Central atoms are the least electronegative
Elements in period 3 and below can have expanded octet
To draw lewis structure for molecules
o Count electrons
o Draw at least one bond between each element first
o Fill in remaining electrons
o If octet not met, draw double, then triple bonds
2.6 Resonance and Formal Charge
Length of resonance bonds are equal; either 1.5 or 1 ⅓ instead of double bond or single
bond
Resonance shown by double-sided arrow
Formal charge = num of valence electrons - num of assigned electrons
o Dominant structure = least number of nonzero charges; negative charges assigned
to electronegative elements
2.7 VSEPR and Hybridization
Valence Shell Electron Pair Repulsion
o Bonds and lone pairs will arrange themselves as far apart as possible
Geometry depends on how many electron domains (bonds / lone pairs) around central
atom
o Steric number
o Lone pairs are more repulsive
VSEPR CHART AND BOND ANGLES
Bond polarity
o In a molecule, if the bonds are polar, but equal in magnitude and opposite in
direction, the molecule is NOT polar
o Symmetrical shapes (cannot be polar): linear, trigonal planar, tetrahedral, trigonal
bipyramidal, square planar, octahedral
Atomic orbitals (s,p,d,f) are only in unbonded atoms
o When atoms bond, they form hybrid orbitals
o 4 electron domains = sp^3; 3 ED = sp^2; 2 ED = sp
o single bond = 1 sigma; double bond = 1 pi 1 sigma; triple bond = 1 sigma 2 pi
UNIT 3: Properties of Substances and Mixtures
3.1 Intermolecular and Interparticle Forces
Intra- in the molecule (covalent bonds); inter- attraction between two different molecules
