Kaplan — General Chemistry Flashcards
Molecules
Combination of elements via covalent bonding
Chemical bonds
Strong attractive forces formed via the interaction of valence electrons of the combining atoms
Octet rule
States that an atom tends to bond with other atoms so that is has 8 electrons in its outermost shell
Exceptions to octet rule (3)
- Incomplete octet — stable with fewer than 8 electrons → hydrogen, helium, lithium, beryllium, boron
- Expanded octet — element in period 3 or greater → phosphorous, sulfur, chlorine
- Odd numbers of electrons — molecules that cannot distribute to give 8 to each atom → nitric oxide
Ionic bonding
One or more electrons from an atom with a low ionization energy (e.g. metal) are transferred to an atom with a high electron affinity (e.g. non-metal) & the resulting electrostatic attraction between opposite charges is what holds the ions together
Covalent bonding
Electron pair is shared between two atoms
Polar covalent bonding
Sharing of the electron pair is unequal
Non-polar covalent bonding
Sharing of the electron pair is equal
Coordinate covalent
Both of the shared electrons are contributed by only one of the two atoms
When a lone pair of one atom attacked another atom with an unhybridized p-orbital to form a bond
Cation
Positively charged atom
Atom that loses the electron
Anion
Negatively charged atom
Atom that gains the electrons
Crystalline lattice
Compound form of ionic constituents consisting of repeating positive and negative ions
Attractive forces between oppositely charged are maximized, repulsive forces between ions of like charge are minimized
Bond order
Number of shared electron pairs shared between two atoms
Bond length
Average distance between the two nuclei of atoms in a bond
Single > double > triple
Bond energy
Energy required to break a bond by separating its components into their isolated, gaseous atomic states
Triple > double > single
Polarity
When two atoms have a relative difference in electronegativity
Atom with high electronegativity have a larger share of electron density
Dipole moment
Vector quantity given by equation p = qd, where p is dipole moment, q is magnitude fo charge, d is displacement vector separating the two charges
Units of dipole moment
Debye units (coulomb-meters)
Bonding electrons
Electrons involved in a covalent bond & located in valence shell
Non-bonding electrons
Electrons located in valence shell & not involved in covalent bonds
Lewis structure
System of notation developed to keep track of bonded and non-bonded electron pairs
Formal charge
Formal charge = number of electrons in atom’s valence shell — number of non-bonding electrons — half of the number of bonding electrons
How to make a Lewis structure
(1) Draw backbone of the compound
(2) Count all the valence electrons of the atoms
(3) Draw single bonds between central atoms and atoms surrounding it
(4) Complete the octets of all surrounding atoms
(5) Complete the octet of the central atom next, adjusting the single bonds to double or triple bonds
Resonance structures (Lewis)
All possible resonance structures connected by a double-headed arrow
Resonance hybrid
Actual structure of compound formed by relative combination of all resonance structures
Stability of different resonance structures
- Small or no formal charges
- Less separation between opposite charges
- Negative charges on electronegative atoms
Number of electrons for hydrogen
Stable with 2
Number of electrons for helium
Stable with 2
Number of electrons for lithium
Stable with 2
Number of electrons for beryllium
Stable with 4
Number of electrons for boron
Stable with 6
Valence shell electron pair repulsion (VSEPR)
Predicts the molecular geometry of covalently bonded molecules
States that the 3-D arrangement of atoms surrounding a central atom is determined by the repulsions between bonding and non-bonding electrons in the valence shell of the central atom
Regions of electron density
Lone pairs and bonds
2 regions of electron density
Linear
3 regions of electron density
Trigonal planar
4 regions of electron density
Tetrahedral
5 regions of electron density
Trigonal bipyramidal
6 regions of electron density
Octahedral
Electron geometry
Spatial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs
Molecular geometry
Spatial arrangement of only the bonding pairs of electrons
Coordination number
Number of atoms that surround and are bonded to a central atom
Ideal bond angle
Determined by the VSEPR model but can be altered by repulsion from non bonding pairs
Molecular orbital
Overlap between two atomic orbitals
Bonding orbital
Signs of 2 atomic orbitals are the same
Anti-bonding orbital
Signs of 2 atomic orbitals are different
Sigma bond
Head to head overlap of two orbitals that allows for free rotation
Pi bond
Caused by parallel electron cloud densities
Do not allow free rotation
London dispersion forces
Shifting polarities that cause transient bonds
Dipole-dipole interactions
Occur between the oppositely charged ends of polar molecules
Hydrogen bonds
Specialized subset of dipole-dipole interactions between a hydrogen bonded to NOF and a NOF atom
Compound
Pure substance composed of two or more electrons in a fixed proportion
Molecule
Combination of two or more atoms held together by covalent bonds
Can be the same elements or different elements
Formula unit
Empirical formula of the compound
Subunit of an ionic compound
Formula weight
Weight of formula unit
Molecular weight
Sum of atomic weights of all the atoms in a molecule and is in units of atomic mass units (amu) per molecule
Mole
Quantity of any substance equal to the number of particles found in 12 grams of carbon-12
Avogadro’s number (N_A)
6.022 x 10^23 mol^(-1)
Molar mass
Mass of one mole of a compound
Expressed in g/mol
Equivalent weight
How many moles of the thing we are interested in (protons, hydroxide ions, electrons, ions) will one mole of a given compound produce?
Example: when talking about hydrogen ions, 1 N HCl is 1 M HCl and 1 N H2CO3 is 0.5 M H2CO3
Gram equivalent weight
Equals molar mass / n, where n is number of particles of interest
Number of equivalents
Mass of compound / gram equivalent weight
Normality
Molarity times n
Law of constant composition
Any pure sample of a given compound will contain the same elements in identical mass ratio
Empirical formula
Simplest whole-number ratio of elements
Molecular formula
Exact number of atoms of each element in the compound and is a multiple of the empirical formula
Percent composition
Mass of element in the formula / molar mass in percentage
Combination reaction
Two or more reactions forming one product
Decomposition reaction
Single reactant breaks down into two or more productions
Combustion reaction
Hydrocarbon + oxidant → carbon dioxide + water
Single-displacement reaction
When an atom or ion in a compound is replaced by an atom or ion of another element
Double-displacement reaction
Two different compounds swap places with each other to form two new compounds
Neutralization reaction
Acid + base → salt + water
Stoichiometric coefficients
Indicate the relative number of moles of a given species involved in the reaction
Limiting reactant
Reactant that limits the amount of product that can be formed in the reaction
Excess reactants
Reactants that remain after all the limiting reagent is used
Theoretical yield
Maximum amount of product that can be generated as predicted from the balanced equation
Actual yield
Amount of product one actually obtains during the reaction
Percent yield
Actual yield / theoretical yield in percentage form
Different ions’ charges in nomenclature
Denoted by Roman numerals following the ion’s name or change suffix to -ous or -ic for less and greater charge respectively
Example: iron(II) vs. iron(III) & ferrous vs. ferric
Mono-atomic anion names
Dropping the ending of the name of the element and adding -ide
Oxyanion names
One with less oxygen ends in -ite
One with more oxygen ends in -ate
If there are more than 2 oxyanions in a series, the one with less is called hypo- and the one with more is per-
Oxyanion names with hydrogen
Add hydrogen or dihydrogen in front of the anion’s name
Can add bi- instead of hydrogen
Ammonium
NH4^+
Acetate
C2H3O2^-
Cyanide
CN^-
Permanganate
MnO4^-
Thiocyanate
SCN^-
Chromate
CrO4^2-
Dichromate
Cr2O7^2-
Borate
BO3^3-
Oxidation states
Different charges that elements can adopt
Electrolytes
Solutes that enable solutions to carry currents