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MCAT Review > Kaplan — General Chemistry > Flashcards

Kaplan — General Chemistry Flashcards

(92 cards)

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1
Q

Molecules

A

Combination of elements via covalent bonding

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2
Q

Chemical bonds

A

Strong attractive forces formed via the interaction of valence electrons of the combining atoms

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3
Q

Octet rule

A

States that an atom tends to bond with other atoms so that is has 8 electrons in its outermost shell

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4
Q

Exceptions to octet rule (3)

A
  1. Incomplete octet — stable with fewer than 8 electrons → hydrogen, helium, lithium, beryllium, boron
  2. Expanded octet — element in period 3 or greater → phosphorous, sulfur, chlorine
  3. Odd numbers of electrons — molecules that cannot distribute to give 8 to each atom → nitric oxide
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5
Q

Ionic bonding

A

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

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6
Q

Covalent bonding

A

Electron pair is shared between two atoms

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7
Q

Polar covalent bonding

A

Sharing of the electron pair is unequal

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8
Q

Non-polar covalent bonding

A

Sharing of the electron pair is equal

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9
Q

Coordinate covalent

A

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

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10
Q

Cation

A

Positively charged atom

Atom that loses the electron

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11
Q

Anion

A

Negatively charged atom

Atom that gains the electrons

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12
Q

Crystalline lattice

A

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

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13
Q

Bond order

A

Number of shared electron pairs shared between two atoms

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14
Q

Bond length

A

Average distance between the two nuclei of atoms in a bond

Single > double > triple

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15
Q

Bond energy

A

Energy required to break a bond by separating its components into their isolated, gaseous atomic states

Triple > double > single

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16
Q

Polarity

A

When two atoms have a relative difference in electronegativity

Atom with high electronegativity have a larger share of electron density

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17
Q

Dipole moment

A

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

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18
Q

Units of dipole moment

A

Debye units (coulomb-meters)

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19
Q

Bonding electrons

A

Electrons involved in a covalent bond & located in valence shell

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20
Q

Non-bonding electrons

A

Electrons located in valence shell & not involved in covalent bonds

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21
Q

Lewis structure

A

System of notation developed to keep track of bonded and non-bonded electron pairs

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22
Q

Formal charge

A

Formal charge = number of electrons in atom’s valence shell — number of non-bonding electrons — half of the number of bonding electrons

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23
Q

How to make a Lewis structure

A

(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

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24
Q

Resonance structures (Lewis)

A

All possible resonance structures connected by a double-headed arrow

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25
Resonance hybrid
Actual structure of compound formed by relative combination of all resonance structures
26
Stability of different resonance structures
- Small or no formal charges - Less separation between opposite charges - Negative charges on electronegative atoms
27
Number of electrons for hydrogen
Stable with 2
28
Number of electrons for helium
Stable with 2
29
Number of electrons for lithium
Stable with 2
30
Number of electrons for beryllium
Stable with 4
31
Number of electrons for boron
Stable with 6
32
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
33
Regions of electron density
Lone pairs and bonds
34
2 regions of electron density
Linear
35
3 regions of electron density
Trigonal planar
36
4 regions of electron density
Tetrahedral
37
5 regions of electron density
Trigonal bipyramidal
38
6 regions of electron density
Octahedral
39
Electron geometry
Spatial arrangement of all pairs of electrons around the central atom, including both the bonding and the lone pairs
40
Molecular geometry
Spatial arrangement of only the bonding pairs of electrons
41
Coordination number
Number of atoms that surround and are bonded to a central atom
42
Ideal bond angle
Determined by the VSEPR model but can be altered by repulsion from non bonding pairs
43
Molecular orbital
Overlap between two atomic orbitals
44
Bonding orbital
Signs of 2 atomic orbitals are the same
45
Anti-bonding orbital
Signs of 2 atomic orbitals are different
46
Sigma bond
Head to head overlap of two orbitals that allows for free rotation
47
Pi bond
Caused by parallel electron cloud densities Do not allow free rotation
48
London dispersion forces
Shifting polarities that cause transient bonds
49
Dipole-dipole interactions
Occur between the oppositely charged ends of polar molecules
50
Hydrogen bonds
Specialized subset of dipole-dipole interactions between a hydrogen bonded to NOF and a NOF atom
51
Compound
Pure substance composed of two or more electrons in a fixed proportion
52
Molecule
Combination of two or more atoms held together by covalent bonds Can be the same elements or different elements
53
Formula unit
Empirical formula of the compound Subunit of an ionic compound
54
Formula weight
Weight of formula unit
55
Molecular weight
Sum of atomic weights of all the atoms in a molecule and is in units of atomic mass units (amu) per molecule
56
Mole
Quantity of any substance equal to the number of particles found in 12 grams of carbon-12
57
Avogadro’s number (N_A)
6.022 x 10^23 mol^(-1)
58
Molar mass
Mass of one mole of a compound Expressed in g/mol
59
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
60
Gram equivalent weight
Equals molar mass / n, where n is number of particles of interest
61
Number of equivalents
Mass of compound / gram equivalent weight
62
Normality
Molarity times n
63
Law of constant composition
Any pure sample of a given compound will contain the same elements in identical mass ratio
64
Empirical formula
Simplest whole-number ratio of elements
65
Molecular formula
Exact number of atoms of each element in the compound and is a multiple of the empirical formula
66
Percent composition
Mass of element in the formula / molar mass in percentage
67
Combination reaction
Two or more reactions forming one product
68
Decomposition reaction
Single reactant breaks down into two or more productions
69
Combustion reaction
Hydrocarbon + oxidant → carbon dioxide + water
70
Single-displacement reaction
When an atom or ion in a compound is replaced by an atom or ion of another element
71
Double-displacement reaction
Two different compounds swap places with each other to form two new compounds
72
Neutralization reaction
Acid + base → salt + water
73
Stoichiometric coefficients
Indicate the relative number of moles of a given species involved in the reaction
74
Limiting reactant
Reactant that limits the amount of product that can be formed in the reaction
75
Excess reactants
Reactants that remain after all the limiting reagent is used
76
Theoretical yield
Maximum amount of product that can be generated as predicted from the balanced equation
77
Actual yield
Amount of product one actually obtains during the reaction
78
Percent yield
Actual yield / theoretical yield in percentage form
79
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
80
Mono-atomic anion names
Dropping the ending of the name of the element and adding -ide
81
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-
82
Oxyanion names with hydrogen
Add hydrogen or dihydrogen in front of the anion’s name Can add bi- instead of hydrogen
83
Ammonium
NH4^+
84
Acetate
C2H3O2^-
85
Cyanide
CN^-
86
Permanganate
MnO4^-
87
Thiocyanate
SCN^-
88
Chromate
CrO4^2-
89
Dichromate
Cr2O7^2-
90
Borate
BO3^3-
91
Oxidation states
Different charges that elements can adopt
92
Electrolytes
Solutes that enable solutions to carry currents

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