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MCAT > Chemistry Fundamentals and Periodic Trends > Flashcards

Chemistry Fundamentals and Periodic Trends Flashcards

1
Q

Metric Units to measure:

Length
Mass
Time
Amount
Temperature
Electrical Current
A 
Meter (m) 
Kilogram (kg) 
Second (s) 
Mole (mol) 
Kelvin (K)
Ampere (A)
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2
Q

Order of Magnitude

A

Factor of 10

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3
Q

Density

A

Mass/Volume (m/v)

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4
Q

Molecular Formulas

A

Give identities and number of atoms in a molecule

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5
Q

Empirical Formulas

A

Smallest whole numbers that give the same ratio of atoms in a molecule

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6
Q

Formula Weight

A

Sum of the atomic weights of all atoms in an IONIC compound

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7
Q

Molecular Weight

A

Term for molecules

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8
Q

Atomic Mass Unit (AMU)

A

Unit for atomic weight. One AMU is 1/12th mass of an atom of Carbon-12

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9
Q

Mole

A

A particular number of things, collection of 6.022x10^23 of something (Avogadro’s Number)

moles = mass (g)/molecular weight

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10
Q

Molar Mass

A

Same as molecular weight but in grams/mol

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11
Q

Percentage Composition by Mass

A

Use empirical formula to find molecule’s percent mass composition

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12
Q

Molarity (M)

A

Measure of concentration that calculates the moles of solute in liters of a solution

Molarity = # moles/#liters

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13
Q

Mole Fraction

A

Fraction of moles of a given substance

Mole fraction = Xs = # moles of substance S/total # moles in a solution

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14
Q

Law of Conservation of Mass (Matter)

A

Amount of matter, and therefore mass, does not change in a chemical reaction.

Modify stoichiometric coefficients to balance Eq and have equal mass on both sides of arrow

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15
Q

Limiting Reagent

A

Limits the extent of a reaction (which reactant you are going to run out of first given an eq)

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16
Q

Notation of States of Matter

s
l
g
aq

A

Phases of substances of reaction

solid
liquid
gas
aqueous

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17
Q

Oxidation States

A

Atom’s ownership of its valence electrons change when it forms a compound

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18
Q

Rules for assigning Oxidation States (7)

A
  1. Oxidation of any element in standard state is 0
  2. Sum of oxidation states of atoms must equal molecule/ion’s overall charge
  3. Group 1 metals have +1 state, group 2 metals have +2
  4. F has -1 oxidation state
  5. H has a +1 state when with an element more electronegative than carbon, -1 with less electronegative, 0 with carbon
  6. Oxygen has -2 oxidation state (exception of peroxides with -1)
  7. Rest of halogens have -1 state, oxygen family has -2 oxidation state
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19
Q

Order of Electronegativity

A

F>O>N>Cl>Br>I>S>C>H

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20
Q

Atoms

A

Smallest unit of any element contain a central nucleus with nucleons (protons +1/neutrons 0) and surrounding electron cloud -1.

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21
Q

Number Neutrons

A

= Mass Number - Atomic Number (Number Protons)

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22
Q

Isotopes

A

Varying number of neutrons between two atoms

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23
Q

Charged Ions

A

Anion (negatively charged)

Cation (positively charged)

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24
Q

Atomic Weight

A

Weighted average of masses of naturally occurring isotopes

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25
Radioactive decay - 3 types
Radioactive unstable nuclear undergo transformation to make them more stable Alpha/beta/gamma
26
Alpha Decay
Large nucleus wants to be more stable by reducing number of protons and neutrons - emits alpha particle (2 protons and 2 neutrons) 4 a 2
27
Beta Decay
beta- decay = unstable nucleus that contains too many neutrons, may convert neutron into a proton and electron which is ejected 0 B -1
28
Positron Emission
B+ decay = unstable nucleus containing too few neutrons, converts proton into a neutron and positron which is ejected 0 B +1
29
Electron Capture
Capture electron from closest electron shell and use it in the conversion of a proton to a neutron 0 e- -1
30
Gamma Decay
Nucleus in excited energy state (after any sort of alpha or beta decay) can relax to its ground state by emitting energy in the form of one or more photons of electromagnetic radiation (called gamma photons) No change in atomic number/mass number, the gamma ray is emitted.
31
Half-life
Time it takes for one-half of some sample of the substance to decay. Shorter the half-life, the faster the decay. Decay is exponential over time
32
Decay Constant
Inversely proportional to half life
33
Nuclear Binding Energy
Energy released when individual nucleons (Protons/neutrons) bound together to form the nucleus. Also equal to energy required to break up the intact nucleus. Greater binding energy equals more stable
34
Mass Defect
When nucleons bind together, some mass converted to energy, so the mass defect is the difference between total mass of separate nucleons and the mass of hte nucleus
35
Einstein's Equations for Mass- Energy Equivalence
E = mc^2
36
Emission Spectrum
Provides identity for a particular element through separation into component wavelengths Gives energetic 'fingerprint'
37
Energy
Ephoton = hf = h(c/wavelength) ``` h = Planck's constant f = frequency c = speed of light ```
38
Planck's Constant
h = 6.63 x 10^-34
39
Bohr Model
Electron orbit nucleus in a circular path. Distance between n and n+1 decreases with increasing n
40
Photons
Absorbed to make an electron jump from ground to excited state (positive energy change) Emitted to create a negative energy change (lower energy level, electron emits photon)
41
Quantum Model
Accounts for electron-electron interactions that exist in many-electron atoms
42
Energy Shell
(n) of an electron is analogous to circular orbits of Bohr model. higher the shell the higher the energy
43
Energy Subshell
Changes from circular orbits to 3D region around nucleus wherein electron likely to be found. spdf (shape and energy described)
44
Orbital Orientation
Each subshell contains one or more orbitals of same energy (degenerate orbitals)
45
Electron Spin
Every electron has two possible spin states (up and down), every orbital can only acount for 2 electrons one up and one down.
46
Afbau Principle
Electrons occupy lowest energy orbitals available
47
Hund's Rule
Electrons in same subshell occupy available orbitals singly before pairing up
48
Pauli Exclusion Principle
There can be no more than two electrons in any given orbital
49
Diamagnetic
All electrons spin paired
50
Paramagnetic
Not all electrons spin paired
51
Electron Configurations
Follow the periodic table (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc/)
52
Anomalous Configurations
Some atoms achieve lower energy state by having fully filled or half filled d-shell (5 or 10)
53
Isoelectronic
Same electron configurations (not necessarily same number of electrons)
54
Shielding Effect
Also known as nuclear shielding effect - each filled shell between nucleus and valence electrons shields electrons from positive charges in the nucleus Decrease of effective nuclear charge
55
Atomic /Ionic Radius
As you go across a period or up a group, atomic radius increases Cations smaller than neutral atom, anions larger than neutral
56
Ionization Energy
Removal of an electron that requires energy input Across period/up group, increases Second ionization energy always greater than the first
57
Electron Affinity
Addition of an electron (releases energy) Across a period and up a group, increases
58
Electronegativity
Atoms ability to pull electron density towards itself Increase across period up a group
59
Acidity
Willingness to donate a proton High electronegativity and large radius make favorable
60
With electronegativity trend, each of the following Atomic Radius Ionization Energy Electron Affinity Acidity
Decreases Increases Becomes more negative Increases across, but decreases up

MCAT (7 decks)

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