MCAT Chemistry

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MCAT Chemistry
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Barron's MCAT flashcards
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  1. Periodic Table
  2. Atom
    • The basic unit of matter, building block for molecules
    • Has a core nucleus of positively charged protons and neutrons
    • The nucleus is surrounded by a cloud of negatively charged electrons
    • An atom of one element always has the same number of protons called the atomic number (Z)
  3. Subatomic Particles
    • Electrons: negatively charged particles orbiting the nucleus
    • Protons: positively charged particles within the nucleus
    • Neutrons: unchanged particles balancing protons in the nucleus, slightly larger than a proton

    Mass: neutron = proton >> electron
  4. Atomic Mass Unit, Mass Number, and Atomic Weight
    • Atomic mass unit: precisely measured as the equivalent of one-twelfth the weight of a carbon-12 atom, roughly equivalent to mass of a proton or neutron; units are Daltons (Da)
    • Mass number (A): the sum of the number of protons and neutrons in an atom
    • Atomic weight: weighted average of all naturally occurring isotopes
  5. Moles and Molar Weight
    • Mole: a standard (SI) unit to describe an elemental number of particles in 12 grams of Carbon-12 = Avagadro's number: 6.02x10^23
    • One mole always has 6.02x10^23 particles: One mole of carbon-12 has many atoms as one mole of gold
    • Molar weight: the mass of one mole of a particular element based on the atomic weight of the element. If an element has an atomic mass of 24.3 amu (magnesium), then one mole of magnesium weighs 24.3 grams. Based on carbon-12-- 1 mole of C-12 is 12g.
  6. Elements, Isotopes
    • An element is distinguished by its atomic number, which is the number of protons in the nucleus
    • Isotopes are atoms of the same element, with different numbers of neutrons in the nucleus. Remember, atoms of the same element always have the same atomic number = number of protons
  7. Ions
    • A neutral atom has an equal number of protons and electrons
    • Ion: an atom that has gained or lost electrons becomes charged
    • Cation: positive ion: loss of electrons
    • Anion: negative ion: gain of electrons
    • Calculate Charge: protons (+1) cancel out electrons (-1) in a neutral atom; charged atoms (ions) have different number of protons and electrons
  8. Quantum Mechanics and Quantum Numbers
    • Quantum mechanics states that electrons in an atom can only exist at specific energy levels, which are said to "quantized"
    • Quantum numbers: a set of four numbers describing that energy state of an electron
    • 1. Principal quantum number
    • 2. Azimuthal quantum number
    • 3. Magnetic quantum number
    • 4. Spin quantum number
  9. Principal (first) Quantum Number (n)
    • Called the principal number, describes the "shell" in which electrons of each shell (n) will exist
    • n=0,1,2,3,4; progressively increasing with increasing energy level
    • n corresponds to the row on the periods table; n=1 is row 1 (hydrogen and helium)
  10. Azimuthal (second) Quantum Number (l)
    • Defines the "sub-shell" in which an electron is found; describes shape of orbital
    • -l=0 called s orbital: spheres
    • -l=1 called p orbital: dumbbells
    • -l=2 called d orbital
    • l=3 called f orbital
    • -l ranges from 0 to n-1; if n=4; l=0,1,2,3
  11. Magnetic (third) Quantum Number (m)
    • Defines the unique energy state of an electron; number of orbitals in each shell (n)= number of m
    • m is found in the range of -l to l
    • l=0; m=0; s sub-shell has 1 orbital
    • l=1; m=-1,0,1; p sub-shell has 3 orbitals
    • l=2; m=-2,-1,0,1,2; d sub-shell has 5 orbitals
    • l=3; m=-3,-2,-1,0,1,2,3; f sub-shell has 7 orbitals
  12. Spin (fourth) Quantum Number (s)
    • Called the "spin number", this defines the unique magnetic spin direction of an electron
    • s=-1/2 (spin-down) or 1/2 (spin-up)
    • Spin number limits electrons in the same orbital from existing in exactly the same configuration
  13. Electron Configuration and Orbitals
    • Electron configuration: electrons fill energy levels within an atom in an ordered pattern; therefore, two atoms of the same element with the same charge always have the same electron configuration. Isoelectric atoms are two different elements with the same electron configuration
    • Orbitals: define a three-dimensional region in which the electron may be found orbiting the nucleus; defined specifically by quantum numbers
  14. Pauli Exclusion and Aufbau Principle
    • Pauli exclusion: no two electrons may exist in the same quantum state; therefore, may not have the same quantum numbers
    • Aufbau principle: the electron configuration of an atom is built up by adding electrons in an orderly pattern. Electrons are added progressively into orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s
  15. Electron Energy Levels, Absorption, and Emission Spectra
    • Electrons exist in a ground state that can be excites by a photon (single particle of light) to higher energy levels, but only by photons with specific amounts of energy will cause excitation
    • Absorption spectra: certain wavelengths of light (photons with a specific energy level) will cause electrons to move to higher levels. The wavelengths of light causing excitation are specific for each element
    • Emission spectra: when electrons fall down energy levels they emit specific wavelengths of light unique to each element
  16. Effective Nuclear Charge and Shielding
    • Effective nuclear charge: the net positive charge, and therefore attraction, that orbiting valence electrons experience from the nucleus of a particular element
    • Shielding: the negative charge of electrons in closer (lower energy) orbitals "shield" the electrons of higher orbitals from the positive charge of the nucleus. Greater shielding results in less attraction and therefore, these electrons are more easily removed from the atom
  17. Valence Electrons and Octet Rule
    • Electrons in the outermost shell of an atom are called valence electrons
    • Valence electrons form chemical bonds, and therefore determine the reactive properties of individual atoms
    • Atoms with the same number of valence electrons therefore have similar chemical properties
    • Octet rule: most atoms tend to bond in a way to obtain eight valence electrons; ions become more stable when they obtain a valence octet
  18. Periodic Trends: Atomic and Ionic Radius
    • Atomic radius: a measure of the size of an element
    • -Decreases across a period: increasing positive charge (protons)
    • -Increases down a group as new valence shells are added
    • Ionic radius: similar in concept to the atomic radius, this describes the size of ions of particular elements
    • -Increasing with the addition of electrons because of large orbitals without an increase in positive nuclear charge
  19. Periodic Trends: First and Second Ionization Energy
    • Ionization energy (IE): the energy necessary to remove a valence electron from an element
    • -Increases across a period because greater positive charge in the nucleus has stronger attraction on valence electrons
    • -Decreases down a column because of greater shielding of valence electrons from increasing number of orbitals
    • First IE is the energy to remove one electron, the second IE is the energy to remove another electron
    • -Second IE is greater when removing one electron causes an ion to reach full octet
  20. Periodic Trends: Electron Affinity and Electronegativity
    • Electron affinity (EA): the energy released in adding an electron to an atom
    • -Increases across a period because these elements want to gain electrons to complete their octet; toward the left elements want to lose electrons to reach the more stable octet
    • Electronegativity: the strength of an atom's pull on electrons of other atoms also increases across a period (similar to IE and EA) because the valence shells get closer to being an octet and more stable by taking on electrons
  21. Periodic Groups: Alkali Metals and Alkaline Earth Metals
    • Alkali metals: Group I monovalent elements (first column of periodic table: Li, Na, K,...), which are highly reactive with a very low ionization energy (IE) and therefore easily give up an electron to form a more stable completed octet shell single cations
    • Alkaline earth metals: Group II divalent elements (second column: Be, Mg, Ca,...), less reactive than alkali metals; however, also have low IE (especially second IE) and tend to give up two electrons to have an octet and double charged cations
  22. Periodic Groups: Halogens and Noble Gases
    • Halogens: Group VII elements (seventh column of periodic table: F, Cl, Br,...), which are highly reactive and some of the most electronegative elements because they need only one electron to complete their octet. Make strong acid when combined with hydrogen
    • Noble gases: Group VIII elements (last column of periodic table: He, Ne, Ar,...), which are colorless, odorless, and non reactive. These gases have a completed octet and rarely participate in chemical reactions. For the MCAT, consider noble gases completely inert in all systems
  23. Periodic Groups: Transition Metals
    • Transition metals: d-block elements, which have partially filled d orbitals and tend to be electropositive with variable oxidation states
    • These elements are found in the middle of the periodic table and often act as the central atoms in coordinate molecules
  24. Periodic Groups: Nonmetals and Metalloids
    • Nonmetals: group of elements that are poor conductors; dull and brittle, and have low melting/boiling points unlike metals: H, C, N, P, halogens, and noble gases
    • Metalloids: a group of elements that have properties of both metals and nonmetals:B, Si, Ge, As, Sb, Te, Po; sometimes called semiconductors-- look at the periodic table, these exist between transition metals and nonmetals
  25. Molecules and Molecular Formula
    • A molecule is a stable group of atoms held together by chemical bonds-- covalent or ionic
    • The molecular formula is established by counting the number of atoms for each element in the molecule
  26. Formula Weight and Molecular Weight
    • The formula weight of a molecule is the sum of the atomic mass times the number of each element component: C=12, H=1, O=16; 6(12) + 12(1) + 6(16)=180 amu (glucose)
    • The molecular weight is based on the formula weight, so one mole of glucose weighs 180g; which has 6.02x1023 glucose molecules
  27. Covalent Bonds and Molecular Orbits
    • Covalent bond: formed by the roughly equal sharing of electrons between two atoms, each contributing unpaired valence electrons. Double and triple bonds form when multiple valence electrons are shared between two atoms
    • Molecular orbital (MO): when atoms form bonds by sharing electrons the orbitals of the individual atoms overlap to create molecular orbitals, which similarly describe the electron distribution around the molecule
  28. Polar Covalent Bond and Formal Charge
    • Polar covalent bond: there is unequal sharing of electrons between two different elements in a covalent bond with more electron density near the more electronegative atom in a molecule
    • Formal charge (FC): exists when an atom within a molecule has greater or fewer valence electrons than it would normally have when not bonded; it can be calculated by FC= normal # valence e-- lone pair e-- 1/2 # of bonded valence e-
  29. Ionic Bonds
    • Ionic bonds form when one atom has a significantly greater electronegativity than its bonding partner. In this case, the more electronegative atom takes a valence electron, creating a cation and anion pair
    • The anion and cation are bonded together as a "salt" that may be dissolved in water, causing dissociation into separate anions and cations in solution
  30. Dipole and Dipole-Dipole Attraction/Repulsion
    • Dipole: created when there is a difference in electronegativity in a covalent bond so that electrons are pulled more toward one atom than the other, creating a negative and positive polarity (the dipole moment points toward the negative polarity)
    • Dipoles of two molecules can interact, with like dipoles repelling and unlike dipoles attracting to create a weak intermolecular bond
  31. Hydrogen Bonds
    • Hydrogen bonds occur between two molecules with a strongly electronegative element (must be O, N, F) bound to hydrogen
    • A dipole-dipole interaction, in which the hydrogen of one molecule (positive dipole) is attracted to the negative dipole at the electronegative atom of the second molecule
    • Hydrogen bonds are perhaps the most important bonds in chemistry and biology; they are the cause of surface tension, the reason ice floats, and coordination of the three-dimensional structure of DNA and proteins
  32. VSEPR Theory: Geometric Patterns
    • VSEPR: valence shell electron-pair repulsion theory states that electron pairs (bonded or nonbonded) repel each other and therefore move as far apart as possible in three dimensions while remaining bonded to a central atom
    • The geometric configuration of molecules therefore depends on the number of electron pairs -- with nonbounded electron pairs preferentially exerting greater repulsion than bonded electron pairs
  33. Magnetic Properties: Diamagnetic vs. Paramagnetic
    • Paramagnetic: atoms or molecules with unpaired valence electrons; when an external magnetic field is applied, a temporary magnetization is created and these atoms are attracted to the magnet
    • Diamagnetic: atoms or molecules with all paired electrons; they are repelled from a magnet
  34. Lewis Dot Structures
    • Lewis dot structures diagram the valence electrons of an atom or molecule with a dot representing each electron around the atom and lines to represent bonded electron pairs between molecules
    • Individual atoms: 
    • Molecules: 
    • In molecules, single bonds are represented by single line (-), double bonds by double lines (=), and triple bonds by (☰)
  35. Resonance
    • Multiple Lewis structures may accurately represent the electron structure for a single molecule, and therefore electrons may be distributed into several patterns
    • Resonance diagrams demonstrate these multiple patterns of electron localization and the resonance structure shows the "average" electron distribution in a molecule using dotted lines to show partial localization of electrons
  36. Oxidation State
    • Oxidation state: an indicator of ownership of electrons in a molecule; alternatively it is an indicator of the charge of each atom in a molecule if the bonds were considered to be completely ionic
    • The oxidation state is a relative charge; for ionic molecules, like NaCl, the oxidation state of each component is clear- Na+ and Cl-
    • However, oxidation state for covalent molecules is not as clear; though more electronegative atoms tend to have negative oxidation while those with lower electron affinity have positive oxidation state according to a set of rules
  37. Oxidation Rules
    • The rules for assigning oxidation state are applied in the following order of most-to-least importance:
    • 1. The sum of oxidation states is always neutral for neutral atoms or equal to the formal charge of charged molecules
    • 2. Group 1 metals are +1 and group 2 metals are +2
    • 3. Halogens (group VII) have a -1 oxidation state
    • 4. Hydrogen has a +1 oxidation state
    • 5. Oxygen has a -2 oxidation state, along with other group VI elements
  38. Chemical Reactions: Combination and Decomposition
    • Chemical reaction: a process that changes the chemical nature of one or more reactants to one or more products
    • Combination reaction: A + B -> C
    • Decomposition reaction: Z -> X + Y

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