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A property such as density that is independent of the amount of the given substance |
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| A property that depends on the amount of given substance, such as mass |
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| atoms of the same element with the same number of protons but different numbers of neutrons and consequently different masses |
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| (A) the mass of the element |
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The atomic number of the element (Z) that determines the number of protons and electrons in an atom |
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Natural abundance (percent abundance) |
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| the relative percentage of a particular isotope in a naturally occuring sample with respect to other isotopes of the same element |
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| The electrode is an electrochemical cell where oxidation occurs electrons flow away from the anode |
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| The electrode is an electrochemical cell where reduction occurs, electrons flow toward the cathode |
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| Highly reactive metals in Group 1A of the periodic table |
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| Fairly reactive metals in group 2A of the periodic table |
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| Law of Conservation of Mass |
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| A law stating that mass is neither created nor destroyed in a chemical reaction |
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Law of Definite Proportion |
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| A law stating that all samples of a given compound have the same proportions as their constituent elements |
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Law of Multiple Proportions |
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| A law stating that when two elements (A and B) form two different compounds, the masses of element B that combines with one gram of element A can be expressed as a ratio of small whole numbers |
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| A polyatomic atom containing a nonmetal covalently bonded to one or more oxygen atoms |
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| The smallest electrically neutrol collection of ions in an ionic compound |
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| A chemical formula that shows the simplest whole number ratio of atoms in a compound |
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The chemical formula that shows the actual number of atoms of each element in a molecule of a compound |
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| A positive or negative whole number that represents the "charge" an atom in a compound would have if all shared electrons were assigned to the atom with a greater attraction for those electrons |
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| The reactant that has the smallest stoichiometric amount in a reactant mixture and consequently limits the amount of product in a chemical reaction |
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| A substance that causes the reduction of another substance ; a reducing agent loses electrons and is oxidized |
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| The reactant where there is some left over at the end of the reaction |
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| The pressure due to any individual component in a gas mixture |
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| The proportionality constant of the ideal gas law (r = .0821) |
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| The law that combines the relationship of Boyle's, Charles, and Avogadro's laws into one comprehensive equation of state with the proportionality constant of R in the form PV = nRT. |
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| A model of an ideal gas as a collection of point particles in constant motion undergoing completely elastic collisions |
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electromagnetic radiation |
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| A form of energy embodied in oscillating electric and magnetic field |
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| The range of wavelengths emitted by a particular element used to identify the element |
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| The smallest possible packet of electromagnetic radiation with an energy equal to hv |
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| the observation that many metals emit electrons when light falls upon them |
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| The observation that the wavelength of a particle is inversely proportional to it's momentum |
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Heinsburg Uncertainty Principle |
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| The principle stating that due to the wave-particle duality, it is fundamentally impossible to precisely determine both the position and veleocity of a particle at any given time |
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| principal quantum number (n) |
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Definition
| An integer that expresses the overall size and energy of an orbital. The higher the quantum number n, the greater the average distance between the electron and the nucleus and the higher in energy |
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| angular momentum quantum number (l) |
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| An integer that determines the shape of an orbital |
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| magnetic quantum number (m) |
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| An integer that specefies the orientation of an orbital |
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| principle shells (levels) |
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| The group of orbitals with the same value of n |
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| Those orbitals in the same principle level with the same value of n and l |
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| The purpose of line spectra |
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| Spectra are used to identify elements |
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| Bohr's explanation of spectral lines: |
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| Light has energy equal to the difference in energy of the electron in two orbitals |
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| both wave and particle properties |
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| The fourth quantum number, which denotes the electrons spin order as 1/2 (up arrow) or -1/2 (down arrow) |
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| Pauli Exclusion Principle |
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| The principle that no two electrons in an atom can have the same four quantum numbers |
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| The principle stating that when electrons fill degenerate orbitals, they first fill them singly with parallel spins |
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| The principle that indicates the pattern of orbital filling in an atom |
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| A notation that shows the particular orbitals that are occupied by electrons in an atom |
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Effective Nuclear Charge (Zeff) |
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| The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons |
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| The way of expressing radius for metals |
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| covalent radius (bonding atomic radius) |
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| Defined in nonmetals as one-half the distance between two atoms bonded together, and in metals as one-half the distance between two adjacent atoms in a crystal of the metal |
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| those electrons in a complete principle energy level and those in the complete d and f sublevels |
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| The energy required to remove an electron from an atom or ion in its gaseous state |
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| The energy charge associated with the gaining of an electron by an atom in its gaseous state |
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In a Lewis Structure where the central atom has more than 8 valence electrons |
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| The charge that an atom in a Lewis Structure would have if all the bonding electrons were shared equally between the bonded atoms |
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| The actual structure of a molecule that is intermediate between two or more resonance structures |
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| The bond formed when a ligand donates electrons to an empty orbital of a metal in a complex ion |
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| A covalent bond between two atoms with significantly different electronegativity values, resulting in an uneven distribution of electron density |
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| covalent bond in which electronegativity values are smilar |
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| A measure of the separation of positive and negative charges in a molecule |
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| An advanced model of chemical bonding in which electrons result in quantum-mechanical orbitals localized on individual atoms that are a hydridized blend of standard atomic orbitals; chemical bonds result from an overlap of thse orbitals |
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| 1/2 (bonding electrons - antibonding electrons) in MO theory |
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| An intermediate force between an ion and the oppositely charged end of a polar molecule |
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| An intermediate force exhibited by polar molecules that results from the uneven charge distribution |
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| ion-induced dipole forces |
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| An intermediate force between an ion and the oppositely charged end of a polar molecule |
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dispersion forces (London forces) |
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| Intermediate force that is present in all molecules |
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| A strong dipole-dipole attractive force between a hydrogen bonded to O,N,or F and one of these electronegative atoms on a neighboring molecule |
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| The phase transition from liquid to gas |
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| The phase transition from gas to liquid |
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| The phase change from solid to liquid |
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| the phase change from liquid to solid |
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| The phase transition from solid to gas |
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| The phase change from gas to solid |
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| The partial pressure of a vapor in dynamic equilibirum with a liquid |
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| The pressure required to bring about a transition to a liquid at the critical temperature |
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| The temperature above which a liquid cannot exist, regardless of pressure |
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| The temperature and pressure above which a supercritical fluid exists |
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| The unique set of coordinates at which all three phases of a substance are equally stable and in equilibrium |
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| The map of the phase of a substance as a function of pressure and temperature |
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Mega (M) = 10^6 Kilo (k) = 10^3 Deci (d) = 10^-1 Centi (c) = 10^-2 Milli (m) = 10^-3 micro (u) = 10^-6 nano (n) = 10^-9 pico (p) = 10^-12 |
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E = hv h = plank's constant (6.626 x 10^-34) |
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| Restrictions on the quantum numbers |
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moles of solute liters of solution |
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Initial and Fianl Gas Problems |
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Determining Oxidation State |
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| Paramagnetic vs. Diamagnetic |
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Paramagnetic = there are unpaired orbitals Diamagnetic = all the orbitals are paired up |
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Atomic radius decreases across a period. Atomic radius increases down a group |
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The radius of a cation is much smaller than that of the corresponding atom. The radius of an anion is much larger than that of the corresponding atom. |
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| Trends in Ionization Energy |
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Ionization energy increases acorss a period. Ionization energy decreases down a group. |
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| Trends in Electronegativity |
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Electron affinity increases across a period. Electron affinity decreases down a group. |
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| Using electronegativity to pedict whether a bond is ionic, polar covalent, or nonpolar covalent. |
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Ionic = elements with very different electronegativities (more than 2) Nonpolar covalent bonds = elements with very similar electronegativities Polar covalent bond = those with intermediate electronegativity (less than 2) Florine is the most electronegative element |
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The formal charge of an atom in a Lewis Structure is the charge the atom would have if all bonding electrons were shared equally between bonding atoms. Formal charges should add up to be zero. |
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| Single bonds are longer than double bonds and double bonds are longer than triple bonds. |
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| Triple bonds are stronger than double bonds and dobule bonds are stronger than single bonds. |
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- In its compounds, the oxidation state of flourine is -1 - In its compounds, the oxidation state of hydrogen is +1 (except when hydrogen is bonded to metals, its oxidation state is -1) -In its compounds, the oxidation state of oxygen is +2 (when oxygen atoms are bonded to each other, its oxidation state is -1) - Group 7A = -1 - Group 6A = -2 - Group 5A = -3 |
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Single bond = 1 sigma bond Double bond = 1 sigma + 1 pi Triple bond = 1 sigma + 2 pi |
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Dispersion Forces: The weakest of the intermolecular forces, are present in all molecules and atoms and increase with increasing with increasing molar mass. These forces are always weak in small molecules but can be significant in molecules with high molar masses. Dipole-Dipole Forces: present in polar forces Hydrogen Bonds: The strongest of the intermolecular forces that can occur in pure substances (second only to ion-dipole forces), are present in molecules containing hydrogen bonded directly to flourine, oxygen, or nitrogen Ion-dipole Forces: present in mixtures of ionic compounds and polar compounds. These are very strong and are especially important in aqueous solutions of ionic compounds. |
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- The rate of vaporization increases with increasing temperature The rate of vaporization increases with increasing surface area - The rate of vaporization increases with decreasing strength of intermolecular forces |
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