Gen Chem 1 Test Questions – Flashcards
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Unlock answersintensive property |
A property such as density that is independent of the amount of the given substance |
| extensive property |
| A property that depends on the amount of given substance, such as mass |
| isotope |
| atoms of the same element with the same number of protons but different numbers of neutrons and consequently different masses |
| mass number |
| (A) the mass of the element |
| atomic number (Z) |
The atomic number of the element (Z) that determines the number of protons and electrons in an atom |
Natural abundance (percent abundance) |
| the relative percentage of a particular isotope in a naturally occuring sample with respect to other isotopes of the same element |
| anode |
| The electrode is an electrochemical cell where oxidation occurs electrons flow away from the anode |
| cathode |
| The electrode is an electrochemical cell where reduction occurs, electrons flow toward the cathode |
| alkali metals |
| Highly reactive metals in Group 1A of the periodic table |
| alkaline earth metals |
| Fairly reactive metals in group 2A of the periodic table |
| Law of Conservation of Mass |
| A law stating that mass is neither created nor destroyed in a chemical reaction |
Law of Definite Proportion |
| A law stating that all samples of a given compound have the same proportions as their constituent elements |
Law of Multiple Proportions |
| 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 |
oxoanions |
| A polyatomic atom containing a nonmetal covalently bonded to one or more oxygen atoms |
| formula unit |
| The smallest electrically neutrol collection of ions in an ionic compound |
| empirical formula |
| A chemical formula that shows the simplest whole number ratio of atoms in a compound |
| molecular formula |
The chemical formula that shows the actual number of atoms of each element in a molecule of a compound |
oxidation number |
| 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 |
limiting reagent |
| The reactant that has the smallest stoichiometric amount in a reactant mixture and consequently limits the amount of product in a chemical reaction |
| reducing agent |
| A substance that causes the reduction of another substance ; a reducing agent loses electrons and is oxidized |
| reactant in excess |
| The reactant where there is some left over at the end of the reaction |
| partial pressure |
| The pressure due to any individual component in a gas mixture |
| ideal gas |
| The proportionality constant of the ideal gas law (r = .0821) |
| ideal gas law |
| 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. |
| kinetic molecular theory |
| A model of an ideal gas as a collection of point particles in constant motion undergoing completely elastic collisions |
electromagnetic radiation |
| A form of energy embodied in oscillating electric and magnetic field |
| emission spectrum |
| The range of wavelengths emitted by a particular element used to identify the element |
| photon |
| The smallest possible packet of electromagnetic radiation with an energy equal to hv |
| photoelectric effect |
| the observation that many metals emit electrons when light falls upon them |
| de Broglie wavelength |
| The observation that the wavelength of a particle is inversely proportional to it's momentum |
Heinsburg Uncertainty Principle |
| 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 |
| principal quantum number (n) |
| 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 |
| angular momentum quantum number (l) |
| An integer that determines the shape of an orbital |
| magnetic quantum number (m) |
| An integer that specefies the orientation of an orbital |
| principle shells (levels) |
| The group of orbitals with the same value of n |
| sublevels (subshells) |
| Those orbitals in the same principle level with the same value of n and l |
| The purpose of line spectra |
| Spectra are used to identify elements |
| Bohr's explanation of spectral lines: |
| Light has energy equal to the difference in energy of the electron in two orbitals |
| Electrons possess.... |
| both wave and particle properties |
| Spin quantum number (ml) |
| The fourth quantum number, which denotes the electrons spin order as 1/2 (up arrow) or -1/2 (down arrow) |
| Pauli Exclusion Principle |
| The principle that no two electrons in an atom can have the same four quantum numbers |
| Hund's Rule |
| The principle stating that when electrons fill degenerate orbitals, they first fill them singly with parallel spins |
| Aufbau Principle |
| The principle that indicates the pattern of orbital filling in an atom |
| electron configuration |
| A notation that shows the particular orbitals that are occupied by electrons in an atom |
Effective Nuclear Charge (Zeff) |
| The actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons |
metallic radius |
| The way of expressing radius for metals |
| covalent radius (bonding atomic radius) |
| 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 |
| core electrons |
| those electrons in a complete principle energy level and those in the complete d and f sublevels |
ionization energy |
| The energy required to remove an electron from an atom or ion in its gaseous state |
| electron affinity |
| The energy charge associated with the gaining of an electron by an atom in its gaseous state |
| expanded octet |
In a Lewis Structure where the central atom has more than 8 valence electrons |
| formal charge |
| The charge that an atom in a Lewis Structure would have if all the bonding electrons were shared equally between the bonded atoms |
| resonance hybrid |
| The actual structure of a molecule that is intermediate between two or more resonance structures |
| coordinate covalent bond |
| The bond formed when a ligand donates electrons to an empty orbital of a metal in a complex ion |
| polar covalent bond |
| A covalent bond between two atoms with significantly different electronegativity values, resulting in an uneven distribution of electron density |
| nonpolar covalent bond |
| covalent bond in which electronegativity values are smilar |
| dipole |
| A measure of the separation of positive and negative charges in a molecule |
valence bond theory |
| 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 |
Bond Order |
| 1/2 (bonding electrons - antibonding electrons) in MO theory |
| Ion-dipole forces |
| An intermediate force between an ion and the oppositely charged end of a polar molecule |
| dipole-dipole forces |
| An intermediate force exhibited by polar molecules that results from the uneven charge distribution |
| ion-induced dipole forces |
| An intermediate force between an ion and the oppositely charged end of;a polar molecule |
dispersion forces (London forces) |
| Intermediate force that is present in all molecules |
| hydrogen bonding |
| 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 |
| vaporization |
| The phase transition from liquid to gas |
| condensation |
| The phase transition from gas to liquid |
Melting |
| The phase change from solid to liquid |
| freezing |
| the phase change from liquid to solid |
| sublimation |
| The phase transition from solid to gas |
| deposition |
| The phase change from gas to solid |
| vapor pressure |
| The partial pressure of a vapor in dynamic equilibirum with a liquid |
| critical pressure |
| The pressure required to bring about a transition; to a liquid at the critical temperature |
| critical temperature |
| The temperature above which a liquid cannot exist, regardless of pressure |
| critical point |
| The temperature and pressure above which a supercritical fluid exists |
| Triple Point |
| The unique set of coordinates at which all three phases of a substance are equally stable and in equilibrium |
| phase diagram |
| The map of the phase of a substance as a function of pressure and temperature |
| The prefixes: |
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 |
| Enery of a photon |
E = hv h = plank's constant (6.626 x 10^-34)
|
| Restrictions on the quantum numbers |
| Shapes of the orbitals |
| Avogadro's Number |
| 6.022 x 10^23 particles |
| Strong Acids |
HCl HBr HI HNO3 H2SO4 |
| Strong Bases |
Metal salts of.... OH O2- |
| Molarity (M) |
moles of solute liters of solution |
| The Ideal Gas Law |
| PV = nRT |
Initial and Fianl Gas Problems |
| P1V1 = P2V2 |
| Sig Fig Rules |
Determining Oxidation State |
| Properties of a Gas |
| Paramagnetic vs. Diamagnetic |
Paramagnetic = there are unpaired orbitals ; Diamagnetic = all the orbitals are paired up |
| Trends in Atomic Radius |
Atomic radius decreases across a period. ; Atomic radius increases down a group |
| Trends in Ionic Radius |
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. |
| Trends in Ionization Energy |
Ionization energy increases acorss a period. ; Ionization energy decreases down a group. |
| Trends in Electronegativity |
Electron affinity increases across a period. ; Electron affinity decreases down a group. |
| Using electronegativity to pedict whether a bond is ionic, polar covalent, or nonpolar covalent. |
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 |
| Formal Charge |
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. |
| Bond Length |
| Single bonds are longer than double bonds and double bonds are longer than triple bonds. |
| Bond strength |
| Triple bonds are stronger than double bonds and dobule bonds are stronger than single bonds. |
| Atoms with one bond: |
H F Cl Br I |
| Atoms with two bonds: |
O S Se Te Be |
| Atoms with three bonds: |
N P As B Al |
| Atoms with four bonds: |
C Si Ge |
| Oxidation Number: |
- 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 |
| Sigma vs. Pi bonds |
Single bond = 1 sigma bond Double bond = 1 sigma + 1 pi Triple bond = 1 sigma + 2 pi |
| Properties of A Liquid |
| Intermolecular Forces |
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. |
| Trends in Vaporization |
- 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 |
