Chemistry Lecture 1
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114 terms
Terms | Definitions |
|---|---|
 atoms | all mass consist of these tiny particles |
| protons and neutrons | these are equal in size and mass; one having positive charge and the other having a neutral charge |
| neutral | an atom has a ____ charge; having equal number of protons as electrons |
| elements | these are the building blocks of all compounds and cannot be decomposed into simpler substances via chemical means. |
| A | the mass number is represented by this letter; protons + neutrons |
| Z | the atomic number or number of protons is represented by this letter; this is also the identity number of the element |
| isotopes | two or more atoms of the same element that contain different numbers of neutrons; they have similiar chemical properties |
| atomic weight or molar mass (MM or M) | this is given in atomic mass units; this is the actual mass of an element |
| mole | aka Avogardo's number: 6.022 X 10^23 is the number of carbon atoms in 12 grams of 12C |
| 1 | 6.022 x 10^23 amu = _ gram(s) |
| dalton | Biochemists refer to amu as a ___. |
| moles= | grams/atomic or molecular weight |
| period | each horizontal row of the periodic table |
| groups or families | each vertical column of the periodic table |
| metals | large atoms that tend to lose electrons to form positive ions or positive oxidation states; |
| metallic character | ductility, malleability, thermal and electrical conductivity, and luster |
| ductility | easily stretched |
| malleability | easily hammered into thin strips |
| mercury | all metals except this one exist as solids in room temperature |
| metals | typically form ionic oxides |
| BeO | this is one exception; not an ionic oxide |
| nonmetals | these are diverse in appearance and chemical behaviors; they form negative ions |
| nonmetals | molecular substances are typically made from only ___. |
| nonmetals | form covalent oxides |
| family or group | elements in the same ___ on the periodic table tend to have similar chemical properties; tending to make the same number of bonds, and exist as similiarly charged ions. |
| hydrogen | unique and its chemical and physical characteristics do not conform well to any family; its a nonmetal; under most conditions it is colorless, odorless, diatomic gas. |
| alkali metals | as pure substances they are soft metallic solids with low densities and low melting points; highly reactive; reacting with most nonmetals to form ionic compounds; in nature they exist only as compounds. |
| alkaline earth metals | form 2+ cations |
| 4A | these elements can form four covalent bonds with nonmetals; can form two additional bonds with Lewis bases |
| carbon | forms strong pi bonds to make strong double and even triple bonds (group 4A element) |
| 5A | these elements can form 3 covalent bonds; can form 5 covalent bonds by using their d orbitals; can further bond with a Lewis base to form a 6th covalent bond; can not make pi bonds |
| nitrogen | forms strong pi bonds to make double and triple bonds (group 5A element); can also form 4 covalent bonds by donating its lone pair of electrons to form a bond |
| weak; double | Phosphorus can form only ___ pi bonds to make ___ bonds. |
| 6A | called chalcogens |
| oxygen | the second most electronegative element; divalent; can form strong pi bonds to make double bonds; |
| S8 | the most common form of pure sulfur is the yellow solid; metal sulfides such as Na2S are the most common form of sulfur found in nature; sulfur can form 2,3,4, or even 6 bonds; it has the ability to pi bond making strong double bonds |
| 7A | radioactively stable group of elements; called the halogens; they are highly reactive; tend to gain electrons; can take on oxidation states as high as +7 when bonding to highly electronegative atoms like oxygen |
| flourine and chlorine | diatomic gases at room temperature |
| bromine | diatomic liquid at room temperature |
| iodine | diatomic solid at room temperature |
| flourine | when in compounds this element always has an oxidation state of -1; meaning it only makes one bond |
| gaseous hydrogen halides | hydrogen combines with all the halogens to form these; they are water soluble forming hydrohalic acids |
| ionic halides | halogens react with metals to form these. |
| noble gases | nonreactive sometimes called inert gases; the only elements normally found in nature as isolated atoms; they are all gases at room temperature |
| hydrogen, oxygen, nitrogen, and the halogens | the elements that tend to exist as diatomic molecules |
| size | the ___ of an atom has a significant effect on its chemistry |
| small atoms | have less room to stabilize charge by spreading it out; they don't have d orbitals avaiable to them for bond formation, thus they cannot form more than 4 bonds |
| large atoms | can't form pi bonds easily |
| ion | when an element has more or fewer electrons than protons |
| anions | nonmetals form these negative ions |
| cations | metals form these positive ions |
| cations | significantly smaller than their neutral atom counterparts |
| anions | significantly larger than their neutral atom counterparts |
| Coulomb's Law | this describes the electrostatic forces holding an electron to its nucleus; the distance between the electron and the nucleus is r. q1 we might plug in the positive charge of the nucleus; Z for q2 (the charge of the given electron). F=kq1q2/r^2 |
| shield | sometimes the first electron will___some of the nuclear charge from the second electron so that the second electron doesn't feel the entire nuclear charge, Zeff. |
| effective nuclear charge (Zeff) | the amount of charge felt by the second electron; increases when moving from left to right, each additional electron is pulled more strongly toward the nucleus |
| atomic radius | with each added shell the atom grows larger; this increases from the top of the periodic table to the bottom. |
| ionization energy | when an electron is more strongly attached to the nucleus, more energy is required to detach it; the energy needed is called this; generally increasing from left to right and from bottom to top |
| first ionization energy | the energy necessary to detach an electron from a neutral atom |
| second ionization energy | the energy for the removal of a second electron from the same neutral atom; it is always much greater than the first because when one electron is removed, the effective nuclear charge on the other electrons increases. |
| electronegativity | the tendency of an atom to attract an electron in a bond that it shares with another atom; the most commonly used form of measurement of this is the Pauling scale that ranges from .79 for cesium to 4.0 for flourine; tends to increase left to right and bottom to top; its related to Zeff similiar to ionization energy; the values for noble gases are undefined |
| electron affinity | the willingness of an atom to accept an additional electron; the energy released when an electron is added to a gaseous atom; tends to increase left to right and from bottom to top and is related to Zeff; |
| metallic character | tends to increase from right to left and top to bottom |
| ampere | electric current is measured with this unit |
| candela | luminous intensity is measured with this unit |
| 1 N= | 1 kg m s^-2 |
| 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 |
| Femto-(f) | 10^-15 |
| bond length | the point where the energy level is the lowest; two atoms will only form a bond if they can lower their overall energy level by doing so; nature tends to seek the lowest energy state |
| bond dissociation energy or bond energy | the energy necessary to achieve a complete separation of a chemical bond |
| compound | a substance made from two or more elements in definite proportions |
| empirical formula | in all pure compounds, the relative number of atoms of one element to another can be represented by a ratio of whole numbers |
| molecules | in molecular compounds, groups of atoms form these repeated, separate and distinct units |
| molecular formula | in molecular compounds, the exact number of elemental atoms in each molecule can be represented by this. |
| ionic compounds | named after their cation and anion; if the cation is metal and capable of having different charges then its name is followed by a Roman numeral in parentheses; an older method for naming cations is to add -ic to the ending of the cation with the greater positive charge and -ous to the cation with the smaller charge; if the cation is made from a nonmetal, the cation name ends in -ium |
| monatomic anions and simple polyatomic anions | given the suffix -ide [example: (H-), hydride ion] |
| polyatomic anions with multiple oxygens | end with the suffix -ite or -ate depending upon the relative number of oxygens; the more oxygenated species will use the -ate suffix. [example: (NO2-), nitrite ion] |
| polyatomic anions with multiple oxygens | if there are more possibilities, the prefixes hypo-, per- are used to indicate fewest and most oxygens. [example: (ClO-), hypochlorite] |
| acids | named based on their anions; if the name of the anion ends in -ide, then the name will start with hydro- and end in -ic |
| oxyacid | if the acid is an/a ____, the ending -ic is used for the species with more oxygens and -ous for the species with fewer oxygens. [examples: (H2SO4), sulfuric acid and (H2SO3), sulfurous acid] |
| binary molecular compounds | compounds with only two elements, the name begins with the name of the element that is farthest to the left and lowest in the periodic table; the name of the second element is given the suffix -ide and a Greek number prefix is used on the first element if necessary [example: (N2O4), dinitrogen tetroxide] |
| physical reactions | when a compound undergoes a reaction and maintains its molecular structure: melting, evaporation, dissolution, and rotation of polarized light |
| chemical reactions | when a compound undergoes a reaction and changes its molecular structure to form a new compound: combustion, metathesis, and redox |
| limiting reagent | reactant that would be completely used up if the reaction were to run to completion |
| theoretical yield | the amount of product produced when a reaction runs to completion |
| percent yield | actual yield divided by the theoretical yield, times 100 |
| Combination Reaction | A + B-->C |
| Decomposition | C-->A + B |
| Single Displacement or Single Replacement | A + BC--> B + AC |
| Double Displacement or Double Replacement or Metathesis | AB + CD--> AD + CB |
| quantum mechanics | elementary particles can only gain or lose energy and certain other quantities in discrete units |
| quantum numbers | a set of four numbers is the address or ID number for an electron in a given atom; no two electrons in the same atom can have the same 4 numbers |
| principal quantum number, n | the first quantum number designates the shell level, the larger the number the greater the size and energy of the electron orbital; given by the period number, for transition metals is one shell behind the period, for the lanthanides and actinides lags two shells behind the period. |
| valence electrons | the electrons which contribute most of an element's chemical properties, are located in the outermost shell of an atom |
| azimuthal quantum number, l | the second quantum number designates the subshell; the orbital shapes s,p,d, and f; if this number = 0 we are in the s subshell; this number = n-1; each subshell has a peculiar shape to its orbitals; the shapes are based on probability functions of the position of the electron |
| magnetic quantum number, ml | the third quantum number designates the precise orbital of a given subshell; each subshell will have orbitals with this quantum numbers from -l to +l; thus for the first shell with n=1 and l=0 this quantum number =0 |
| electron spin quantum number, ms | the fourth quantum number has values of -1/2 or +1/2; any orbital can hold up to two electrons and no more; if two electrons occupy the same orbital they have the same first three quantum numbers |
| The Pauli exclusion principle | this states that two electrons in the same atom can not have the same four quantum numbers; because two electrons in the same orbital have identical 1st, 2nd, and 3rd quantum numbers, they must have opposite electron spin quantum numbers |
| Heisenberg Uncertainty Principle | this arises from the dual nature (wave-particle) of matter; it states that there exists an inherent uncertainity in the product of the position of a particle and its momentum, and that this uncertainty is on the order of Planck's constant. [delta x (delta p) = h |
| Planck's constant | 6.63 X 10^-34 J s |
| Aufbau principle | this states that with each new proton added to create a new element, a new electron is added as well. |
| ground state | the electron configurations for atoms whose electrons are all at their lowest energy level |
| transition metals | for these elements, ions are formed by losing electrons from the subshell with the highest principle quantum number first; their electron configurations are not always that easy to predict due to degenerate orbitals |
| electrons | these can momentarily absorb energy and jump to a higher energy level creating an atom in an excited state |
| Hund's rule | the mutual repulsion between two particles creates an increase in potential energy; explaning why two electrons can fit into one orbital; electrons will fill any orbital in the same subshell until all orbitals in that subshell contain at least one electron, and the unpaired electrons will have parallel spins |
| Planck's quantum theory | the father of quantum mechanics demonstrates that electromagnetic energy is quantized (comes only in discrete units related to the wave frequency); if we transfer energy from one point to another via an electromagnetic wave, and we wish to increase the amount of energy that we are transferring without changing the frequency, we can only change the energy in discrete increments given by : delta E= hf |
| photoelectric effect | electrons are emitted from matter as a consequence of their absorption of energy from electromagnetic radiation of very short wavelength, such as visible or ultraviolet light. |
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