Section 1
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Neutralization Reactions
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Mar 1, 2020
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Section 1
(50 cards)
Neutralization Reactions
Special type of double replacement reaction in which one reactant is an acid and the other is a base. The products are a salt and water
Metalloids
Metals bordering the staired line separated the metals from the nonmetals. These exhibit some of the properties of metals and some of nonmetals
Complete Electron Configuration (i.e. electronic configuration)
(shell #)(type of orbital)^(# of electrons) E.g. F = 1s^2 2s^2 2p^5
Periodic Chemical Properties of Binary Compounds
Metal oxides of groups 1 & 2 are basic--form hydroxides in solution Non-metal oxides in groups 14-17 are acidic--form acids in solution
Arrhenius Equation
K = Ae^(Ea/RT)
Anions
Atoms that have gained one or more electrons and carry a negative charge (-). Increased mutual repulsion makes anions larger than their parent atoms. Effective nuclear charge decreases
Single Replacement Reactions
An element may react with a compound to produce a different element and new compound
Atomic Radii
Going across the row, the radii gets smaller - the more electrons, the more they go closer to the nucleus; going down a column, the radii gets larger - there is another electron level. This is the only trend that doesn't increase to the top right
Highly Flammable Compounds
Organic Compounds.
Kinetic Molecular Theory
Ideal gases: Random Motions, Elastic Collisions, Volume is negligible, No Attraction or Repulsion, Energy Proportional to Temperature; When gases are put under pressure and cooled, they become real and this theory no longer applies
Decomposition Reactions
A large molecule decomposes into a elements or into smaller molecules
Number of Neutrons
The atomic mass minus the atomic number
Pauli Exclusion Principle
An orbital can hold only two electrons, and they must have opposite spins
Halogens
The 2nd to last column
Ksp
The solubility product, denotes the fact that the equilibrium is between the solid and its solution products
Graham's Law of Effusion
√(m1/m2) = √(vrms2/v4ms1)
Synthesis Reactions
Reactions of two or more elements to form a compound
Atomic Symbol placements
The element's symbol is in the middle, the charge is in he upper right hand corner, the atomic number is in the lower left-hand corner, and the mass number is in the upper left hand corner.
Addition Reactions
A simple molecule or an element is added to another molecule
Condensation
The conversion of a gas into a liquid
Number of Protons
Atomic Number
Effective Nuclear Charge (Z_eff)
# of protons in nucleus - shielding effect of electrons (the extent to which the inner electrons protect the outer electrons from the + charge of the nucleus)
Typical Ionic Charges
Group 1 -- +1 Group 2 -- +2 Group 13 -- +3 Group 15 -- -3 Group 16 -- -2 Group 17 -- -1
Maximum Number of Electrons in Each Shell
2(n^2)
Ionization Energy
The energy required to remove the least tightly held electron in a given species. Determined by Coulomb's Law: F = (q1 * q2)/(r^2). Measured in kj mol^-1 and are always positive (endothermic). Going across the row, the ionization energy increases; going down the column the ionization energy decreases
Electron Spin Quantum Number (m_s)
An electron can have either a clockwise spin or counterclockwise spin (+ 1/2 or - 1/2)
Alkali Metals
First column on the left
Accuracy v. Precision
Accuracy - closeness to the true value. Precision - closeness of a group of repeated measurements
Cations
Atoms that have lost one or more electrons and carry a positive charge (+). Decreased mutual repulsion makes cations smaller than their parent atoms. Effective nuclear increases
Heisenberg Uncertainty Principle
The position and momentum of an electron can never be simultaneously, exactly known
Orbital Notation
Uses boxes to represent the orbitals and arrows to represent electrons (remember: s can have 1 orbital, p can have 3, d can have 5, and f can have 7)
Formation Reactions
Same as synthesis except the product must have a coefficient of 1.
Principal Quantum Number (n)
related to the size and energy of the orbital. As n increases, the orbital becomes larger and and the electron spends more time farther from the nucleus
Noble Gas Configuration
Write the previous noble gas in brackets and then continue with complete electron configuration
Isoelectronic
Atoms and Ions that have identical electron configurations
Combustion Reactions
An Organic compound reacts with oxygen to form carbon dioxide and water
Noble Gases
The last column on the table.
Electron Affinity
The energy change that accompanies the addition of an electron to an atom. Increases left to right and slightly increases bottom to top
Allotrope
One or more distinct forms of an element; classification as an allotrope is based on structure and physical properties
Number of Electrons
Atomic Number if the charge is 0
Electronegativity
The ability of an atom to attract electrons to itself within a covalent bond. Increases left to right and bottom to top
Kf
The formation constant, a special term used to describe the formation of complex ions
Double Replacement Reactions
Two compounds react and the cation in one compound replaced the cation in the second compound and vice versa
Transition elements
In the middle of the periodic table.
Photoelectric Effect
The use of high energy X-rays or UV photons to eject electrons from atoms
Molecular Formula Multiplier
molar mass / mass of empirical formula
Angular Momentum Quantum Number (l)
Integral values from 0 to n-1 where s=0, p=1, d=2, and f=3; relate to the shape of atomic orbitals
Magnetic Quantum Number (m_l)
Integral values between l and -l (including 0); related to the orientation of the orbital in the space relative to the other orbitals of the atom. This gives the number of possible orbitals for each sub-shell (s, p, d, or f): s can only have 1 (m_l = 0), p can have 3 (m_l = -1, 0 1), d can have 5 (m_l = -2, -1, 0, 1, 2), and f can have 7 (m_l = -3, -2, -1, 0, 1, 2, 3)
Hund's Rule of Maximum Multiplicity
If there is more than one orbital with the same energy (degenerate orbitals), then 1 electron is placed into each orbital before any pairing takes place. All similar orbitals have a similar energy
Alkaline Earth Metals
Second Column on the left
Section 2
(50 cards)
Substintitutial Alloys
atoms of a comparable size to pure metal, substitute themselves for those pure metal atoms; less malleable
Boyle's Law
Pressure is inversely proportional to volume
Valence Bond Theory
Orbitals overlap to form bonds between atoms; one e- from each of the bonded atoms accomodates the overlapping orbital; both e- in the overlap are attracted to the nucleus of each atom
Covalent Network Solids
High melting and boiling point; rigid and hard; no conduction from one layer to another
Arrhenius Base
a substance that releases OH- ions when placed in water
1 lone pair
Endothermic reactions
absorb energy from the surroundings; delta H is positive; breaking bonds is endothermic
STP
1 mol = 22.4 L
Oxidation states
the charges that an atom actually has when in an ionic compound or the charge that the atom would have if the compound were ionic
Dalton's Law
The total pressure of a mixture of ideal gases is the sum of their partial pressures
Spectator Ions
ions that are found in exactly the same form on both sides of the equation
Dipole-Dipole
Intermolecular Coulombic attractions between permanent dipoles in adjacent molecules; molecules with polar bonds (caused by differences in electronegativity) and dipoles that do not cancel out via symmetry will have permanent dipoles; intermediate in strength
Interstitial Alloys
small atoms fill the spaces (/interstices) between the bigger atoms, making the alloy more rigid and less malleable than the pure metal
Covalent Bonds
When two close, non-metal atoms share one or more pairs of valence e- (intramolecular); identical electronegativities -> nonpolar; different electronegativities -> polar (more electronegative atom will develop a - partial charge
Beer-Lambert Law
Used to asses the concentration of a solution by relating the absorbance of the solution to the concentration of the solution. A = abc where A = absorbance, a = absorptivity coefficient, b = path length, and c = concentration
Work done in a cylinder and piston system
P * delta V
Standard Enthalpy of Fusion
the energy change when 1 mol of a substance is converted from a solid to a liquid
Heating and Cooling Curves
Polyprotic acids
can donate more than one proton per molecule
Arrhenius Acid
a substance that contains H and releases H+ ions (must be aqueous)
Bronsted-Lowry Base
proton accepto
Rate Equation
Rate of reaction = k [A]^x [B]^y [C]^z
Gay-Lussac's Law
Pressure is directly proportional to temperature
Hydrogen Bonds
Strong intermolecular bonds formed when H is bonded to a more electronegative, and relatively small element
Bronsted-Lowry Acid
proton donor
Sigma Bonds
Single bonds formed by overlap along the internuclear axis of the bond: overlap of 2 s orbitals, overlap of an s and p orbital, or overlap of 2 p orbitals
Chromatography
The process of separation of components based upon their relative affinities (intermolecular attraction) for a stationary phase and a moving phase
Monoprotic acids
can only donate one proton per molecule
Cathode
reduction occurs here; e- in reactants; e- flow from anode to here
Specific Heat Capacity (C_p)
(quantity of heat supplied) / (mass of object * temp change)
Ionic solids
low vapor pressure, high melting point, and high boiling point bc of ionic lattices; brittle; only conduct electricity when molten or in solution
Exothermic reactions
release energy to the surroundings; delta H is negative; making bonds is exothermic
Amphiprotic
can act either as an acid or a base (e.g. water)
Concentration cell
e- flow from the lower concentration half-cell to the higher concentration half-cell in order to achieve equilibrium
Capillary Action
Combination of cohesive forces within the liquid itself and adhesive forces between the liquid and the walls of the tube
Charles' Law
Volume is directly proportional to the Kelvin temperature
Avogadro's Law
Volume is directly proportional to the number of moles of gas present
Anode
oxidation occurs here; e- in products; e- flow from here to cathode
London Dispersion Forces
Small intermolecular Coulombic forces that are caused by the movement of electrons within the covalent bonds of molecules that would otherwise have no permanent dipole; increase with increasing # of electrons and surface area; exist between all molecules, but are really weak
Collision Theory
Reacting molecules must collide with sufficient energy and correct orientation to be successful; most are not
Molecular Solids
low melting points; soft; non-conductors
Pi Bonds
Occur in double and triple bonds; formed by sideways overlap above and below the internuclear axis of the bond; overlap of 2 p orbitals; double bonds contain 1 sigma and 1 pi; triple bonds contain 1 sigma and 2 pi
Metallic Bonds
Valence e- are delocalized and are mobile around a regular arrangement of cations--sea of e-; leads to ductility, malleability, and thermal and electrical conductivity
Infrared Spectroscopy
Used to identify the presence of different types of covalent bonds present in larger molecules
Neutralization Reaction
Arrhenius acid and base react to produce a salt and water
Formal Charge on an atom within a Lewis Structure
(# of valence e- around that atom in the free atom) - (# of non-bonding e- around that atom in a Lewis Structure) - 0.5(# of bonding e- around that atom in a Lewis structure)
Maxwell-Boltzman Distribution
Particles at a higher temperature have greater kinetic energies
Mass Spectrometry
Used to identify the varying masses of the isotopes in a naturally occurring sample of any given element. The process goes: ionization (always +), accelerations, defection, and then detection
0 lone pairs
Ionic Bond
When valence e- are transferred from 1 atom to another; between metals and non-metals that aren't close (intramolecular)
Section 3
(22 cards)