Chemistry 111 Final Exam

the scientific discipline that studies the composition, properties, and transformations of matter
homogeneous mixture of more than one substance; can have gas, liquid, or solid solution
simplest substance that has a distinct chemical identity, cannot be broken down further; pure substance
consists of 2 or more elements held together by chemical bonds
Homogeneous Mixture
same throughout; example clean air, wine, brass
Heterogeneous Mixture
contains regions that are different in structure and property; example rocks, particles in air
Physical Property
property which can be measured without changing the identity of the substance; example color, odor, density changes of state
Chemical Property
property which describes how a substance changes its identity to form other substances; example burning coal
Intensive Property
does not depend on the amount of the substance present; example color, density
Extensive Property
depends on the amount of substance present; example mass, volume, length
Equivalence Statement
an equation defining units where the two units are set equal
Conversion Factor
a ratio relating the same quantity in two systems of units that is used to convert the units of measurement
has mass and occupies space
definite shape, definite volume
no definite shape, definite volume
no definite shape, no definite volume
Pure Substance
Element or compound
homogeneous or heterogeneous
separation by state
separation using boiling points of two substances in the mixture
separation using how well a substance adheres (sticks) to another substance, absorption
meters squared, m2
meters cubed, m3
1 kilogram
2.205 pounds
1 inch
2.54 centimeters
1 Liter
1 decimeter cubed, 1 dm3
1 Liter
1.057 quarts
Degree Kelvin
Degree Celcius + 273.15
mass/volume, g/cm cubed
Atomic Number
the number of protons in the nucleus of an atom of an element
Mass Number
the sum of the number of protons and neutrons in the nucleus of a particular atom
atoms of the same element (same number of protons) with different numbers of neutrons in the nucleus
Peta (P)
1015, 1 Pm = 1×1015m
Tera (T)
1012, 1 Tm = 1×1012 m
Giga (G)
109, 1 Gm=1×109 m
Mega (M)
106, 1Mm = 1×106 m
Kilo (k)
103, 1 km = 1×103 m
Deci (d)
10-1, 1 dm = 1×10-1 m
Centi (c)
10-2, 1 cm = 1×10-2 m
Milli (m)
10-3, 1 mm = 1×10-3 m
Micro (u)
10-6, 1 um = 1×10-6 m
Nano (n)
10-9, 1 nm = 1×10-9 m
Pico (p)
10-12, 1 pm = 1×10-12 m
Femto (f)
10-15, 1 fm = 1×10-15 m
Atto (a)
10-18, 1 am = 1×10-18 m
Zepto (z)
10-21, 1 zm = 1×10-21 m
Molecular Formula
indicates the active number and kind of atoms in a molecule
Emperical Formula
gives the relative number and kind of atoms in a compound shown as lowest whole number ratio
Structural Formula
gives arrangement of atoms in a compund
combination of more than one non-metallic elements
a positively charged atomic species, missing an electron
a negatively charged atomic species, an extra electron
Polyatomic Ion
an electrically charged group of two or more atoms
Ionic Bond
force of attraction between oppositely charged ions
Covalent Bond
the attraction involving the sharing of electrons between two nuclei
polyatomic anions which contain oxygen
Oxidation Number
the apparent charge on an individual atom in a compound or a polyatomic unit
invisible particles called “atamos”
Robert Boyle
1661, definition of element
Joseph Priestly
1774, isolation of oxygen
Antoine Lavoisier
1789, law of conservation of mass
Joseph Proust
1806, law of definite proportions
John Dalton
1808, postulates of atomic theory and law of multiple proportions
Joseph Gay-Lussac
1809, determining chemical formulas
Amedeo Avogadro
1811, equal volumes of gases contain an equal number of particles
Jon-Jakob Berzelius
early 1800s, symbols for elements
Dmitri Mendeleev
1869, periodic table
Wilhelm Roentgen
1896, discovered xrays
Henri Becquerel
1896, proposed concept of radioactivity
J. J. Thomson
1903, measured charge-to-mass ratio of an electron
Robert Millikan
1909, oil drop experiment to measure the charge of an electron
Ernest Rutherford
1911, proved existence of nucleus of the atom
Henry Mosley
1913, devised the atomic number system
First Postulate
each element is composed of atoms
Second Postulate
atoms of the same element are identcal
Third Postulate
atoms of different elements are different in some fundamental way and have different properties
Fourth Postulate
atoms are neither created or destroyed in chemical reactions
Fifth Postulate
compounds are formed when atoms of more than one element combine
Sixth Postulate
in a compound, the relative numbers and kinds of atoms are constant
First Law of Atomic Theory
law of definite proportions
Second Law of the Atomic Theory
law of conservation of matter
Third Law of the Atomic Theory
law of multiple porportions
Fourth Law of Atomic Theory
avogadro’s number
Weight of one 12C atom
12 atomic mass units (amu)
-1 unit charge, weighs 5.486×10-4 amu
+1 charge, weighs 1.0073 amu
zero charge, weighs 1.0087 amu
composed of protons, neutrons, and electrons, consists of a dense core (protons and neutrons) surrounded by an electron cloud
What defines and element?
the number of protons in the nucleus
Neutral Atom
equal number of protons and neutrons
Mass Number
total number of protons and neutrons
Isotope Superscript
mass number
Isotope Subscript
atomic number
positive ion, has fewer electrons than protons
negative ion, has more electrons than protons
atomic size, 1 A = 1×10-10 m
Group 1A on the Periodic Table
alkali metals
Group 2A on the Periodic Table
alkaline earth metals
Group 5A on the Periodic Table
Group 6A on the Periodic Table
Group 7A on the Periodic Table
Group 7A on the Periodic Table
Group 8A on the Periodic Table
noble gases
Active Metals
groups 1A and 2A
Main Group Elements
groups 1A trhough 8A
Transition Metals
group of 10 columns
Inner Transition Metals
group of 14 columns
first series-lanthanides
second series-actinides
B, Si, Ge, As, Sb, Te, and At
Ionic Compounds Consist of
1) metal + non-metal
2) metal + polyatomic anion
3) polyatomic cation + non-metal
4) polyatomic cation + polyatomic anion
Molecular Compound
atoms held together by covalent bonds, consist of two or more non-metals bonded together, can also result in polyatomic ions
Naming Binary Molecular Compounds
often formed between non-metals, more positive element first then the more negative ending in -ide, use prefixes mono-, di-, tri-…
Naming Ionic Compounds
Cation first then anion
Naming Ionic Compound Cations
can only form one cation-name of element
can form more than one cation-romen numerals in parentheses to indicate charge
-ous for cation with lower charge, ic for cation with higher charge
Naming Ionic Compound Anions
name of monotomic anion-element with suffix -ide
Naming Polyatmic Anions
-ite, lower amount of oxygen
-ate, higher amount of oxygen
hypo–ite, fewer oxygen than -ite form
per–ate, more oxygens than -ate form
Naming Acids
Derived from the name of the anion
replace -ide with -ic and prefix hydro-
replace -ate with -ic
replace -ite with -ous
Atomic Weight
the weight and average of the masses of the naturally occurring isotopes of an element
Formula Weight
the sum of the atomic weights for all of the atoms that appear in the formula of a compound
Molecular Weight
the sum of the atomic weights for all the atoms in a molecule
the number of particles in an amount of matter equal to the number of atoms in exactly 12 grams of the Carbon-12 isotope, 1 mole = 6.022×1023 particles
Molar Mass
the mass of one mole of a substance in grams, equal to the formula wieght
Avogadro’s Number
the number of 12C atoms in exactly 12 g of 12C, equals 6.022×1023 molecules
a starting substance in a chemical reaction
a substance produced in a chemical reaction
Limiting Reagent
the substance which is completely consumed in a chemical reaction and thereby limits the amount of product formed
Actual Yield
the amount of product actually obtained in the laboratory
Theoretical Yield
Quantity of product that is calculated to be produced when all of the limiting reagent has reacted
Percent Yield
actual yield/theoretical yield x 100
Percent Composition of an Element in a Formula
(# of atoms in a formula)(atomic weight)/formula weight x 100
Percent Composition of a Unit in a Formula
formula weight of the unit/total formula weight x 100
Formula Units per Molecule
molecular weight/formula weight
Combination Reaction
A + B –> AB
Decomposition Reaction
AB –> A + B
Single Displacement Reaction
A + BX –> AX + B
Double Displacement Reactions
AX + BY –> AY + BX
reactant + O2 –> products
the dissolving medium of a solution
a substance dissolved in a solvent to form a solution
a mixture of substances that has a uniform composition
the quantity of solute present in a given quantity of solvent or solution
the concentration of a solution expressed as mole of solute per liter of solution, M
substances which form ions in aqueous solution (water)
Strong Electrolyte
completely forms ions, falls apart in solution
substance that does not form ions in an aqueous solution
Weak Electrolyte
partially dissociates in aqueous solution
Arrhenius Acid
a substance capable of increasing the H+ ion concentration in an aqueous solution
Arrhenius Base
a substance capable of increasing the OH ion concentration in an aqueous solution
Neutralization Reaction
a chemical reaction in which excess OH ions from a base combined with H+ ions from an acid to produce water
the maximum amount of a substance that will dissolve at a given temperature
Molecular Equation
a chemical equation showing complete neutral compound formulas of reactants and products
Ionic Equation
a chemical reaction in which all soluble strong electrolytes are expressed as ions in aqueous solution
Spectator Ion
ions which appear in exactly the same form with the same stochiometric coefficient
Net Ionic Equation
ionic equation from which the spectator ions have been removed
an ionic compound formed by replacing one or more hydrogens of an acid by other cations
solid formed as a product from an aqueous solution
Driving Force
a chemical change which promotes a chemical reaction
Loss of one or more electrons
gain of one or more electrons
Activity Series
a ranking of metals in order of their ease of oxidation
number of moles of solute/volume of the solution in L, moles/L, intensive property
M1 V1 = M2 V2
Identifying Strong Electrolytes
most salts, most common bases, some acids (HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4)
Identifying Weak Electrolytes
most acids and ammonia (base)
Solubility Rule 1
all nitrates, acetates, chlorates, ammonium and compounds of the group 1A elements are soluble
Solubility Rule 2
all chlorides, bromides, and iodides are soluble except for the silver, lead, and mercurous compounds
Solubility Rule 3
all sulfates are soluble except the compounds of silver, mercurous, lead, and heavy alkaline earths (Ca, Ba, Sr)
Solubility Rule 4
all carbonates, sulfites, and phosphates are insoluble except the compounds formed with the group 1A elements and ammonium
Solubility Rule 5
all hydroxides are insoluble except compounds of group 1A elements, ammonium, and the heavy alkaline earths (Ca, Ba, Sr)
Solubility Rule 6
all sulfides are insoluble except groups 1A, 2A, and ammonium
Name the Driving Forces
1) formation of a precipitate
2) formation of a lowly ionized species (water, weak acids and bases)
3) formation of a gas
4) redox process
5) formation or breakage of one or more covalent bonds
Ease of Oxidation
metals on the left side of the periodic table are more easily oxidized than those to the right
Reducing Agent
reductant, the substance which causes another species to be reduced and is thereby itself oxidized
Oxidizing Agent
oxidant, the substance which causes another species to be oxidized and is thereby itself reduced
a chemical equation that shows either the reduction or the oxidation process separately
Redox Process
a chemical reaction which occurs by a TRANSFER of electrons from one species to another
Electronic Structure
the arrangement of electrons in an atom
maximum push or pull
the time required for one cycle of a wave to pass a specific point in space, in seconds
Electromagnetic Radiation
the periodic displacement of the electromagnetic field
the number of cycles which pass a specific point in space in one second, 1/s
the distance between corresponding point on a wave, in meters
Line Spectrum
a spectrum containing radiation of only specific wavelengths
a separation or categorization of light according to its frequency or wavelength
a quantum (discrete bundle) of radiant energy (light)
Continuous Spectrum
a spectrum containing radiation distributed over all wavelengths in the range of interest
Ground State
the lowest energy state the electron can be in
Excited State
an energy state of an electron other than the ground state
Electron Density
the probability of finding an electron at a certain point in space
Quantization of Energy
the specific allowed values of the energy
a region in space with a high (90%) electron density
Degenerate Orbitals
a situation in which two ore more orbitals have the same energy
Quantum Numbers
values which describe the probable location and energy of an electron in an atom
point of no deflection on a wave
Sir Isaac Newton
1687, laws of classical physics
Max Planck
1900, quantization of energy
Albert Einstein
1905, explained photoelectric effect
Neils Bohr
1914, model of hydrogen atom
Louis DeBroglie
1923, duality principle
Wolfgang Pauli
1924, exclusion principle
Werner Heisenberg
1926, uncertainty principle
Planck’s Quantum Theory
on an atomic scale, matter gains or loses energy in discrete amounts called quanta
Planck’s Constant (h) = ?
6.626×10-34 J sec
1 J = ?
1 kg m2/sec2
Photoelectric Effect
light of a sufficiently high frequency can stimulate emission of electrons from the surface of a metal
Delta E = ?
RH [(1/ni2) – (1/nf2)
RH = ?
2.18×10-18 J
Delta E for Emission of Energy
Delta E for Absorption of Energy
Momentum (p) = ?
mass x volume
Wavelength of a Particle
planck’s constant/mass x volume
Duality Principle
both matter and electromagnetic radiation can exhibit wave-like and particle-like behavior
Heisenberg Uncertainty Principle
it is impossible to determine the exact location and momentum of a particle simultaneously
Principle Quantum Number
n = 1,2,3,… (integer values)
defines the shell in which the electron is located
Azimuthal (Orbital Angular Momentum) Quantum Number
l = 0, 1, 2, 3,…, n-1 (integer values)
defines shape of the orbital, each n has its own set of l values
l = 0
s orbital
l = 1
p orbital
l = 2
d orbital
l = 3
f orbital
Magnetic Quantum Number
ml = -l, (-l+1), …, 0, …, (l-1), ldefines the orientation of the orbital in space, each l has its own set of ml values
Electron Spin Quantum Number
ms = -1/2, +1/2defines the direction of spin of the electron
Valence Electrons
the outermost electrons of an atom, higher in energy than the core electrons, used in bonding
Core Electrons
the electrons that are not in the outermost shell of an atom
Electron Configuration
listing of the populates subshells in an atom (in order of lowest to highest energy)
Effective Nuclear Charge
the net positive charge experienced by an electron in a many-electron atom
Ionization Energy
the energy required to remove an electron from a gaseous atom when the atom is in its ground state
Electron Affinity
the energy change that occurs when and electron is added to a gaseous atom or ion
the measure of the ability of an atom participating in a bond to pull electron density toward itself
Metallic Character
an extent to which an element exhibits the physical and chemical properties of a metal
Isoelectronic Series
a series of atoms, ions, or molecules having the same number of electron

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