Chemistry 112 – Flashcards

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Term
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Definition
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Chemistry
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The study of matter
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Matter
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Anything that has mass and volume
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Mass
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A measure of the gravitational force acting on an object
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Weight
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A measure of the mass of an object on Earth
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Physical Properties
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Properties that can be observed or measured without changing the composition of the matter
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Chemical Properties
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Properties observed when one attempts to change matter into some other type of matter
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Chemical Composition
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What chemicals are present
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Chemical Reactivity
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The ability to interact with other chemicals
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Physical Change
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A change in the state of matter (does not alter the chemical makeup)
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Chemical Change
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Substance changes into something new
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States of Matter
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Solid, Liquid, Gas
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Change of State
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Conversion of matter from one state to another
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Pure Substance
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Matter that is uniform in its chemical composition and properties
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Mixture
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A blend of two or more pure substances in any ratio, each retaining their identity
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Element
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Pure substance that cannot be broken down chemically into simpler substances
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Chemical Compounds
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Two or more elements combined chemically in specific ratios to form a pure substance
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Atom
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A single molecule of an element
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Molecule
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Two or more elements combined chemically in specific ratios to form a pure substance
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Chemical Formula
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A notation for a chemical compound using symbols and subscripts to show how many atoms of each element are present
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Physical quantities
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Measurements of physical properties such as height, volume and temperature requiring both a number and a unit
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SI Units
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Scientific standard set of units closely related to metric units
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Atomic Theory 1
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All matter is composed of atoms
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Atomic Theory 2
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The atoms of a given element differ from the atoms of all other elements
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Atomic Theory 3
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Chemical compounds consist of atoms combined in specific ratios
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Atomic Theory 4
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Chemical reactions change only the way the atoms are combined in compounds; the atoms themselves are unchanged
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Proton
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Subatomic particle with a positive charge
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Electron
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Subatomic particle with a negative charge
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Neutron
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Subatomic particle with no charge
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Nuclear strong force
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The force holding protons and neutrons together in an atom’s nucleus
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Mesons
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Exchanged between protons and neutrons creating nuclear strong force
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Atomic number
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(Z) – The number of protons in each atom of an element
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Mass Number
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(A)– The total number of protons and neutrons in a atom
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Isotopes
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Atoms with identical atomic numbers (Z) but different mass numbers (A); same number of protons, but varying number of neutrons
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Atomic Weight
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Average mass of an element and all of its naturally occurring isotopes
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Periodic Law
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When all the elements are placed in order of increasing atomic number, elements with similar chemical properties will occur at regular (periodic) intervals.
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Periods
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Seven horizontal rows of the periodic table
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Groups
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18 vertical columns of the periodic table
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Main Groups
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The two groups on the far left (1&2) and the six on the far right (13-18)
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Transition Metal Groups
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Elements in groups 3-12
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Inner Transition Metal Groups
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The 14 unnumbered groups shown at the bottom of the table
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Alkali metals
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Group 1 elements
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Alkaline earth metals
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Group 2 elements
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Halogens
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Group 7 elements
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Nobel Gases
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Group 8 elements
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Valence Electrons
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Electrons in the outermost shell of an atom
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Distinguishing Electron
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The last electron added to an element
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Electrically neutral
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All atoms of elements in the periodic table are this because they contain equal numbers of protons and electrons.
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Cation
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A positively charged ion due to the loss of one or more electrons
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Anion
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A negatively charged ion due to the gain of one or more electrons
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Ion
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An atom that has gained or lost one or more electrons and now has a charge
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Ionization energy
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The energy required to remove one electron from a single atom in the gaseous state
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Electron affinity
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The energy released on adding an electron to a single atom in the gaseous state
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First ionization energy
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The amount of energy required to remove the first electron from an atom
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Second ionization energy
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The amount of energy required to remove the second electron from an atom
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Ionic bond
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The electrical attractions between ions of opposite charge in a crystal
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Ionic compound
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A compound that contains ionic bonds
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Octet Rule
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After bonding, each atom will have 8 electrons in its valence shell
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Crystal lattice
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The most stable form of an ionic compound, a crystal of many ions in a rigid, three-dimensional arrangement
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Ionic solids
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Individual cation/anion bonds cannot be determined, therefore, collectively called this
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Polyatomic ions
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Ions that are composed of more than one atom
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Molecular weight
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The sum of the atomic weights for all the atoms in the molecule; the average mass of a substance’s molecules
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Formula weight
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The sum of the atomic weights for all the ions in the compound
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Covalent bond
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The bond formed when atoms share electrons
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Molecule
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A group of atoms held together by covalent bonds
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Repulsive interaction
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Like-charged atomic particles repel each other – nuclei and electrons
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Attractive interaction
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Oppositely charged atomic particles attract each other nucleus/electrons
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Single bond
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A covalent bond formed by sharing one electron pair
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Double bond
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A covalent bond formed by sharing two electron pairs
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Triple bond
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A covalent bond formed by sharing threre electron pairs
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Diatomic
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Two-atom molecules (H2, Cl2, N2, O2, F2, Br2, I2)
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Coordinate Covalent Bonds
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The covalent bond that forms when both electrons are donated by the same atom. This creates a charged molecule
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Molecular formula
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A formula that shows the numbers and kinds of atoms in one molecule of a molecular compound H2O
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Formula Unit
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The formula that identifies the smallest neutral unit in an ionic compound NaCl
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Structural formula
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A molecular representation that shows the connections among atoms by using lines to represent covalent bonds. H-O-H
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Polar covalent bonds
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In molecules of different elements, electrons are attracted more strongly by one atom that by the other and thus are shared unequally.
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Electronegativity
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The ability of an atom to attract electrons in a covalent bond
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Electronegativity difference <0.5
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Covalent bonds
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Electronegativity difference >0.5 & <2.0
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Increasingly polar covalent bonds
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Electronetativity difference > 2.0
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Increasingly ionic bonds
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Intermolecular Forces
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Weak bonds that form between molecules (other than ionic and covalent)
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Dipole-dipole
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An intermolecular force where positive and negative ends of polar molecules are attracted to each other. This results in higher boiling points
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London dispersion forces
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Averaged over time, electron dispersion is uniform. A snap-shot in time may reveal more polarity of electrons and thus a momentary polarity to the molecule
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Reactant
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A substance that undergoes change in a chemical reaction and is written on the left side of the reaction arrow in a chemical equation
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Product
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A substance that is formed in a chemical reaction and is written on the right side of the reaction arrow in a chemical equation
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Law of Conservation of Mass
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Matter is neither created nor destroyed in chemical reactions.
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Combination/Addition reaction
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Molecules A and B combine/react to make C (A+B->C)
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Decomposition Reaction
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Molecule A breaks down into molecules B & C (A->B+C)
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Single Replacement
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One molecule replaces another (A+BC->AC+B)
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Double Replacement
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Both molecules break down and form new molecules (AB+CD->AC+BD)
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Mole
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The amount whose mass in grams is numerically equal to its molecular or formula weight
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Avogadro’s Number
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The number of molecules or formula units in a mole Na = 6.022 x 1023
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Potential energy
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Stored energy ie a coiled spring
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Kinetic energy
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Energy in motion ie hands of the clock moving
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Bond dissociation energy
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The amount of energy that must be supplied to break a bond and separate the atoms in an isolated gaseous molecule
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Law of Conservation of energy
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Energy can be neither created nor destroyed in any physical or chemical change
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Heat of reaction
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The difference between the energy absorbed in breaking bonds and that released in forming bonds; represented by ?H
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Enthalpy change
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Heat of reaction
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Endothermic Process
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A chemical change (like bond breaking) that absorbs heat and has a positive ?H
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Exothermic Process
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A chemical change (like bond formation) that releases heat and has a negative ?H
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Spontaneous process
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A process that, once started, proceeds without any external influence
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Entropy
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A measure of the disorder of a system; ?S
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Free energy change
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Used to describe spontaneity of a process; ?G
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Exergonic
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A spontaneous reaction or process that releases free energy and has a negative ?G
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Endergonic
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A non-spontaneous reaction or process that absorbs free energy and has a positive ?G
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Activation energy
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The amount of energy/heat needed to start a reaction; Ea
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Catalyst
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A substance that accelerates a chemical reaction but is itself unchanged in the process
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Reversible reactions
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A reaction which easily goes in either direction; indicated by a double arrow in equations.
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Chemical equilibrium
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A state in which the rates of forward and reverse reactions are the same
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Le Chatelier’s Principle
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When a stress is applied to a system at equilibrium, the equilibrium shifts to relieve the stress. The stress can be any change in concentration, pressure, volume, or temperature that disturbs original equilibrium
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Kinetic-Molecular theory of gases #1
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A gas consists of many particles, either atoms or molecules, moving about at random with no attractive forces between them
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Kinetic-Molecular theory of gases #2
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The amount of space occupied by the gas particles themselves is much smaller than the amount of space between particles.
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Kinetic-Molecular theory of gases #3
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The average kinetic energy of gas particles is proportional to the Kelvin temperature
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Kinetic-Molecular theory of gases #4
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Collisions of gas particles, either with other particles or with the wall of their container, are elastic; that is, the total kinetic energy of the particles is constant
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Ideal Gas
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A gas that obeys all the assumptions of the kinetic-molecular theory
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Boyle’s law
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The volume of a gas is inversely proportional to its pressure for a fixed amount of gas at a constant temperature – More pressure, less volume, same temp.
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Charles’s law
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The volume of a gas is directly proportional to its Kelvin temperature for a fixed amount of gas at a constant pressure. – More heat, more volume, same pressure
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Gay-Lussac’s law
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The pressure of a gas is directly proportional to its Kelvin temperature for a fixed amount of gas at a constant volume. – More heat, more pressure, same volume
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Avogadro’s law
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The volume of a gas is directly proportional to its molar amount at a constant pressure and temperature
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Dalton’s law
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The total pressure exerted by a gas mixture is the sum of the partial pressures of the components in the mixture.
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Vapor
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A molecule near the surface of a liquid can break free of the liquid and escape into this gaseous state
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Vapor pressure
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The contribution that the gas molecules make to the total pressure of the gas above the liquid according to Dalton’s law
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Boiling
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Bubbles of vapor form under the surface and force their way to the top of the liquid
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Surface tension
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The resistance of a liquid to spread out and increase its surface area; caused by the difference between the forces experienced by molecules at the surface and these experienced by molecules in the interior.
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Specific heat
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The capacity to absorb a large quantity of heat while changing only slightly in temperature
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Heat of vaporization
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The ability of water to carry away a large amount of heat with it evaporates
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Bronsted-Lowry acid
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Any substance that is able to give a hydrogen ion to another molecule or ion, and need not occur in water
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Bronsted -Lowry base
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a substance that accepts a hydrogen ion from an acid, and need not occur in water
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Amphoteric
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Substances like water, which can react as either an acid or a base depending on the circumstances.
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Dissociation
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The splitting apart of an acid in water to give H+ and an anion
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pH
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A measure of the acid strength of a solution; the negative common logarithm of the H3O+ concentration
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Buffer
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A chemical reaction that keeps hydrogen ions from getting too high
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Nuclear reaction
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A reaction that changes an atomic nucleus, usually causing the change of one element into another
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Nuclear decay
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The spontaneous emission of a particle from an unstable nucleus
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Transmutation
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The resulting change of one element into another
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Alpha emission
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The emission of 2 protons and 2 neutrons as an ? particle from an unstable radioactive nucleus, resulting in a positive charge
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Beta emission
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The result of a neutron decomposing into a proton and an electron, retaining the proton in the nucleus and emitting the electron as a ? particle, resulting in a negative charge
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Gamma emission
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The emission of photons which have no charge
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Positron emission
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The conversion of a proton in the nucleus into a neutron plus an ejected positron, a ‘positive electron,” which has the same mass as an electron but a positive charge
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Electron Capture
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A process in which the nucleus captures an inner-shell electron from the surrounding electron cloud, thereby converting a proton into a neutron.
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Half-life
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The amount of time required for one-half of the radioactive sample to decay
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Nuclear fission
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The fragmenting of heavy nuclei
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Nuclear fusion
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The joining together of light nuclei
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Chain reaction
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A reaction that is self-sustaining
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Critical Mass
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The minimum amount of radioactive material needed to sustain a nuclear chain reaction
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