Review (more in-depth) Flashcards
Avogadro (S)
Determined the number of objects in one mole is 6.022 x 10^23. He also found that equal volumes of gases at the same temp. and pressure contain the same number of molecules.
Born (S)
Showed that the probability of the location of an electron can be determined.
Boyles (S)
Found that volume and pressure of a gas vary inversely if the amount and temp. is held constant.
Bronsted and Lowry (S)
Defined an acid as a proton receiver and the base as the proton giver.
Chadwick (S)
Discovered neutrons.
Charles (S)
Determined that the temp. and volume of a gas are directly related if amount and pressure is held constant.
Dalton (S)
Modern atomic theory. (Ptotal=P1+P2…Pn)
Dobereiner (S)
First attempt to classify elements by grouping elements with similar properties precursor to the periodic table.
Gay-Lussac (S)
Determined that temp. and pressure of a gas are directly related if amount and volume are held constant.
Graham (S)
Determined that the ratio of the rates of the diffusion of gases is equal to the square root of the inverse ratio of their molecular masses.
Heisenberg (S)
Uncertainty Principle (It is impossible to know the exact position of an electron in an atom)
Hund (S)
Electrons will fill all empty orbitals before pairing up in one orbital.
Le Chatelier (S)
Determined that if stress is applied to a system at equilibrium that the system will shift as to relieve the stress.
Lewis (S)
Used dot diagrams to represent an atom and its (outermost) valence electrons. Also defined an acid as an electron receiver and a base as the electron giver.
Mendeleev (S)
Proposed that the properties of elements were a function of their atomic masses and thus formulated the periodic table.
Millikan (S)
Used oil drop experiment to determine charge on an electron.
Mosely (S)
Revised the periodic table and found that properties of elements are based on atomic number rather than atomic mass.
Pauli (S)
Exclusion Principle (two electrons can never occupy the same space)
Planck (S)
Quantum Theory (energy is given off in packets called quanta)
Rutherford and Bohr (S)
Planetary model of atomic structure; Rutherford’s gold foil experiment showed that the atom is made of mostly space and with dense nucleus.
Schrodinger (S)
Wave Equation (used to determine the position of an electron as a wave rather than a particle)
Thompson (S)
Discovered electrons using cathode ray tube. Determined charge/mass ratio.
pH
-log[H3O+]=-log[H+]
pOH
-log[OH-]
Hydronium
H3O+ or H+, characteristics of acids
Hydroxide
OH-, characteristic of bases
Acetic Acid
HC2H3O2
Nitric Acid
HNO3
Hydrochloric Acid
HCl
Sulfuric Acid
H2SO4
Acid
Below 7 on the pH scale, [H+]>[OH-]
Base
Above 7 on the pH scale, [OH-]>[H+]
Neutral
pH of 7, [H+]=[OH-]
Indicator
Weak acid or base whose conjugate is a different color and changes colors in acids/bases.
Titration
Addition of a solution of known concentration to a solution of unknown concentration to discover the unknowns concentration.
Endpoint
End of titration, visual determination where there are equivalent concentrations of acid and base.
Equivalence pt.
Mathematical determination where there are equivalent concentrations of acid and base.
Standard Solution
A solution with a known concentration.
Concentration
Measurement of amount of solute relative to solvent or solution.
Dilute
A “weak” solution, a solution of low concentration, with large amounts of solvent.
Arrhenius Acids and Bases
donate H+ (acid) or OH- (base)
Bronsted-Lowry Acid
Proton donor.
Bronsted-Lowry Base
Proton receiver.
Conjugate Acid
Formed when a base gains a proton.
Conjugate Base
Formed when an acid donated a proton.
Monoprotic
Acid containing more than 1 ionizable H+
Diprotic
Acid containing 2 H’s
Triprotic
Acid containing 3 H’s
Polyprotic
Acid containing more than 1 ionizable H+
Lewis Acid
Electron acceptor.
Lewis Base
Electron donor.
Ammonia
NH3 (weak base)
Atomic Number
The number of protons an element has that indicated what element it is.
Average Atomic Mass
The average mass number of an element determined by considering the masses and relative proportions of the isotopes of that element.
Mass Number
The number of protons and neutrons in one atom that varies with isotopes as the neutrons change.
Anion
The negative ion formed by a gain of 1 or more electrons.
Cation
A positive ion formed by a loss of 1 or more electrons.
Cathode Ray Tube
Tube used to determine the relative size and charge of an electron and proton “Crooke’s Tube”
Transition Metal
A metal found in the center of the periodic table that can have multiple oxidation numbers so it has a roman numeral with it when naming compounds.
Halogen
An element found in Group 17 with a -1 oxidation number.
Alkali Metals
An element found in Group 1 with +1 oxidation. They are very reactive.
Metal
Found to the left of the stairstep line on the periodic table. Metals are ductile, malleable, lusterous, conductors and have + oxidation numbers.
Nonmetal
Found to the right of the stairstep line and do not have characteristics of metals. Often, yet not always, gases.
Metalloid
Elements having a common side with the stairstep line excluding aluminum that have some characteristics of metals.
Ion
An atom that has an unequal number of electrons and protons thus producing a charge. A cation is the result of too few electrons, while an anion is the result of two many electrons.
Atom
The smallest piece of an element that retains the properties of an element. It is composed of protons, neutrons, and electrons.
Isotope
A form of an atom with a different number of neutrons thus having a different atomic mass than other atoms of this element.
Neutron
A subatomic particle found in the nucleus with a mass number of 1 and no charge.
Proton
A subatomic particle found in the nucleus with a mass number of 1 and a + charge.
Electron
A subatomic particle found outside the nucleus in a cloud with a mass number of 0 and a - charge.
Electron Affinity
The attraction an atom has for an electron.
Electronegativity
The attraction an atom has for a shared pair (bonding) of electrons.
First Ionization Energy
The energy required to remove the most loosely held electron from an atom.
Wave Particle Duality
Waves have particle properties and particles have wave properties.
Octet Rule
Atoms are most stable with eight electrons in the outer energy level.
Periodic Table
Arrangement of elements according to atomic number and electron configuration (first organized by Mendeleev)
Quantum Numbers
Numbers used to locate an electron in an atom.
Principal Quantum Number
Designates the energy level and size of the electron cloud and corresponds to periods on the periodic table.
Orbital Quantum Number
Designates the energy sublevel and shape of the electron cloud (s, p, d, etc.)
Magnetic Quantum Number
Designated the orientation of the electron cloud
Spin Quantum Number
Designates the direction of the electron and allows to distinguish between the two electrons in an orbital.
Atomic Radius
Distance from the nucleus to the outermost electron energy level.
Sublevels
Differing electron orbitals within an energy level discernible by shape. S is a sphere, P has 2 lobes, and D has 4 lobes.
Energy Levels
Specific energy possessed by a group of electrons in an atom corresponding to distance from the nucleus.
Electron Cloud
Space that an electron has a high probability of occupying in an atom.
+ Charge
Charge of a proton (nucleus)
- Charge
Charge of an electron (cloud surrounding nucleus)
Orbitals
Space that can be occupied by electrons with the same energy level, sublevel, and orientation. Only two electrons can occupy an orbital.
Noble Gases
Elements located in Group 18 on the periodic table. They are inert because they possess a full outer energy level of electrons.
Diatomic Molecules
The seven elements (gases) that when they are by themselves in elemental form are found in pairs. When combined with other elements or forming ions they need not be in pairs.
Single Bond
One shared pair of electrons between two atoms.
Double Bond
Two shared pair of electrons between two atoms.
Metallic Bond
Bond between two metal atoms held together by a “sea” of delocalized electrons.
VSEPR
Valence Shell Electron Pair Repulsion theory for molecular shape.
Tetrahedral
Shape of an atom with 4 equally spaces bonds with a 109.5 degree angle.
Bent
Shape of an atom with 2 bonds and unshared pairs of electrons.
Linear
Shape of an atom with 180 degree angles.
Polar
Uneven charge distribution. Has a more negative side and a more positive side. There are stronger intermolecular forces between polar molecules.
Nonpolar
Even distribution of charge.
+ Delta
Partial positive charge.
- Delta
Partial negative charge.
Intermolecular Forces
Attractive forces between molecules.
Dipole-Dipole
Intermolecular force between two polar molecules. The more positive side is attracted to more negative side of other molecule.
Ionic Compound
A compound formed by two ions held together by an attraction between the opposite charged ions. Bond due to transferring electrons causing the formation of ions and found when metals and nonmetals combine.
Molecular Compound
Compound formed by covalent bonds (sharing electrons). Found when two nonmetals combine or in polyatomic ions.
Hydrogen Bond
A dipole-dipole attraction between molecules when hydrogen is part of the compounds. Water leads to hydrogen bonding. These are strong intermolecular attractions, but they are much weaker than covalent and ionic bonds.
Covalent Bond
Bond created when atoms share electrons found between two nonmetals.
Ionic Bond
Bond created when ions are attracted.
Van der Waals
Term for intermolecular forces (generally weak) Ex. London Dispersion forces, Hydrogen bonding, and dipole-dipole attractions.
London Dispersion
An intermolecular force found in nonpolar molecules.
Lewis Structures
Drawing of covalent compound that represents bonds with dashes and unshared electrons with dots.
Lewis Dot Diagrams
Diagrams representing the valence electrons of an element
Energy
The ability to do work. Many different kinds.
Chemical Energy
Energy that molecules contain by being bonded together.
Law of Conservation of Mass/Energy
(E=mc^2) Einsteins law that says energy and matter are equivalent and can be converted from one to the other, while total amount is constant.
E=mc^2
Energy=(mass)(speed of light)^2
Exothermic
A reaction that releases energy to its surrounding environment.
Endothermic
A reaction that absorbs energy from its surrounding environment.
Activation Energy
Energy required to start a reaction.(Decreased by catalysts)
Calorie
Unit for heat. Nutritional value is 1000 Calories.
calorie
Unit for heat.
Joule
(J) SI unit for heat.
Entropy
(S) Measure of disorder
Enthalpy
(H) heat content, △H is negative for exothermic reactions.
Rate Law
Represents the effect of concentration on the speed of a reaction.
STP
Standard temperature and pressure. 1 arm and 273 K.
Boyle’s Law
P1V1=P2V2. In which P and V are inversely related.
Charles’ Law
V1/T1=V2/T2. In which P and T are directly related.
Graham’s Law
R1/R2=sq. root of M2/M1. Rate of effusion/diffusion is inversely proportional.
Dalton’s Law
Pt=P1+P2…
Molar Volume
(At STP) 1 mol/22.4 L for any gas.
Ideal Gas Law
PV=nRT (atm, L, moles, K)
R
Ideal gas constant: .0821 L•atm/mol•K
Atm
Atmospheres unit of pressure at sea level=1 atm
Torr
Pressure unit at sea level=760 Torr
Mm of Hg
Pressure unit at sea level= 760 mm Hg
Kpa
Pressure unit at sea level=101.3 kPa
Matter
A material that takes up space and has inertia.
Mixture
A material consisting of 2 or more substances that can be homogeneous or heterogeneous.
Solution
A homogeneous mixture.
Homogeneous
A substance that is the same throughout.
Heterogeneous
A substance that is not the same throughout
Chemical Property
A property of a substance that deals with how the substance reacts with its surroundings.
Physical Property
A property of a substance that does not deal with how it reacts.
Phase
A distinct section of matter with uniform properties that is different from its surroundings
Precipitate
A solid that forms out of mixing two solutions due to its insoluble nature.
Compounds
A substance composed or two or more atoms held together by chemical bonds.
Elements
A substance composed of one atom, and always have the same number of protons to determine the type of element.
Formula Mass
The total mass of a formula measured in amu. For 1 atm/molecule and g for one mole.
Molecular Formula
Formula for the actual compound that exists.
Empirical Formula
Formula that indicates the ratio of atoms of each element present in a substance, but doesn’t give the actual molecule that exists.
Formula
A group of symbols to indicate the number and kind of atoms in a compound.
Avogadro’s Number
6.022•10^23. Avogadro’s number of something is one mole. 1 mole of atoms is equal to the formula mass in grams of that atom.
Mole
The SI unit for measurement of the quantity of a substance = to Avogadro’s numbers of atoms/molecules in that substance.
Percent Composition
The percentage of each element found in a compound as calculated by comparing masses of the elements
Hydrate
A molecule that has water molecules attached to it and is indicated by a dot in between the number of water molecules and the actual substance.
Equation
A representation of the number of moles, reacting substances, products, and the state that each substance is found.
Chemical Reaction
A reaction where a chemical change has taken place that can be represented with an equation.
Balanced Equation
A chemical equation that indicates the appropriate number of miles of each substance that will react and be formed to not break the law of conservation of mass/energy.
Reactants
The substances that are present before the reaction occurs.
Products
The substances that are present after the reaction occurs.
Single Displacement Reaction
A reaction where an element and an ionic compound trade partners to form a new ionic compound while the other element of the original compound is left alone.
Double Displacement Reaction
A chemical reaction where two compounds trade partners to form two new compounds.
Synthesis Reaction
A chemical reaction where two elements are combined to form a compound.
Decomposition Reaction
A chemical reaction where one compound is broken down into two or more elements.
Combustion Reaction
A chemical reaction where a hydrocarbon is combined with oxygen to form carbon dioxide and water.
Coefficients
A number before a substance in an equation to indicate the number of moles of that substance that will react or be formed.
Subscripts
Small number lowered after an atom within a compound to indicate how many atoms of that element are found in a compound. This cannot be changed to balance equations, because it determines what kind of substance is present.
Limiting Reagent
The reactant that is completely used up during a reaction if it goes to completion.
Enzyme
Catalysts found in the body.
Catalysts
Chemicals that speed up a reaction, without being permanently changed in the reaction.
Reactants
Substances on the left of a chemical equation. The starting materials.
Products
Substances on the right of the chemical equation. The ending materials.
Reversible Reaction
A reaction that can proceed in both directions (with double arrow)
Equilibrium
A dynamic condition, where opposing reactions occur at equal rates.
Reaction Rate
The speed of a chemical reaction, dependent upon temperature, nature of the reactants, concentration and surface area.
Precipitates
A solid formed by the mixing of two aqueous solutions.
Net Ionic Equation
Reaction where all soluble compounds are broken down into ions and only actual participants are written.
Spectator Ions
Ions that don’t participate in the chemical reaction.
Binary Compound
An compound consisting of two elements
Polyatomic Ion
A group of atoms covalently bonded that have a + or - charge which will lead them to forming ionic bonds.
Anode
The positive end of a cathode tube that attracts the anions.
Cathode
The negative end of a cathode ray tube that attracts the cations.
Oxidation Numbers
A number that is assigned to an atom to indicate whether that atom wants to gain or lose electrons. A + means it wants to lose, a - means it wants to gain.
Oxidizing Agent
Causes oxidation
Reducing Agent
Causes reduction.
Redox Reaction
Reaction where electrons are transferred.
Reduction
Gaining electrons.
Oxidation
Losing electrons.
Solute
The dissolved material (often solid)
Solvent
The dissolving material (often water)
Molarity
A unit for concentration that is moles of solute per liter of solution.
Molality
A unit for concentration: moles of solvent per kilogram of solvent
Freezing Point Depression
△tf = Kfm, freezing point is lowered by the addition of a solute
Boiling Point Elevation
△tb=Kbm, boiling point is raised by the addition of a solute.
Vapor Pressure Lowering
(Raoult’s law): Psoln = XsolvPsolv vapor pressure is lowered by the addition of a solute.
Miscible
Two substances that will dissolve in each other.
Alloy
A solution made of two metals.
Colligative Properties
Properties that depend on the amount of solute in the solution.
Solubility
Measure of the amount of solute that can be displaced in a specific amount of solvent.
Stoichiometry
A type of problem where we use mass and volume values to determine the amount of each reactant that will react and the amount of products that will form. It utilizes the mole/mole ratio determined by a balanced equation.
Stoichiometry
Relationship between the amount of reactants and products in the chemical reaction.
Mole to Mole Ratio
Conversion from substance A to B found by viewing coefficients
Quantitative
Data concerning numerical measurement of a quantity
Qualitative
Data concerning the general measurement of a quantity.
Mass
The amount of matter.
SI
Standard international units for metrics.
Kelvin
The standard international unit for temperature.
Accuracy
How correct a measurement is
Precision
How consistent the same measurement is gotten.
Density
Mass divided by volume.
Buret
A calibrated glass tube that allows for careful release of liquids needed for titration.
Graduated Cylinder
Calibrated container for measuring volume of a liquid.
Erlenmeyer Flask
Flask with narrow neck and triangular body.
Florence Flask
Flask with narrow neck and round body.
Percentage Yield
The percentage of actual product vs. expected product.
Volume
Unit derived from length cubed.
