Modules 1-7

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Chemistry
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The study of the composition of matter and the changes it undergoes.
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Organic Chemistry
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Study of carbon compounds.
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Inorganic Chemistry
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Study of compounds that do not contain carbon.
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Biochemistry
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Study of processes that take place in organisms.
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Analytical Chemistry
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Focuses on composition of matter.
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Physical Chemistry
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Study of mechanism, rate, and energy changes in matter when undergoing change.
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Pure Chemistry
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Pursuit of chemical knowledge for its own sake.
Ex: Study of metalloids and their properties.
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Applied Chemistry
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Chemistry in use.
Ex: Use of metalloids in semiconductors for computer chips.
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Qualitative Data
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Descriptions.
Ex: The solution is blue, the base feels slippery, or there are bubbles formed in the reaction.
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Quantitative Data
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Measurements.
Ex: The test tube has a mass of 3.2 grams or the length of my shoe is 20 centimeters long.
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Scientific Method
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1. State problem and collect data.
2. Formulate hypothesis.
3. Perform experiments.
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Independent Variable
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The manipulated variable.
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Dependent Variable
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The measured variable.
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Control
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Remains unchanged in the experiment.
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Observations
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Witnessed and can be recorded.
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Theories
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Interpretations or explanations.
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Law
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Summary of observed behavior.
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Matter
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Anything that has mass and takes up space.
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Three states of matter
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1. Solid
2. Liquid
3. Gas
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Solid
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Particles packed tightly, usually dense, incompressable, and do not flow. Definite volume and definite shape.
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Liquid
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Have a definite volume and the ability to flow. Do not have a definite shape. Takes shape of container it fills.
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Gas
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Particles move rapidly and independently, fill the container they are in, and have the ability to flow. No definite shape or volume.
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Fixed Composition a.k.a. Pure Substance
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Always has the same composition and internal properties. Can write a formula for it.
Ex: H2O, NaCl, C12H22O11, and N2
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Variable Composition a.k.a. Mixture
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Cannot write a formula for it.
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Atoms
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Smallest unit of matter.
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Elements
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Same kind of atoms.
Ex: Carbon, oxygen, and hydrogen
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Compounds
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Chemical combination of elements, forms molecules or crystals.
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Homogeneous
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The same throughout the sample. All elements and compounds. If a mixture, it’s called a solution.
Ex: Saltwater, sugar water, and air.
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Heterogeneous
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Composition can vary within the sample. All mixtures that aren’t solutions are heterogeneous.
Ex: Salad, pizza, soil.
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Physical Properties
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Can be measured.
Ex: Size, shape, color.
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Physical Changes
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Can be observed without changing the substance.
Ex: Melting, freezing, boiling.
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Chemical Properties
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A description or characteristic of the substance that when observed causes the substance to be changed.
Ex: Flammable, corrosive, and reactivity.
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Chemical Changes
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Composition of matter changes, a chemical reaction has occurred.
Ex: Decompose, rust, and burn.
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Other states of matter
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Plasma and Bose Einstein Condensate (BEC)
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Plasma
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Similar to gas, but some particles are ionized. Found in stars and neon lights.
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Bose Einstein Condensates (BEC)
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At very low temperatures, atoms become blobs.
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Kinetic Molecular Theory (KMT)
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Based on five assumptions:
1. Gases are made up of many tiny particles that are far apart.
2. The movement of the gas produce elastic explosions.
3. Gas particles possess kinetic energy because they are in constant motion.
4. There are no attraction or repulsion forces between particles.
5. The average kinetic energy is dependent upon the temperature.
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Six Properties of Gas
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1. Expansion-Gases will fill their container.
2. Fluidity-Gases flow.
3. Low Density-Gases have low density.
4. Compressibility-Gases can be compressed into a smaller container.
5. Diffusion-Gases will completely mix in a container.
6. Effusion-Gases can be forced through a tiny opening.
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Heat
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Total kinetic energy. Measured in joules(J).
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Temperature
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Average kinetic energy. Measured in C or K. K=C+273
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Direction of Heat
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Heat travels from warm to cold until equilibrium is met.
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Heat, temperature, and gases
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As heat is added, temperature increases and the particles in the gas expand and move faster. Thus, the gas expands.
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What causes a phase change?
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Change in temperature and pressure.
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Liquid to Gas
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Vaporization-conversion of liquid to gas
Evaporation-conversion of liquid to gas at the surface and is not boiling, occurs in an open container
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Other Changes of State
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Sublimation-solid directly to gas
Deposition-gas directly to solid
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Triple Point
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Temperature and pressure at which all three phases can exist in equilibrium.
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Critical Point
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The temperature and pressure at which the liquid and gas phase become indistinguishable.
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Proton
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Mass: 1 amu
Charge: Positive
Symbol: p+
Location: Nucleus
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Electron
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Mass: 0 amu
Charge: Negative
Symbol: e-
Location: Energy Levels
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Neutron
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Mass: 1 amu
Charge: Neutral
Symbol: n0
Location: Nucleus
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Average Atomic Mass
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[(Atomic Mass)(%Abundance)+(Atomic Mass)(%Abundance)]/100
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Facts about protons, electrons, and neutrons.
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-Number of protons=atomic number
-Number of electrons=number of protons
-Number of neutrons=atomic mass- number of protons
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Isotopes
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Same atomic number, but different atomic mass. Thus, there are a different number of neutrons than usual.
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Electromagnetic Spectrum (EM Spectrum)
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Longest wavelength is the radio waves, then it is the microwaves, after that is infrared, next is visible light, then ultraviolet, after that is X rays, and the shortest wavelength is Gamma rays.
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Wave-Particle Nature of Light
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Light travels as waves, but carries packets of energy called photons.
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Wavelength
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Distance from the top of one wave to another.
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Frequency
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How often the wave rises and falls at a specific point.
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Emission Spectrum
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Spectrum of light released from excited atoms of an element.
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Endothermic
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Absorbs/requires input of energy. Energy “enters” the system.
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Exothermic
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Releases/emits/produces energy. Energy “exits” the system.
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Quantum Model
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Probability of locating an electron at any place.
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Heisenberg Uncertainty Principle
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It is impossible to know both the velocity and position of an electron at the same time.
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Principle Energy Level
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Indicates main energy level occupied by e-. Always a whole number.
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Sublevels
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Indicates shape or type of orbital.
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Orbitals
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Each orbital holds a pair of electrons. Both electrons will spin in opposite directions.
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Electron Configurations
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Shows the electron arrangement in an atom and always represents the lowest possible energies.
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Aufbau Principle
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electrons fill orbitals that have the lowest energies first.
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Electron Configuration of Pb
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1s2 2s2 2p6 3s2 3p6
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Shorthand Notation
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Uses noble gases as a reference point.
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Orbital Notation
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Uses lines to represent orbits and arrows to represent the spin of each electron.
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Hund’s Rule
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e- spread out within equivalent orbits.
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Radioactivity
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Radiation is emitted during radioactive decay.
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Three Types of Radiation
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1. Alpha Radiation
2. Beta Radiation
3. Gamma Radiation
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Alpha Particles
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Helium nuclei emitted from a radioactive source. Consist of two protons and two neutrons which gives it a +2 charge.
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Beta Particles
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Same properties as electrons.
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Gamma Ray
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High energy photons emitted by radioactive isotopes. Most powerful form of radiation.
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Half Life
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Time it takes for half of a radioactive isotope to decay.
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Nuclear Fission
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Splitting of a nucleus into smaller parts. Initiated by hitting fissionable isotopes with neutrons. Only two fissionable isotopes are Uranium-235 and Plutonium-239.
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Nuclear Fusion
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When nuclei combine to form a nucleus of greater mass.
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Nuclear Waste
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Mainly, spent fuel rods from nuclear power plants. Made up of Uranium-235 and Plutonium-239.
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Period
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Rows of periodic table. Currently, seven exist.
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Group
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Columns of periodic table. Also called families. Have similar properties. Currently, there are eighteen.
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Alkali Metals
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-Group 1
-Include all except H
-Soft metals
-Very reactive, not found in nature as free elements
-1 valence electron
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Alkaline Earth Metals
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-Group 2
-Stronger, denser, and harder than Alkali metals
-Still too reactive to be found as free elements.
-2 valence electrons
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Transition Metals
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-Groups 3-12
-Typical metals
-Less reactive
-Number of valence electrons vary
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P-Block Elements
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-Groups 13-18
-Include metals, metalloids, and nonmetals
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Halogens
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-Group 17
-Most active nonmetals
-React with metals to form salt
-7 valence electrons
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Metalloids
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-Along zigzag line/ staircase
-Both metal and nonmetal properties
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Lanthanides
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-Period 6
-Shiny metals
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Actinides
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-Period 7
-All are radioactive
-No isotopes are known to be stable
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Valence Electrons
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Outer most electrons that are available to be gained or lost. 8 valence electrons=chemically stable. Elements react to reach eight valence electrons.
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Ions
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An atom or group of bonded atoms with a positive or negative charge.
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Positive Ions
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Called cations. Form when an atom loses electrons.
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Negative Ions
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Called anions. Form when an atom gains electrons.
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Atomic Radius
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Half the distance between the nuclei of identical atoms that are bonded together. Size of the atom. Trend increases down and to the left.
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Ionization Energy
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Energy required to remove one electron from a neutral atom. Trend increases up and to the right.
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Electronegativity
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Measure of the ability of an atom in a chemical compound to attract electrons closer to it. Noble gases aren’t included. Trend increases up and to the right.
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Metallic Activity
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Trend increases down and to the left.
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Ionic Radii
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Positive ions are smaller than an atom of the same element because they lose an electron. Negative ions are larger because they gain an electron.
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Ionic Bonds
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-Transfer of electrons
-Metal and nonmetal
-Cation+anion
-Metals lose their valence e- and nonmetals gain a valence e-
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Lewis Dot Diagram
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Visual interpretation of the number of valence electrons in an atom or ion.
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Properties of Ionic Bonding
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-Hard, brittle, crystalline solids at room temp.
-High melting and boiling points
-Don’t conduct as solids
-Do conduct when melted or dissolved in water
-Mot are soluble in water
-Called salts
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Metallic Bonding
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-Cations in a sea of electrons
-Malleable-can be flattened
-Ductile-can be pulled into wire
-Bendable/shapeable
-Alloys-metal mixtures
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Chemical Formula
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Show kinds and numbers of atoms in smallest representative unit.
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Oxidation Numbers
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The charges that an atom takes on to obey the Octet Rule.
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Polyatomic Ion
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Group of covalently bonded ions that have an overall charge.
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Properties of Molecular Compounds
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-Formed when atoms share electrons
-Occurs between nonmetals
-Most have low melting points
-Not very conductive
-Soft
-Dissolve in water to form solution
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Single Bond
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Share one pair of electrons. Group 17
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Double Bond
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Share two pairs of electrons. Group 16
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Triple Bond
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Share three pairs of electrons. Group 15
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Coordinate Covalent Bond
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Shared pair of electrons in which both came from the same atom.
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Bond Dissociation Energy
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Indicates the strength of a bond. The higher the energy, the stronger the bond.
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Mono-
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1
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Di-
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2
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Tri-
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3
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Tetra-
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4
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Penta-
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5
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Hexa-
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6
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Hepta-
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7
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Octa-
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8
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Nona-
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9
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Deca-
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10
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Valence Shell Electron-Pair Repulsion Theory
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Valence electron pairs repel each other.
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Polar Bonds
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Unequal sharing of electrons creates partially positive and partially negative ends. Called a dipole.
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Nonpolar Bonds
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No partially charged regions are created resulting in equal sharing of electrons.
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Dispersion Forces
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Occur between nonpolar forces. Weakest intermolecular forces.
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Dipole-Dipole Forces
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Occur between polar molecules.
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Hydrogen Bonding
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Occur between molecules with an H-F, H-O, H-N bond. Strongest form of intermolecular bonds.
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Mole
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Defined as the number of carbon atoms in exactly twelve grams of carbon-12. Abbreviated mol. 1 mol=6.02×10^23. Used to measure matter. Used because atoms and molecules are too small to measure.
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Avogadro Number
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6.02×10^23
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Percent Composition
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Molar mass represents 100% of the mass one mole of a substance. Mass of each element compared to the total mass of the compound.
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Empirical Formulas
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Lowest whole number ratio of elements. All ionic formulas are empirical.
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Subscripts
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Tells how many atoms of a particular element are in a compound.
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Coefficients
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Quantity of molecules in a compound.
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Synthesis
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Two or more substances combine to form a new compound.
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Decompsition
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A single compound undergoes a reaction that produces two or more simpler substances.
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Single Replacement
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A+BX->AX+B
BX+Y->BY+X
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Double Replacement
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The ions of two compounds exchange places in an aqueous solution to form two new compounds.
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Combustion
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A substance combines with oxygen, releasing a large amount of energy in the form of light and heat.
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Mole-Mole Conversions
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Given the number of moles of a substance to find how many moles of another substance are created from that in a chemical reaction.
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Mole-Mass Conversions
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Given the number of moles of a substance to find how many grams of another substance are created from that in a chemical reaction.
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Mass-Mole Conversions
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Given the mass of a substance to find how many moles of another substance are created from that in a chemical reaction.
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Mass-Mass Conversions
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Given the mass of a substance to find how many grams of another substance are created from that in a chemical reaction.
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Mass-Volume Conversions
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Given the mass of a substance to find how many liters of another substance are created from that in a chemical reaction.
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Mole-Particle Conversions
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Given the number of moles of a substance to find how many particles of another substance are created from that in a chemical reaction.
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Limiting Reactant
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Limits the amount of product that is produced.
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Excess Reactant
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Reactant that produces more product than the limiting reactant. Does not affect how much product is created.
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Theoretical Yield
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The amount of product expected to be recovered according to the limiting reactant.
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Actual Yield
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The amount of product recovered in an actual experiment.

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