Modules 1-7

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