Honors Chem Final.docx

  1. Matter and Change
    1. Branches of Chemistry
      1. Organic - carbon
      2. Inorganic - non organic
      3. Physical - energy
      4. Analytical - identify components and composition of matter
      5. Biochemistry - living
      6. Theoretical - Math and computers
    2. Atoms, elements, compounds
      1. Chemical - substance that has definite composition
      2. Matter - anything that has mass and takes up space
      3. Mass - amount of matter
      4. Volume - amount of 3D space
      5. Atom - smallest unit of an element
      6. Element - pure substance that can not be broken down
      7. Compound - substance that can be broken down
    3. Physical and chemical properties
      1. Physical property - Characteristic that can be seen or measured without changing the identity of the substance
        1. Intensive - doesn't depend on amount of matter
          1. MD, BP, density, conductivity
        2. Extensive - depends on amount of matter
          1. Volume, Mass, Amount of energy
      2. Chemical property - Ability for a substance to change into a different substance
    4. Physical and Chemical Changes
      1. Physical Change - Change that doesn’t change the identify of substance
        1. Melting, boiling, cutting, grinding
        2. Used to separate mixtures
      2. Chemical change - Change from one substance to another
        1. Burning, digesting, rusting, corrosion
        2. Used to separate compounds
        3. Evidence of Chem change
          1. Heat, light, precipitate, new products, gas
    5. Heterogeneous Mixtures
      1. Different throughout
    6. Homogeneous mixtures
      1. Same throughout
    7. Pure Substance
      1. Matter either an element or single compound
        1. O2 or H2O
    8. States of matter
      1. Solid, Liquid, Gas, Plasma
        1. Plasma - high temperature where atoms lose electrons
    9. Separation Methods/Techniques
      1. Filtration
      2. Crystallization/evaporation
      3. Distillation
      4. Chromatography
    10. Properties of Metals, Non metals, and Metalloids
      1. Metals
        1. Left side of PT
        2. Malleable - ability to become a sheet
        3. Ductile - drawn into thin wire
        4. Tensile strength - how hard to pull apart
        5. Luster - shiny
        6. Conductivity - electrical and thermal
      2. Non-Metals
        1. Right side of PT
        2. Poor conductors
        3. Brittle
      3. Metalloids
        1. On staircase
          1. Except Al
        2. Properties of metals and nonmetals
        3. Semiconductors
  2. Dimensional Analysis
    1. Scientific Method
      1. Observing
        1. Qualitative/Quantitative
      2. Hypothesis
        1. Must be testable
      3. Testing
        1. Experiment
      4. Formulating theories
        1. Theories are broad generalizations
      5. Publish
    2. SI Units of Measurement
      1. International System of Units
      2. Types
        1. Mass - Gram
        2. Time - Second
        3. Temperature - Kelvin
        4. Amount of substance - Mol
        5. Volume - Liters or cm^3 (1000)
        6. Density - g/cm^3
    3. Dimensional Analysis
      1. Converting Units to other units
    4. Sig Figs
      1. Non-zero digits are significant
      2. Zeros between nonzero digits are significant
      3. Zeros appearing in front of all nonzero digits are not significant
      4. Zeros at the end aren't significant unless there is a decimal point.
    5. Sig Fig Math
      1. Addition and Subtraction
        1. Number of sig figs to the least decimal place.
        2. 12.11+3.6=15.71 but with sig figs = 15.7
      2. Multiplication and division
        1. Number of sig figs of least sig fig given.
    6. Scientific Notation
      1. # * 10^#
    7. Density
      1. Mass/Volume
      2. Intensive physical property
    8. Accuracy and Precision
      1. Accuracy - how close it is to the correct
      2. Precision - how close all the attempts are to each other
    9. Percentage Error
      1. ( |e-a| / a ) *100
        1. E = what experiment yields
        2. A = what should have been yielded
    10. Direct proportion
      1. y/x = k
        1. Closer k=1 the stronger the positive linear relationship
        2. Closer k=-1 the stronger the negative linear relationship
  3. History of the Atom
    1. Law of Conservation of Mass
      1. Antoine Lavoisier
      2. Mass can’t be created or destroyed
    2. Law of Definite Proportions
      1. Joseph Proust
      2. Chemicals are the same proportion
        1. Water is always H2O
    3. Law of Multiple Proportions
      1. John Dalton
      2. Have to have whole number of atoms
    4. Dalton’s Atomic theory
      1. All matter is composed of atoms
      2. All of a given element is the same
        1. Wrong - Isotopes
      3. Atom’s can’t be split, created, or destroyed
        1. Wrong - Protons, Elections, Neutrons
      4. Atoms of different elements combine in fixed rates to form compounds
      5. Chemical reactions combine, rearrange, and separate atoms
    5. Rutherford’s Gold Foil
      1. Ernest Rutherford
      2. Shoot Alpha particles at gold foil. They hit something and bounced away
      3. Discovered the Nucleus.
    6. Cathode Ray Tubes
      1. JJ Thomas
      2. Like a wire carrying electric current
      3. Rays were deflected away from Negative charge
      4. Discovery of electrons
    7. Plum Pudding Model
      1. JJ Thomas
      2. Negative electrons were spread evenly throughout the the atom
    8. Electron Cloud
      1. Quantum Model
      2. Modern Model
    9. Isotopes
      1. Same number of protons and electrons
      2. Different number of neutrons
    10. Ions
      1. Different number of protons
    11. Average Atomic Mass
      1. Weighted average of all the natural isotopes
      2. One per element
    12. Relative Atomic Mass
      1. Carbon-12 = 1 amu
    13. Notations
      1. Hyphen notation
        1. Carbon-13
      2. Nuclear Symbol

4

        1. HE

2

    1. Mole
      1. SI unit for amount of substance that contains as many particles as there are atoms in exactly 12 grams of Carbon-12
      2. Avogradro's number
        1. 6.02*10^23
  1. Electron Arrangement
    1. Dual Nature of Light
      1. Einstein
      2. Light is both wave and particle
      3. Photon
        1. Electromagnetic radiation with no mass that carries 1 quantum of energy.
    2. Frequency, Wavelength and Energy
      1. Wavelength - λ (lambda)
        1. Measured in meters
      2. Frequency - 𝛎 (nu)
        1. Number of waves that pass per second
        2. Measured in hertz (Hz) or s^-1
      3. Plank's constant - h
        1. 6.626 * 10^-34 Joules per second
        2. Amount of energy 1 photon has
      4. Amplitude
        1. Wave height from origin to crest
      5. Speed of light - C
        1. 3.00 * 10^8
    3. Equations for light
      1. C = λ𝛎
        1. Speed of light
      2. E = h𝛎
        1. Energy in Joules
    4. Max planck
      1. Discovered the smallest amount of energy called quanta
      2. Plank's constant
    5. Dual Nature of light
      1. Einstein
      2. Light is both wave and particle
      3. Photon
        1. Electromagnetic radiation with no mass that carries 1 quantum of energy.
    6. Atomic Emission Spectra
      1. Electricity or heat is added to atom it
        1. Absorbs energy
        2. Not stable
        3. Release energy (light)
        4. Each element has its own color set (frequencies)
    7. Quantum Theory and Atom
      1. Bohr's proposed H atom
        1. Allowed him to predict its emission spectrum
    8. Bohr Model
      1. Nucleus with electrons orbiting in rings
      2. Assigned quantum number, n, to each orbit
      3. Electron in n=1 orbit is in grand state
        1. No energy radiated
      4. Energy electron move to a higher energy orbit are release light
      5. Only works for hydrogen
    9. Quantum Mechanical Model of the Atom
      1. Erwin Schroninger
      2. Particles are both waves and particles
      3. Electrons exist around nucleus in specific orbits
      4. Quantum numbers tell you
        1. Which main level (principle energy level)
        2. How far away from the nucleus
      5. Sublevels of principal energy levels
        1. s, p, d, f
      6. Determine shape of orbital
    10. Quantum Numbers
      1. N (Principal Quantum number)
        1. 1-7
        2. Tells how far electrons are from the nucleus
        3. N > farther away from nucleus
      2. L (Angular momentum)
        1. Tells shape of orbital
          1. S (sphere)
          2. P (Peanut)
          3. D (Double Peanut)
          4. F (Flower)
      3. Ml (Magnetic)
        1. Each sublevel has a specific number or orbital orientations
          1. S - 1
          2. P - 3
          3. D - 5
          4. F - 7
        2. Each orbital can hold 2 electrons
      4. Ms (Spin)
        1. Each pair of electrons has to have opposite spins
        2. Up or down
        3. 2n2 = entire energy level
        4. 2n = max number of electrons
    11. Heisenberg uncertainty principle
      1. States that is is impossible to know both velocity and position of a particle at the same time
      2. Only thing known is the most probably place an electron to occupy
    12. Electron configurations
      1. arrangement of electrons in atom
      2. Ground state configuration
        1. Most stable
        2. Lowest energy arrangement of electrons
      3. Orbital Diagram (Orbital Notation)
        1. Visual representation of electrons fills orbitals
        2. _ _ _ _ _ _ _ _ _ _
      4. Electron configuration
        1. 1s2 2s2 2p6
      5. Noble gas notation (shorthand)
        1. Use the closest noble gas + electron configuration
          1. Closest noble gas before the end of configuration
        2. [Ar] 4s2 3d6
    13. Principles
      1. Aufbau Principle
        1. Electrons occupy the lowest energy orbit
        2. Orbitals in same sublevel are equal
        3. Sublevels within an energy level have priorities (2s than 2p etc..)
      2. Hund's Rule
        1. All same energy orbitals are filled first with electrons containing the same spin before extra electrons can occupy the same orbital with opposite spin.
      3. Pauli Exclusion Rule
        1. Max of 2 electrons may occupy a orbit (same spin)
    14. Valence electrons
      1. Outer energy level / orbital electrons that bond
      2. 1a-8a are just what they are
      3. B groups do not count as easy
        1. Use electron config. To find how many valence electrons in group s and p orbitals have the highest energy level.
  2. Periodic Table and Trends
    1. History of Periodic Table
      1. Lavoisier
        1. gets all elements
      2. John Newlands
        1. proposed arrangement where elements were ordered by increasing atomic mass
        2. Realised when this happened properties repeated
      3. Dmitri mendelev
        1. First periodic table
        2. Predicted elements that didn't exist yet
      4. Henry Mosely
        1. Saw that + charge of the nucleus increased by one unit from one element to the next within theta are arranged on the PT.
        2. Atomic number named
    2. Periodic Law
      1. Physical and chemical properties of the elements are periodic functions of the atomic number
      2. Elements with similar properties recur at regular intervals
    3. Arrangement of the PT
      1. Boxes contain
        1. Element name
        2. Symbol
        3. Atomic number
        4. Average atomic mass
      2. Columns of elements are called groups or families
        1. Have similar properties
      3. Rows of elements are called periods
    4. Properties of Metals/Nonmetals
      1. Metals
        1. Alkali metals
          1. Elements in group 1 except hydrogen
          2. Very reactive
        2. Alkaline earth metals
          1. Group 2
          2. Highly reactive elements but not as alkali
        3. Transition metals
          1. Middle block
        4. Inner transition metals
          1. Lanthanide series

Bottom part of PT

First row

          1. Actinide metals

Bottom party of PT

Bottom row

      1. Nonmetals
        1. Group 17 highly reactive elements
          1. halogens
        2. Group 18 are unreactive or inert
          1. Noble gasses
      2. Metalloids
        1. Have physical properties and chemical properties of metals and nonmetals
        2. Located along staircase
    1. Periodic Trends
      1. Electrons configuration and the PT
        1. s-block
          1. Groups 1 and 2
        2. P-block
          1. Groups 13-18
        3. D-block
          1. Groups 3-12
        4. F block
          1. Inner transition metals
        5. Their period of the element indicates energy level.
      2. Ions
        1. Charged element
        2. Types
          1. Cation

Positive

Loss an electrons

          1. Anion

Negative

Gained electron

      1. Atomic/Ionic Radius and Atomic/Ionic Radii
        1. Single (Radius)
          1. Elements or single ion
        2. Molecule (Radii)
          1. Distance from nuclei to half the distance of the other nuclei
        3. Decrease in atomic/Ionic radius/radii from left to right
          1. Because of increased + charge in the nucleus
        4. Increases as you move down a group
          1. Outermost orbital size increases down a group, making atom bigger
      2. Ionization Energy
        1. Energy required to remove an electron from an atom
        2. First ionization energy
          1. Energy required to remove the first electron
        3. Second ionization energy etc..
          1. More energy required for each one
          2. Especially after all valence electrons are taken away
        4. Trends
          1. Increases Left to right across a period

Amount of valence electrons

          1. Increases down to up

Less energy required to remove an electron on a far orbital

      1. Electronegativity / Electron Affinity
        1. Ability of an element to attract an electron in a chemical bond
          1. High number means it is easier to get 1
        2. Trends
          1. Increases down to up
          2. Increases left to right
  1. Ionic Formulas and Metallic Bonds
    1. Ion Formation
      1. Positive
        1. Cation
        2. Metal
        3. Reactive because they lose Valence electrons easily
      2. Transition metal Ions
        1. +2 and +3 ions
        2. Can be greater than +3
      3. Negative
        1. Anion
        2. Nonmetals
    2. Valence Electrons
      1. They want 8
      2. Amount of electrons in the outer energy level
    3. Properties of Ionic Compounds
      1. Positive and negative ions
      2. Valence electrons shared to create bonds
      3. Bonds create a crystal lattice
      4. Strength of bond relates to MP and BP
    4. Ionic bonds
      1. Ionic bonds form ionic compounds
      2. Between metals and nonmetals
        1. Bond between metal and oxygen is oxide
        2. Salts are the rest
      3. Types of reactions
        1. Reactions that absorb energy are endothermic
        2. Released energy is exothermic
    5. Ionic Formulas
      1. Cation + anion
      2. Monatomic ions are one atom ions
      3. Oxidation number of oxidation state
        1. Charge of a monatomic ion
        2. Cl-1
      4. Polyatomic Ions
        1. OH-1
    6. Metallic Bonds
      1. Crowded lattices of the outer energy levels overlap.
      2. Gives metals their unique characteristics
      3. Sharing their valence electrons for a sea of elections
    7. Alloy
      1. Mixture of elements that have metallic properties
        1. Brass
      2. Types
        1. Substitutional alloys
          1. Original metal's crystal lattice has atoms substituted out for another metal
        2. Interstitial alloy
          1. Smaller atoms go in between the original crystal lattice
  2. Covalent Formulas
    1. Covalent Bonds
      1. Between 2 nonmetals
      2. Diatomics
        1. Go to 7 make a 7 + H
    2. Bonds
      1. How ever many bonds are formed is so that every element has 8 valence electrons
    3. Lewis Structures
      1. Shows elections, elements, and types of bonds formed
    4. Bond Energy, Length, and Strength
      1. The longer the bond the less strength
        1. Bond dissociation energy
    5. Naming covalent compounds
      1. Mono - only for anion
      2. Di
      3. Tri
      4. Tetra
      5. Penta
      6. Hexa
      7. Hepta
      8. Octa
      9. Nona
      10. Deca
    6. Acid Naming and formulas
      1. All acids have H as their cation
      2. Types
        1. Binary
          1. HCL
          2. Add prefix hydro to anion
          3. Change ending of anion to ic
          4. Add acid
          5. Hydrochloric Acid
        2. Oxyacids
          1. Hydrogen + oxyacids
          2. If ate than ic
          3. If ite than ous
  3. Chemical Reactions
    1. Evidence of a reaction
      1. Heat change
      2. Change in color
      3. Change in odor
      4. Gas/bubbles
      5. Appearance of solid
    2. Types of reactions
      1. Synthesis
        1. A+B = AB
      2. Decomposition
        1. AB = A+B
      3. Combustion
        1. CH+O2 = CO2+H2O
      4. Double Replacement
        1. AX+BY = AY+BX
        2. 2 aqueous reactants create 1 aqueous and 1 either solid liquid or gas products
      5. Single Replacement
        1. AX+B = BX + A
        2. Activity series determines if it will occur
        3. Element that replaces compound A has to be more active than it
    3. Understand how to do all of the following reactions
  4. Stoichiometry and Limiting Reactants
    1. Stoichiometry
      1. Study of quantitative relationship in a chemical reaction
    2. Mole Ratio
      1. the ratio between the amounts in moles of any two compounds involved in a balanced chemical reaction.
    3. Limiting Reactants
      1. The chemical that limits more of the reaction from happening
    4. Excess Reactants
      1. The chemical that has extra left over because there is no more limiting react for it to react with
    5. Yield
      1. Actual Yield
        1. What u get form the experiment
      2. Theoretical Yield
        1. What you should have gotten
    6. Empirical Formula
      1. Smallest whole number ratio
    7. Molecular Formula
      1. Actual number of molecules
      2. Multiple of the EF
    8. Percent Composition
      1. Mass of one mol of an element / molar mass of the entire compound
      2. Used to find EF/MF
  6. States of Matter
    1. Kinetic Molecular Theory
      1. All matter is composed of particles that move. The more energy/movement the particles have tells what state they are in.
      2. Gasses
        1. Constant movement
        2. Neither react or repel
        3. Collisions are called elastic collins
      3. Solid
        1. Particles vibrate
    2. Temperature
      1. Effects how much energy is in a particle
        1. Effects what state it is in.
    3. Diffusion/Effusion
      1. Effusion - escapes through tiny opening
      2. Diffusion - diffuses into the air
    4. Units of Pressure
      1. Force / Area
      2. Newton is SI Unit
      3. Torricelli invented the Barometer
        1. Measures pressure
    5. Dalton's Law of partial Pressure
      1. Pressure of a mixture of gases = pressure of every individual gas added up
    6. Lewis Structure = Structural Formula
    7. Resonance Structures
      1. VSEPR model
        1. Theory used to predict the shape of a molecule based on bonds and loose elections
      2. Electron Domain Geometry
        1. Around the central atom
      3. Molecular Geometry
        1. Shape
      4. ED and MG
        1. Linear - 2
          1. Linear - 2

180

        1. Trigonal planar - 3
          1. Trigonal Planar - 2

120

          1. Bent - 2

<120

        1. Tetrahedral - 4
          1. Tetrahedral - 4

109.5

          1. Trigonal pyramidal - 3

107

          1. Bent - 2

105

        1. Trigonal Bipyramidal - 5
          1. Trigonal Bipyramidal - 5

90 120

          1. SeeSaw - 4

120 90

          1. T-Shaped - 3

90 180

          1. Linear - 2

180

        1. Octahedral - 6
          1. Octahedral - 6

90

          1. Square pyramidal - 5

90 180

          1. Square planar - 4

90

    1. Electronegativity
      1. Using electronegativity you can determine the type of bond
        1. X > 1.7 Ionic
        2. X = .4-.17 polar covalent
        3. X < .4 mostly covalent
    2. Polarity
      1. If a molecule has a pos and neg charge than it is polar
      2. Polar Covalent Bond
        1. Unequal share of electrons
        2. Atoms pull on electrons in a molecule unequally
      3. Asymmetrical molecules are polar
      4. Symmetrical molecules are nonpolar
    3. Intermolecular Forces
      1. Attraction force between molecules cause some materials to be solids, liquid, and gasses at room temperature
      2. Dispersion force
        1. London
          1. Weak forces that reacts from temporary shifts in a density of electron in electron cloud
        2. Dipole-Dipole
          1. Attraction between oppositely charged regions of polar molecules
        3. Hydrogen Bond
          1. Special Dipole-Dipole
          2. Only forms form F O N
          3. Very strong
    4. Liquids
      1. Definite volume but not shape
      2. Viscosity
        1. Measure of resistance to flow
          1. As IMF increases viscosity increases

IMF increases when Size, Shape increase or temp decreases

      1. Particle Shape/Size
        1. Larger molecules create greater viscosities
        2. Long chains increases viscosity
      2. Surface tension
        1. Energy required to increase SA at liquid by a given amount
        2. Measure of the inward pull by particles
        3. Surfactants lower ST
      3. Cohesion, Adhesion
      4. Capillary action
        1. Upward movement of liquid in narrow cylinder, or tube
          1. How water moves up a tree
    1. Crystalline solids
      1. Atoms, ions, or molecules in a repeating shape
    2. Amorphous solid
      1. Not repeating pattern
      2. Molten material cools quickly
        1. Plastics
    3. Phase Changes
      1. Require energy if they go up the chain
      2. Release energy when they go down the chain
      3. Sublimation
        1. S to G
      4. Deposition
        1. G to S
      5. Condensation
        1. G to L
      6. Vaporization
        1. L to G
      7. Melting
        1. S to L
      8. Freezing
        1. L to S
    4. Phase Diagram
      1. Graph of Pressure vs. Temp
      2. Shows which phase a substance would be in under those conditions
      3. Triple Point
        1. All 3 phases coexist
      4. Critical Point
        1. Liquid and gas stages are indistinguishable
  1. Gases
    1. Units of Pressure
      1. 1 atm
      2. 760 torr or mmHg
      3. 101.3 kpa
    2. Gas Laws
      1. Boyle’s gas law
        1. PV = PV
      2. Charles Law
        1. V/T = V/T
      3. Gay Lussac's Law
        1. P/T = P/T
      4. Combined Gas Law
        1. PV/T = PV/T
      5. Ideal Gas Law
        1. PV = nRT
          1. R = LP/Mol K
          2. n = moles of substance
        2. D = MP / RT
    3. Gas Stoichiometry
      1. 22.4 L = 1 mol of gas at STP
  2. Solutions
    1. Solutions and solutes/solvents
      1. Solvent is the thing that dissolves the solute
      2. Solute is the thing being dissolved by the solvent
    2. Concentration
      1. Measure of amount of solute in a given amount of solvent
      2. Types
        1. Molarity (M)
          1. Moles of solute / volume of solvent (L)
        2. Molality (m)
          1. Moles of solvent / mass of sheet (kg)
    3. Molarity by dilution
      1. MV = MV
        1. Moles
        2. Volume
        3. V2 is total volume including water
    4. Dissociation in aqueous solutions
      1. Dissociation
        1. Separate cations and anions
      2. NaCL = Na+ + Cl-
    5. Precipitation Reactions
      1. Predict solid formed after double replacement reactions
      2. Soluble
        1. Compound separates into its ions
      3. Insoluble
        1. Compound doesn't separate into its ions
    6. Net ionic equation
      1. Shows only the ions used in the reaction to form the precipitate
      2. Show state of every ion (aq) or (s) etc..
      3. Spectator ions
        1. Ions that don't take part in reaction
      4. Total Ionic equation
        1. Double replacement reaction that has every compound has been pulled apart expect for the non-aqueous compounds
  3. Acids and bases
    1. Properties of acids
      1. Produce H+
      2. Corrode metals
      3. Electrolytes
      4. PH < 7
      5. Turns blue litmus paper red
    2. Properties of bases
      1. Produce OH-
      2. Electrolyte
      3. Feel slippery or soapy
      4. Corrosive
      5. PH > 7
      6. Turns red litmus paper blue
    3. Electrolytes
      1. Strong
        1. Completely dissociates in aqueous solution
      2. Weak
        1. Only dissociations in a small amount in an aqueous solution
    4. Ionization
      1. Process of + or - electrons from one atom or molecule which gives it a charge.
      2. Hydrogen added or taken away from an atom to make it charged.
    5. Arrhenius Acids/Bases
      1. Increase H or OH in an aqueous solution
      2. Acid/Base + water = H3O/OH + ion
    6. Bronsted-Lowry Acids/Bases
      1. Molecule or ion that is a proton donor/acceptor
      2. Made through synthesis
      3. Hydrogen is transferred from one reactant to another
    7. Lewis Acids/Bases
      1. Accepts/gives an electron from its covalent bond
      2. Synthesis
      3. 1 product
    8. Conjugate Acids/Bases
      1. Species that remains after a BL base/acid gains/gives proton
      2. Conjugate Base
        1. Made when It's acid gains a proton
      3. Conjugate Acid
        1. Made when it's base loses a proton
    9. Neutralization Reaction
      1. Reaction of hydronium ions and hydroxide ions to form H2O and Salt
    10. Acid Rain
      1. Greenhouse chemicals react with water and produce a strongly acidic rainwater
    11. Calculating pH and concentration
      1. pH stands for power of hydrogen
      2. pH = -log(H3O+)
      3. pOH = -log(OH-)
      4. KW = (H3O) * (OH)
    12. Indicators
      1. Compounds whose colors are sensitive to pH
        1. Most commonly Phenolphthalein
      2. Used during titrations
      3. Acidity or alkalinity