Chemistry Regents Review Notes

Safety, Scientific Method, and Graphing

• Safety Precautions:

  • Tie back long hair.
  • Rinse arm if a liquid is spilled, then tell the teacher.
  • Dispose of chemicals properly after an experiment.

• Scientific Method:

  • All variables must remain the same except the one being tested.

• Graphing:

  • Use an even scale of numbers and circle final points.

Unit 1: Atomic Structure

• Atoms contain subatomic particles (protons, neutrons, electrons).
• Protons:
  • Have a +1 charge.
• Neutrons:
  • Have a 0 charge.
  • Mass of Neutron = 1 amu = Mass of Proton
• Electrons:
  • Have a charge of -1 and a mass of 0.
• Charge Comparison:
  • Electron charge is -1, proton charge is +1 (same magnitude, opposite charge).
• Table O:
  • Shows symbols, mass, and charges of particles (electrons are represented as beta).
• Location of Protons and Neutrons:
  • Located in the nucleus of an atom.
• Charge of Nucleus:
  • Charge of an atom’s nucleus = (+) number of protons.
• Structure of Atoms:
  • Atoms have a positively charged nucleus and negatively charged electrons located in “clouds” (orbitals) around the nucleus.
• Mass Number:
  • Mass \ number = #protons + #neutrons
• Atomic Number:
  • Atomic number = number of protons (All atoms of the same element have the same atomic number).
• Number of Neutrons:
  • $Number \ of \ neutrons = Mass \ number – atomic \ number$
• Isotopes:
  • Atoms with the same number of protons, different number of neutrons (different mass number).
• Isotope Notation:
  • Top number of isotope notation is Mass number and bottom number is atomic number.
• Other Notations:
  • (C-14 or Carbon-14) Number after an element represents mass number.
• Atomic Mass:
  • The weighted average of all the naturally occurring isotopes for that element.
• Average Atomic Mass Calculation:
  • $Average \ Atomic \ mass = (isotope1 \ mass)(\% \ in \ decimal \ form) + (isotope \ 2 \ mass)(\% \ in \ decimal \ form)$
• Abundance:
  • Whatever whole number the atomic mass is closest to on the Periodic Table means that isotope is most abundant.
• Rutherford’s Gold Foil Experiment:
  • Shows an atom is mostly empty space with a small, dense, positively charged nucleus.
• Thomson and Bohr’s Models:
  • Showed electrons present in an atom.
• Wave-Mechanical Model (Electron Cloud Model):
  • An orbital (cloud) is the most probable location of electrons.
• Neutral Atom:
  • An Atom has the same number of protons and electrons as long as there is no charge (total charge of 0).
• Total (Net) Charge of an Atom:
  • Total \ (Net) \ Charge \ of \ an \ atom = # \ protons - # \ electrons
• Ion:
  • A charged element (it has lost or gained electrons): electron configuration will change if it is an ion (possible charges are found on PT).

Electron Configuration and Energy Levels

• Electron Configuration:
  • Shows location of electron in their shells. Example: 2-8-2 (2 electrons in first energy level, 8 electrons in second energy level, 2 electrons in third energy level) CAN BE FOUND ON PERIODIC TABLE
• Energy Levels:
  • First shell has less energy than 2nd shell.
• Electron Capacity:
  • First shell can have a max of 2 electrons and Second shell can have a max of 8 electrons.
• Valence Electrons:
  • Electrons found in the outer most shell (last number in electron configuration).
• Lewis Dot Diagram:
  • For a single atom just shows the valence electrons (electrons represented by dots, drawn in pairs).
• Excited Electrons:
  • When excited electron (at higher energy level) moves to ground state (lower energy level), a specific amount of energy is emitted (sometimes as light/bright line spectrum)
• Excited Electron Configuration:
  • Is not the same as the configuration on the reference table (there will be one less electron in one energy level and one more electron in another energy level).
• Energy Emission:
  • Energy emitted from excited electron can be used to determine identity of element
• Bright Line Spectrum:
  • When viewing a Bright line spectrum- elements must line up exactly to be part of the mixture in the spectrum

Unit 2: Nuclear Chemistry

• Table O:
  • Shows a positron and beta particles have same mass (0) & opposite charge (Beta negative, positron positive).
• Particle Penetration:
  • Beta particle has less mass and greater penetrating power than a +alpha particle (Gamma radiation has the greatest penetrating power).
• Nuclear Reactions:
  • All Nuclear reactions are transmutations (examples: fission, fusion, decays)
• Unstable Elements:
  • Any element after Po is naturally unstable and will spontaneously decay
• Stable Isotopes:
  • Are not on table N (do not spontaneously emit particles)
• Decay Modes:
  • Decay modes and half lives on table N (show alpha, beta, positron decay)
• Natural Transmutations:
  • Show Spontaneous radioactive decay = 1 reactant ! 2 products (elements must change)
• Nuclear Decays:
  • Release the decay particle
• Completing Nuclear Equations:
  • Remember to add up mass number and atomic numbers on each side of the arrow (must be equal on both sides)
• Fusion:
  • Light nuclei combine to form a heavy nucleus and a lot of energy (energy is sometimes in the form of a neutron)
• Fusion Energy:
  • Fusion produces more energy than fission
• Energy Release:
  • Nuclear reactions (such as fission or fusion) releases more energy than a chemical reaction (redox, substitution, neutralization)
• Mass Conversion:
  • Nuclear reactions: mass is converted into energy
• Half Life:
  • The length of time it takes for ½ mass of a sample to decay
• Half Life Sample:
  • 1 half life = ½ sample remains, 2 half lives = ¼ sample remains, 3 half lives = 1/8 sample
• Half Life Question:
  • Use table or timeline method
• Radioisotope Applications:
  • Radioisotopes are used for dating of geological formations (C-14)
  • I-131 used to diagnose thyroid disorders
  • Radioisotopes can be used to detect diseases
  • Radioisotopes can treat cancer but can also cause mutations in healthy cells (Co-60)

Unit 3: Matter

• Substance:
  • Substance = compound or element
• Elements:
  • Elements cannot be broken down by chemical means (it is an element if it is on the Periodic table/Table T)
• Compounds:
  • Can be broken down by chemical means
• Same Compound:
  • Same compound = same chemical property, different compound = different chemical properties
• Diatomics:
  • 7 diatomics (two of the same atom bonded together) – BrINClHOF (Br2 ,I2 , N2 , Cl2 , H2 , O2 , F2)
• Table T:
  • Melting Point (MP) and Boiling Point (BP) of elements on table T
• Liquid State:
  • MP < Specific Temperature < BP
• Mixture:
  • Can vary in proportion of its components (example: Salt water)
• Homogeneous Mixtures (Solutions):
  • Even distribution of particles (aq-dissolved in water) substance has to be soluble to mix with water
• Heterogeneous Mixtures:
  • Not even throughout. Contains a substance that will not be soluble in water.
• Substance Properties in Mixtures:
  • When substances are mixed, they retain their properties
• Mixture Separation by Physical Means:
  • Mixtures that contain substances with different density and particle size can be separated by physical means
• Separation Techniques:
  • Mixture can be separated by chromatography, distillation and filtration
• Distillation:
  • Separates liquids with different boiling points (water and alcohol)
• Chromatography:
  • Method of separating particles by solubility and polarity
• Evaporation:
  • Separates a salt dissolved in water
• Chemical Property:
  • Is how substances react
• Chemical Change:
  • Results in the formation of a difference substance (example: burning)
• Physical Change:
  • Do not form new compounds, commonly phase changes (change in distance between molecules)
• States of Matter:
  • Solids = atoms close together, liquid in the middle, gas = atoms far apart
• Solid Properties:
  • Solids have a definite shape and definite volume
• Deposition:
  • Gas to solid phase change
• Sublimation:
  • Solid!Gas phase change (ex: CO2)
• Phase Change Diagram:
  • In a phase change diagram, the flat parts represent the phase changes (Potential Energy [PE] changes and Kinetic Energy [KE] remains the same)
  • In a phase change diagram, the sloped lines represent heating or cooling (PE remains the same and KE changes)
• Density:
  • $Density = mass/volume \ (g/L \ or \ g/cm^3)$
• Density and Position:
  • Higher density sinks to bottom of tank
• Element Density:
  • Density never changes for each element (Found on Table S for elements)

Math and Significant Figures

• Significant Figures (Sig Figs):
  • Atlantic/Pacific Rule:
    • Decimal Absent: Count from the first nonzero number on the Atlantic side (right) and count all numbers to the left.
    • Decimal Present: Count from the first nonzero number on the Pacific side (left) and count all numbers to the right.
• Multiplying/Dividing:
  • Answer should be the lowest number of sig figs.
• Percent Error Formula:
  • On table T.

Unit 4: Energy

• Kelvin Conversion:
  • $Kelvin = °C + 273$
  • (0°C = 273 K)
• Kilojoules to Joules:
  • 1 kJ = 1000 J
• Forms of Energy:
  • Chemical, thermal, electromagnetic, electrical, nuclear, mechanical
• Thermal Energy:
  • (heat) is measured in joules (J) = random motion of atoms and molecules
• Average Kinetic Energy:
  • Average kinetic energy = temperature
• Thermal Energy and Mass:
  • When two substances have the same temperature, the substance with the greater mass has more thermal energy
• Heat of Vaporization:
  • Amount of heat required to vaporize a substance (table B for water constants) = 2260 J or $2.26 \times 10^3$ J
• Heat of Fusion:
  • Heat it takes to melt a substance is less than heat of vaporization because it requires less heat to melt a substance than boil a substance.
• Heat Flow:
  • Heat flows from hot to cold
• Heat Equation:
  • $q = mCΔT$
    • q is heat, m is mass, C is specific heat capacity [found on table B for water], ΔT is change in temperature. All info on table T
• Exothermic:
  • Energy exits (is released).
• Endothermic:
  • Energy absorbed (heat is shown on left side of equation) (examples of endothermic phase changes: s!l, l!g, s!g)

Unit 5: Gas Laws

• Pressure Effects:
  • Pressure only effects gases
• Pressure Measurement:
  • Pressure is measured in pascals
• STP Conditions:
  • Standard Temperature and Pressure (STP) on table A (273 K and 1 atm or 101.3 kPa and 0C)
• Pressure and Temperature Relationship:
  • Pressure and temperature have direct relationship (as pressure increases, temperature increases)
• Ideal Gas Conditions: PLIGHT
  • Pressure Low, Ideal gas, High Temperature
• Ideal Gas Motion:
  • Ideal gases move in random, constant, straight line motion
• Ideal Gas Separation:
  • Ideal gases are separated by great distances compared to their size
• Ideal Gas Forces:
  • Ideal gases have no attractive forces
• Gas Collisions:
  • Collisions of gas may result in a transfer of energy
• Table H:
  • Dotted line is the normal boiling of the substance
• Intermolecular Forces (IMF):
  • Gases have the weaker IMF than solids

Unit 6: Periodic Table

• Mendeleev's Table:
  • Organized by atomic mass.
• Modern Periodic Table:
  • Elements arranged in order of atomic number.
• Periods and Groups:
  • Periods are horizontal rows.
  • Groups are vertical rows.
• Liquids on the Periodic Table:
  • Bromine and Mercury.
• Metals:
  • (left of staircase) are good conductors of heat, malleable.
• Metal Properties:
  • Malleable because of the nature of the bonds between the atoms.
• Metal Electrons:
  • Have fewer valence electrons than nonmetals.
• Metal Reactivity:
  • Commonly react with nonmetals.
• Nonmetals:
  • To the right of the staircase.
• Metalloids:
  • Are on the staircase.
• Transition Metals:
  • Form colored solutions in aqueous ion form.
• Halogens:
  • (Group 17) – forms halide ions in solution.
• Noble Gases:
  • (Group 18) are stable because of their stable electron configuration (8 valence electrons), they are not reactive
• Gaining/Losing Electrons:
  • Elements will gain or lose electrons to be like noble gases
• Ions Formed by Metals:
  • MELPS- Metals Electrons Lost form Positive Smaller ions
• Oxidation Numbers:
  • Some groups will form the same oxidation numbers when gaining or losing electrons
• Group Properties:
  • Same group/family = similar chemical properties because they have the same number of valence electrons
• Electronegativity:
  • Attraction for electrons in a chemical bond = electronegativity (found on table T)
• Ionization Energy and Electronegativity Trends:
  • First Ionization energy and electronegativity increases left to right (across a period) and decreases top to bottom (down a group)
• Atomic Radius Trends:
  • Atomic radius (on table T) decreases left to right (across a period) and increases top to bottom (down a group)
• Atomic Radius and Energy Shells:
  • Atomic radius increases down a group because more energy shells are added

Unit 7: Naming, Formulas, and Equations

• Ionic Formulas:
  • Metal comes first, nonmetal second. Criss cross charges to get formula. Example: Calcium phosphate
• Naming Ionics with Table E:
  • Say metal and then ion from table E. Ammonium examples shown
• Roman Numerals:
  • Represent the charge of the metal
• Polyatomic Ions:
  • Names on table E
• Balancing Equations:
  • Number of each atom needs to be the same on both sides of the equation (use tallies). Use coefficients to balance
• Conservation:
  • Mass, charge and energy are conserved in a chemical reaction
• Types of Chemical Reactions:
  • Decomposition, single replacement, double replacement, synthesis
• Synthesis:
  • Two or ore reactants combine to form one product
• Decomposition:
  • One compound is broken down into two
• Single Replacement:
  • One element switches partners
• Double Replacement:
  • Think Do Si Do

Unit 8: Bonding

• BARF:
  • Bond BROKEN energy ABSORBED (endothermic), energy RELEASED bond FORMED (exothermic)
• Compound Types:
  • Ionic compound (metal and nonmetals or table E ions) and molecular compounds are all nonmetals with no table E ions
• Ionic Bonds and Table E:
  • Ionic compounds with a table E ion have covalent and ionic bonds
• Ionic Compound Properties:
  • Hard brittle solids with a high melting point, poor conductors as solids and good conductors in aq (higher concentration = better conductor)
• Metallic Bonding:
  • Metals only!
• Ionic Bonds:
  • Bond between metal and nonmetal or ion from table E
• Ionic Bond Electron Transfer:
  • Transfer electrons from the valence shell of one atom to the valence shell of another atom
• Lewis Dot Diagrams:
  • Have brackets for ionics. Molecular compounds: make sure all elements have eight electrons (except H) and the total number of electrons in diagram = total number of valence electrons
• Covalent Bond:
  • Nonmetals (molecular compound)
• Nonpolar Covalent Bond:
  • Bond between two of the same elements
• Multiple Covalent Bonds:
  • Double (four electrons shared/ 2 pairs), triple (six electrons shared/ 3 pairs)
• Diatomic Molecules:
  • Diatomic Oxygen molecule has double bond, diatomic nitrogen has triple
• Polarity of Bond:
  • Electronegativity difference = polarity of bond (higher the difference, the more polar the bond)
• Molecule Polarity: SNAP
  • Symmetrical Nonpolar, Asymmetrical Polar
• Intermolecular Forces (IMF):
  • Stronger Intermolecular forces (IMF) = High boiling point
• Hydrogen Bonding:
  • Example of IMF: hydrogen bonding (between H and F,O or N)
• Water Molecule Charges:
  • (Oxygen is partially negative and hydrogen is partially positive)
• Charge Attraction:
  • Positive and negative charges attract, like charges repel

Unit 9: Stoichiometry

• Atomic Mass:
  • Atomic mass on the Periodic Table is in grams/ 1 mole
• GFM:
  • (gram formula mass, formula mass or molar mass) is found by getting the mass of each element (multiplied by their subscript if they have one) and add all masses together
• Percent Composition by Mass:
  • $\% \ Composition \ by \ mass= mass \ of \ part / mass \ of \ whole$
  • (Table T)
• Calculating Percent Composition:
  • Can be calculated given a periodic table and chemical formula
• Mole Calculation:
  • $Mole = mass/GFM$
  • (Table T)
• Mole Ratio:
  • Mole : mole ratio = coefficient : coefficient ratio
• Empirical Formula:
  • Simplified formula, molecular is normal not simplified formula
• Conservation of Mass:
  • Mass of products = mass of the reactants

Unit 10: Solutions

• Solubility Dependence:
  • Solubility is dependent on temperature for solids (upward curve on table G)
• Table G:
  • Use when given a temperature and asking about a salt dissolved in water (on the line is saturated, below = unsaturated, above = super)
• Solubility Rule:
  • Likes dissolve likes (example: polar mixes with polar substances)
• Table F:
  • (insoluble compounds will not fully dissolve in water and will be in solid phase, soluble compounds will dissolve in water and be in aq phase)
• Ion Effect:
  • More ions more change in BP/FP
• Freezing Point Change:
  • Add something to water (such as a salt) = freezing point decreases (more salt = lower the freezing point)
• Boiling Point Change:
  • Add something to water (such as a salt) = BP increases (more salt = higher BP)
• Concentration Units:
  • m/L or PPM
• Parts Per Million (PPM):
  • $PPM = (grams \ of \ solute / grams \ of \ solution) \times 1,000,000$
  • [formula on table T]
• Molarity Calculation:
  • $Molarity = moles/L$
  • formula on Table T

Unit 11: Kinetics

• Reaction Conditions:
  • For a chemical reaction to occur you need sufficient energy and proper orientation for new bonds to form
• Effective Collisions:
  • Chemical reactions need effective collisions
• Temperature Effect:
  • Increase temperature, increases reaction rate (more collisions)
• Pressure Effect:
  • Increase pressure: increases reaction rate (more collisions)
• Concentration Effect:
  • Lower concentration slow reaction rate because of less collisions
• Surface Area Effect:
  • Greater surface area = higher rate of reaction
• Potential Energy Diagrams:
  • Show energy of reaction and when a catalyst is add the energy of the reactants and products do not change
• Activation Energy:
  • Goes from the reactants to top of curve
• Reverse Activation Energy:
  • Goes from the products to top of curve
• Exothermic Reaction Energy Diagram:
  • Potential energy diagram starts high and ends low for an exothermic reaction
• Heat of Reaction:
  • Heat of reaction = potential energy of products – potential energy of the reactants
• Heat of Reaction Location:
  • Heat of reaction (ΔH) is in the middle of a potential energy diagram
• Table I:
  • Show heat of reaction (energy absorbed + ΔH energy released – ΔH
• Catalyst Action:
  • Speeds up a reaction by providing a different reaction pathway that lowers the activation energy
• Entropy (Disorder):
  • Gases have the highest entropy, then liquid/aq and solids have the lowest
• Nature's Trend:
  • Nature undergoes changes towards higher entropy and lower energy

Equilibrium

• Equilibrium Definition:
  • Forward reaction ! Reverse reaction " Equilibrium "
• Equilibrium State:
  • At equilibrium the concentration of reactants and products remain constant because reaction rates of the forward and reverse reaction are equal
• Reaction Rates:
  • Rate of forward reaction = rate of reverse reaction
• Closed System:
  • Need a closed system to maintain equilibrium
• Equilibrium Shift:
  • Shift to right make more products, shift to left make more reactants
• Equilibrium Principles:
  • UP and AWAY, DOWN and TOWARDS
• Pressure Shift:
  • Decrease pressure shift to the side with most number of moles
• Temperature Shift:
  • Increase temperature shift away from heat
• Catalyst Effect:
  • Adding catalyst= no shift
• Solution Equilibrium:
  • Rate of dissolving equals rate of crystallization
• Dynamic Phase Equilibrium:
  • Two phase changes going on in a sealed flask

Unit 12: Acids and Bases

• Electrolytes:
  • Acids, bases, salts
• Electrolyte Conductivity:
  • Conduct in aqueous solutions not in solid phase
• Arrhenius Theory:
  • Describes the behavior of acids and bases
• Arrhenius Definitions:
  • Acids yield H+ ion and bases have an OH-
• Acids & Bases on Tables K & L:
  • Know CH3COOH is an acid, all other CH compounds are not electrolytes
• Alternate Acid/Base Theory:
  • BAAD: BASES ACCEPT H+ ACIDS DONATE H+
• Acidity and pH:
  • Most acidic compounds = lowest pH
• pH and Hydronium Ions:
  • Lower the pH (more acid) = more hydronium ions
• pH Scale Changes:
  • Concentration of hydronium ions increases by a factor of 10 when the pH goes down by 1 pH unit
• Neutral Solutions:
  • Moles of OH- and H+ are equal in a neutral solution
• Neutralization Reactions:
  • Acid + base ! salt + water
• Metal and Acid Reactions:
  • Any metals above H2 on table J will react spontaneously with an acid
• Titration Equation:
  • $MaVa= MbVb$
  • remember M is called molarity or concentration
• Titration Purpose:
  • Determines the concentration of unknown acids or bases
• Indicators:
  • Range on the reference table (table M) is when the color is in-between two primary colors

Unit 13: Redox Reactions

• Oxidation Number:
  • Single atom = 0, look on reference table or use algebra for atoms with multiple charges (can also use reverse criss cross if all else fails)
• Electron Transfer:
  • Redox reactions electrons are transferred (charge changes)
• Oxidation and Reduction:
  • Oxidation: electrons are lost (electrons on right), Reduction: electrons are gained (electron on left) LEO the lion says GER
• Half Reactions:
  • Show charges and electrons lost
• Balanced Redox:
  • Balanced redox reaction has same number of electrons lost and gained (atoms and charges balanced)
• Table J Reactivity:
  • Lower on table J is less reactive
• Active Metals:
  • More active metals are higher on table J
• Oxidation/Reduction on Table J:
  • Higher on table J more likely to be oxidized for metal side, lower more easily reduced
• Spontaneous Reactions on Table J:
  • If the oxidized element is higher on table j then the reaction will be spontaneous
• Electrochemical Cell:
  • AN OX RED CAT (anode oxidation, reduction cathode)
• Electron Flow:
  • Electrons flow from anode to cathode
• Anode Mass:
  • Anode mass will decrease because solid metal is going into aqueous
• Voltaic Cell:
  • Chemical energy is converted to electrical energy spontaneously
• Electron Movement:
  • Electrons flow through a wire for voltaic cells
• Ion Movement:
  • Ions move through the salt bridge for a voltaic cell
• Electrolytic Cell:
  • Electrical energy to chemical energy
• Power Source:
  • Battery (power source) is needed for an electrolytic cell to provide electrical energy (electrolysis or electroplating)
• Electroplating:
  • Cathode in an electrolytic cell is what is electroplated (key or spoon)

Unit 14: Organic Chemistry

• Organic Compounds:
  • Must contain at least one carbon and one hydrogen
• Hydrocarbons:
  • Contain only carbon and hydrogen
• Saturated Hydrocarbons:
  • Have all single bonds (2 electrons shared)
• Unsaturated Compounds:
  • Have double or triple bonds between carbons
• Organic Formulas:
  • Empirical, structural, molecular
• Structural Formula:
  • Shows bonds
• Carbon Bonding:
  • Carbon (atomic number 6) can form chains, rings and networks
• Table P:
  • Tells number of carbons (remember carbon always makes 4 bonds!)
• Homologous Series:
  • On table Q (alkanes, alkenes, alkynes)
• Naming Alkenes:
  • (use table Q), start from side with double bond and count number of carbons, put number to represent location of double bond
• Carbon Counting:
  • Count carbons from side closest to double bond or functional group
• Isomers:
  • Have same number of C’s and H’s but are in a different structural arrangement (same molecular formula)
• Isomer Properties:
  • Different structure = different physical/chemical properties
• Functional Groups:
  • Table R
• Functional Group Properties:
  • Different Functional groups have different chemical properties
• Alcohols:
  • Drawing structure of alcohol includes putting an OH
• Halides:
  • Contains group 17 elements
• Organic Acids:
  • Structure has C with a double bonded O and OH
• Amines:
  • Contain N
• Organic Reactions:
  • Esterification, addition, saponification, polymerization, fermentation, etc.
• Addition Reaction:
  • Binary compound (two atoms bonded together) gets added to double/triple bonded compound to reduce the number of bonds
• Saponification:
  • Organic reaction used to make soap
• Polymerization:
  • Adding same compound together to make a chain
• Fermentation:
  • Sugar + enzyme ! ethanol (alcohol) + carbon dioxide

Test Taking Skills

• Use the Reference Tables!
• Read Paragraphs Completely:
  • Information you need MIGHT be in there
• Answer Every Question.
• Best Guess Approach:
  • If you don’t know the answer take the best guess you can make
• Referring to the Reference Table:
  • Refer to the reference table for any question you may have trouble on
• Using the Test to Take the Test:
  • The answer or a hint may be in another question
• First Choice:
  • Your first choice is usually your best choice (unless you read the question incorrectly the first time around)
• Double Check Reference Table:
  • Even if you think you know a chemical symbol, formula or charge….look it up an the reference table anyway (Table S and Periodic Table help a lot with this)
• Skipping Questions:
  • Skip a question if you are having a hard time and go back to it later
• Time Management:
  • You have plenty of time to take the regents (3 hours). Take a 5 minute break after you are done with the test and then look over your answers.
• Healthy Habits:
  • Eat a healthy meal the night before and for breakfast as well.
  • Get a good night’s sleep. A tired mind is not as sharp and clear as a well-rested one.
• Relax!
  • You have seen all this stuff before…it is somewhere in your brain!