Chemistry Regents Review Notes

Chemistry Regents Layout

  • 85 Questions Total: 50 multiple choice + 35 short answer questions split into multiple parts.
    • Part A: 30 multiple choice questions that assess knowledge and understanding of all chemistry topics.
    • Part B-1: 20 multiple choice questions that are content AND skills based. Designed to assess ability to use the Chemistry Reference Tables AND apply, analyze, and evaluate chemistry material.
    • Part B-2: 15 short answer questions that test skills like performing numerical calculations, drawing and interpreting graphs, drawing diagrams, and applying chem concepts to solve short-answer problems.
    • Part C: 20 short answer questions that assess ability to apply knowledge of scientific concepts and skills to address real-world situations (labs or scenarios).
  • Need about 50 out of the 85 questions correct to pass with a score of 65.

Unit 1: Matter

  1. Substance refers to a compound OR an element.
    • Elements are made of only one kind of atom and CANNOT be broken down by chemical means. An element is on the Periodic Table/Table S. Ex. Sn (Tin)
    • Compounds are made of 2 or more different kinds of atoms, CHEMICALLY combined in a FIXED proportion. Can be broken down by chemical means Ex. NH3 (Ammonia)
  2. Same compound = same chemical property, different compound = different chemical properties
  3. Different structure = different chemical and physical properties
  4. 7 diatomics (two of the same atom bonded together, exist in pairs) – BrINClHOF “Brinklehoff” (Br2 ,I2 , N2 , Cl2 , H2 , O2 , F2)
  5. Physical properties include Melting Point (MP) and Boiling Point (BP). These are listed on table S for many elements
    • If given temperature is lower than melting point = solid Ex. At STP (273 K) Li is a solid
    • If given temperature is higher than boiling point = gas Ex. At STP (273 K) He is a gas
    • If given temperature is in between the melting and boiling points = liquid Ex. At STP Hg is a liquid
  6. Mixtures can vary in proportion of its components (example: Salt water)
    • Homogeneous mixtures (unsaturated and saturated solutions and aqueous solutions): even distribution of particles (particles are spread out evenly) (aqueous = dissolved in water)
      • Soluble compounds (dissolve) in water will make a homogeneous mixture (unsaturated or saturated). Check Table F! Ex. NaNO3
    • Heterogeneous mixtures: uneven distribution of particles (particles are UNevenly spread out).
      • INsoluble compounds will make a heterogenous mixture (stay as solid and sink to the bottom so particles are unevenly spread out). Ex. AgCl
  7. When substances are mixed, they retain their original properties
  8. Mixtures that contain substances with different density and particle size can be separated by physical means
  9. Mixture can be separated by chromatography, distillation and filtration
    • Distillation separates liquids with different boiling points (ex. water and alcohol)
    • Boiling or Evaporation separates a salt dissolved in water because water has a much lower boiling point than salt.
    • Chromatography: method of separating particles based on molecular polarity. Ex. paper chromatography to separate different colored ink
    • Filtration separates out insoluble solids or large particles from liquids (ex. Sand from water) due to particle size (sand particles are large)
  10. Chemical properties refer to properties observed during a chemical reaction (NOT during a physical change like melting, freezing, etc.)
  11. Chemical change results in the formation of a difference substance (example: burning). Results in bonds being broken and formed and NEW ARRANGEMENT)
  12. Physical change: do not form new compounds, commonly phase changes (change in distance between molecules only. Molecules otherwise stay the same, no new substance forms)
  13. Deposition = gas to solid phase change, Sublimation: Solid to Gas phase change (ex: CO2)
  14. Solids = close together, neatly organized. Definite shape and volume. Liquid = close together, disorganized. Takes shape of container, has its own definite volume. Gas = far apart, disorganized. Takes the shape and volume of container.
  15. Particle diagrams can show elements, compounds, mixtures and solids, liquids, gas.
    • One type of circle = one atom of an element
    • Two or more different types of circles connected = compound
    • Two or more SAME type of circle = element only
  16. Density = mass/volume (g/Lg/ L or g/cm3g/cm^3). Higher density sinks to the bottom of the container. Density never changes for each element, no matter the size, mass or volume of it you have. (Found on Table S for elements)

Unit 2: Units/Measurement

  1. Percent Error Formula is Found in Table T. Sometimes the actual value is given, sometimes you have to look up the ACTUAL value on Table S for an element (ex. Density, boiling point, atomic radius, etc.)
  2. Significant Figure (sig figs) Rules!
    • Rule #1: All nonzero digits are always significant. Ex. 123 = 3 sig figs
    • Rule #2: Zeros between two nonzero digits are significant. Ex. 10.02 = 4 sig figs
    • Rule #3: Zeros in FRONT of the first nonzero digit are NOT significant. Ex. 0.002 = 1 sig fig
    • Rule #4: ANY digits after the first non-zero in a number WITH a decimal ARE significant. Ex. 200.0 = 4 sig figs
    • Rule #5: Zeros at the end of numbers without a decimal point are NOT significant. 300 = 1 sig fig
    • Rule #7: For numbers in scientific notation, count the significant figures in the number only. Ignore the “x 10exponent” 1.200 x 103 = 4 sig figs.
    • Overall: Zeros in the front of numbers do not count. Zeros in the back of numbers do. Ex. 0.002300 = 4 sig figs
  3. When you multiply or divide numbers, your answer should be expressed with the same # of sig figs as the number with the least # of sig figs in the problem. Ex. 3.0 (2 sig figs) x 7.35 (3 sig figs), answer should be rounded to 2 significant figures because 3.0 has 2 sig figs, which is the least number of sig figs in the problem.
  4. Units, their symbols, and what they measure are listed on Table D. Prefixes, their factors and symbols are on Table C. Ex. 1 kJ = 1 kilojoule = 10310^3 joules so 1 kJ = 1000 J
    • To find how many kJ if you have Joules, divide by 1000.
    • To find how many J if you have kJ, multiply by 1000.

Unit 3: Energy

  1. Forms of energy: chemical, thermal (heat), electromagnetic, electrical, nuclear, mechanical
    • Thermal energy (heat) is measured in joules (J) =due to random motion/vibration of atoms and molecules
  2. Temperature: measure of average kinetic energy (temperature is NOT a form of energy)
    • As temperature increases, average kinetic energy increases. As temperature decreases, average kinetic energy decreases
  3. Kelvin = °C +273 formula on table T
    • 0 ° C = 273 K
    • –273 K = 0 K = absolute 0
  4. When two substances have the same temperature, both substances have the same average kinetic energy BUT the substance with the greater mass has more thermal (heat) energy.
  5. Heat energy ALWAYS travels from HOT (substance with higher temperature) to COLD (substance with lower temperature).
  6. Three Equations on Table T to solve for Thermal/Heat Energy. Pick the CORRECT one based on info given
    • q=mCΔTq = mCΔT
      • q is heat, m is mass, C is specific heat capacity, ΔT is change in temperature
      • Use when temperature is changing, NOT when there is a phase change
    • q=mHfq = mH_f
      • q is heat, m is mass, HfH_f is heat of fusion
      • Use when there is a phase change between liquid and solid
      • Heat of fusion is the amount of heat it takes to melt a substance from solid to liquid
    • q=mHvq = mH_v
      • q is heat, m is mass, HvH_v is heat of vaporization
      • Use when there is a phase change between gas and liquid
      • Heat of vaporization is the amount of heat required to vaporize/boil a substance (liquid to gas).
      • The heat of vaporization is generally greater than the heat of fusion for a substance because it usually requires more energy to vaporize/boil a liquid into the gas phase than it does to melt a solid into the liquid phase.
  7. FOR WATER ONLY: Constant values for Specific Heat Capacity, Heat of Fusion and Heat of Vaporization can be found on Table B!
    • Specific Heat Capacity FOR WATER: 4.18 J/g- oC or 4.18 J/g-K (Same value for both oC or K, DO NOT convert/change the value for specific heat capacity or temperature when doing heat equations)
    • Heat of Fusion FOR WATER (Solid ←→ Liquid) = 334 J/g
    • Heat of Vaporization FOR WATER(Liquid ←→ Gas) = 2260 J/g
    • *FOR ANY OTHER SUBSTANCE (NOT WATER) THE PROBLEM WILL GIVE YOU THE CONSTANT VALUE. Do NOT use any water constant values on table B if the question is not about water.
  8. Endothermic = energy absorbed/gained/required (heat is shown on LEFT side of equation–reactant). Phase changes that are endothermic are:
    • Boiling/Vaporization (Liquid to gas)
    • Melting (Solid to Liquid)
    • Sublimation (Solid to Gas Directly)
  9. Exothermic = energy is released/lost/exiting (heat energy is shown on the RIGHT side of the equation–product). Phase changes that are exothermic are:
    • Condensation (Gas to liquid)
    • Freezing (Liquid to Solid)
    • Deposition (Gas to Solid Directly)
  10. Heating/Cooling Curves (aka Phase Change Graph)
    • If temperature is generally increasing going from left to right, substance is being heated (heating curve)
    • If temperature is generally decreasing going from left to right, substance is being cooled (cooling curve)
    • The diagonal, sloped lines represent heating (if line is going up) or cooling (if line is going down) because temperature is changing.
    • The flat lines represent phase changes because temperature is constant during a phase change.
      • Lower flat line represents melting (solid to liquid) in a heating curve and freezing (liquid to solid) in a cooling curve. Temperature at this line is the melting/freezing point of a substance.
      • Higher flat line represents boiling (liquid to gas) in a heating curve and condensation (gas to liquid) in a cooling curve. Temperature at this line is the boiling/condensation point of a substance.
      • BE CAREFUL! Sometimes the graph only shows PART of a substance’s heating or cooling curve. Read the information provided carefully to see what phases you are starting and ending with in the graph that is given.
    • Kinetic energy always matches what is happening with temperature.
      • If temperature is increasing, kinetic energy is increasing.
      • If temperature is decreasing, kinetic energy is decreasing.
      • If Temperature is constant, kinetic energy is constant.
    • If kinetic energy changes, potential energy is constant. If kinetic energy is constant, potential energy is changing. They can never do the same thing at the same time.
      • If temperature is increasing, potential energy is constant
      • If temperature is decreasing, potential energy is constant
      • If temperature is constant and substance is being HEATED, potential energy increases
      • If temperature is constant and substance is being COOLED, potential energy decreases

Unit 4: Gas

  1. Gases behave MOST like an IDEAL gas at HIGH Temperature and LOW Pressure (like us during the summer!)
  2. Gases behave LEAST like an IDEAL gas at LOW Temperature and HIGH Pressure (like us during January finals time!)
  3. Kinetic Molecular Theory states that an IDEAL gas:
    • move in random, constant, straight line motion
    • are separated by great distances compared to their size
    • have no attractive forces (or weak attractive forces)
    • Individually occupy (take up) negligible volume (each individual particle are so tiny that it practically takes up no space)
    • completely transfer energy when they collide with one another
  4. Avogadro’s Hypothesis: When two samples of gas have the same pressure, volume and temperature, they also have the same number of molecules/particles!
    • P, V, T and # of molecules. If 3 of these 4 are the same, the last one also must be the same.
  5. Pressure is measured in pascals (Pa), kilopascals (kPa), atmospheres (atm)
  6. Pressure only affects substances that are in the gas phase. Pressure cannot compress (decrease the volume) of substances in the solid or liquid phase (at least in Regents Chem)
  7. Pressure is caused by collisions of gas molecules against the walls of the container. **More forceful (powerful) and more frequent (more often) collisions = more pressure!
  8. Combined Gas Law is given on Table T **If a variable is constant, you can cross it out of the formula
    • Temperature is constant: P<em>1V</em>1=P<em>2V</em>2P<em>1V</em>1 = P<em>2V</em>2
    • Pressure is constant: V<em>1/T</em>1=V<em>2/T</em>2V<em>1/T</em>1 = V<em>2/T</em>2
    • Volume is constant: P<em>1/T</em>1=P<em>2/T</em>2P<em>1/T</em>1 = P<em>2/T</em>2
  9. STP (Standard Temperature and Pressure) listed on table A
    • Standard Temperature: 273 K or 0oC (must use 273 K for Gas Law Math Problems)
    • Standard Pressure: 1 atm or 101.3 kPa (pick the unit that matches when doing a Gas Law Problem)
  10. Gas Law Relationships:
    • As pressure increase, volume decrease (INVERSE relationship): less volume (less space for gases) = more collisions against the walls of the container = more pressure
    • As temperature increase, volume increase (DIRECT relationship): think hot air balloon. If you heat up a gas, it will spread out further, more volume.
    • As temperature increases, pressure increase (DIRECT relationship): increase temperature, increase average kinetic energy, increase movement, speed and collisions of particles against the walls of the container = more pressure. *If temperature is involved, there is a direct relationship.

Unit 5: Atomic Structure

  1. There are 3 subatomic particles in an atom
    • Proton:
      • Charge: +
      • Location: Nucleus
      • Mass: 1 u
    • Neutron:
      • Charge: Neutral/No Charge
      • Location: Nucleus
      • Mass: 1 u
    • Electron:
      • Charge: -
      • Location: Orbitals, energy levels surrounding the nucleus
      • Mass: ~0 u
        • (very tiny, 1/1836 u DOES NOT contribute to the mass of an atom)
    • Note about mass: Mass of protons and neutrons are equal (both 1 u).
    • Note about charge: Proton (+1) and electron (–1) have equal but opposite charges = “same magnitude but opposite charge”
  2. TABLE O also shows symbols, mass and charges of these 3 subatomic particles (electrons are represented as beta particle in table O), but it also shows other particles that will be covered in Nuclear Chem.
  3. Important Ways to Figure out Mass number, Number of Protons, Neutrons, Electrons:
    • Atomic Number = # of protons of an atom (use periodic table to look it up. Periodic table is organized by increasing atomic number) **All atoms of the same element have the same atomic number and same number of protons. Ex. All Carbon atoms have 6 protons because the atomic number of Carbon is 6.
    • Nuclear Charge = charge of the nucleus = number of protons in an atom. Ex. Carbon has a nuclear charge of +6 *Although neutrons are in the nucleus also, they have no charge, so they do not affect the nuclear charge.
    • Mass Number = Mass of ONE particular atom of that element = # Protons + # Neutrons **ATOMIC Mass is NOT the same as the mass number. Atomic mass = weighted average mass of ALL naturally occurring isotopes of an element (formula below, after definition of isotope)
    • # Neutrons = Mass number – # Protons
    • # Protons = # Electrons in an atom (because atoms are neutral)
    • Ion: an atom that has lost or gained electrons and has a charge.
      • Since electrons are – : Gain Electrons, become negatively charged. Lose electrons become positively charged. Ex. 1: When Al atom (13 protons, 13 electrons) loses 3 electrons → Al+3 ion (still have 13 protons, but now has 10 electrons) Ex. 2: When S atom (16 protons, 16 electrons) gains 2 electrons → S–2 ion (still have 16 protons but now has 18 electrons)
      • Charge of Ion = # protons – # electrons Ex. 9 protons, 10 electrons → -1 charged ion. (This is F– ion, fluoride)
      • Selected oxidation states on the periodic table show what charges metal ions can become. But when nonmetals become ions, the FIRST charge listed under the selected oxidation state is the charge they typically become. The ones below the first charge for nonmetals are possible oxidation numbers (covered in electrochemistry unit)
  4. Isotopes: atoms of the same element with the same number of protons, but different number of neutrons (and therefore different mass number)
  5. ATOMIC Mass = weighted average Atomic mass of ALL naturally occurring isotopes of an element = (mass of isotope 1)(% in decimal form) + (mass of isotope 2)(% in decimal form) + … (keep going until you included all of the isotopes) *Decimal form = percent / 100 Percent Abundance: how many percent of all atoms of an element is that particular isotope Ex. 78.99% of all Mg atoms in the world has a mass of 23.985 u
    • May also see: (massofisotope1)((mass of isotope 1)(%~NOT~in~decimal~form) + (mass of isotope 2)(%~NOT~in~decimal~form) \over 100
  6. Isotope Notation: Includes…
    • Element Symbol
    • Top left number of isotope notation = Mass number = protons + neutrons
    • Bottom left number = atomic number = # of protons
    • IF there is a charge, include on top right.
  7. Other notations: (C-14 or Carbon-14) number after an element represents mass number.
  8. History of Atom: the order is from least to most complex info about the atom
    • Hard, indivisible sphere
    • Atoms have small, negatively charged particles
    • The center of an atom is a small, dense, + charged nucleus
    • Electrons are located in orbitals & have wavelike properties.
  9. CATHODE RAY TUBE Experiment (Thompson) led to discovery of electrons (negatively charged particle in atoms). Plum pudding model: sea of + charge with electrons spread throughout.
  10. GOLD FOIL Experiment (Rutherford): alpha particles aimed at gold foil Conclusion: Atoms are mostly empty space with a small, dense, positively charged nucleus.
  11. Bohr Model: atom has a nucleus with electrons in energy levels/electron shells surrounding it
  12. Wave-mechanical model (electron cloud model, modern theory of the atom): orbital (electron cloud) is the most probable location of electrons. Electrons have wave-like properties.
  13. Electron Configuration: Shows location of electrons in their energy levels/shells. Ground State Electron Configuration is shown on the periodic table. Example: Na has a ground state electron configuration of 2-8-1 (2 electrons in first energy level, 8 electrons in second energy level, 1 electron in third energy level **Lower energy levels has less energy than higher energy levels. 1st energy level has less energy than 2nd, which has less energy than the 3rd and so on.
  14. Excited State: Electrons gain energy and jump/move from lower energy level to higher energy level. Ex. In Na, 1 electron from the second energy level can jump to the third energy level. The electron configuration changes from 2-8-1 to 2-7-2. Note that the number of electrons stays the same, just the location of the electrons changes.
  15. When excited electrons release energy and return (move back down) from higher to lower energy levels, colored LIGHT is produced. This light can be used to produce an element’s bright line spectrum.
  16. Every element has its own unique bright line spectrum and gives off unique colors. It’s like a fingerprint for an element.
  17. Bright line spectra can be used to identify elements from another and to see if certain elements are in a mixture.
    • If ALL of the spectral lines of an element are in the mixture, that element is in the mixture Ex. Lithium and Strontium are in the mixture because ALL of their lines are also in the mixture. Cadmium is NOT in the mixture because some of its lines (like the first two on the left) are NOT in the mixture.
  18. When atoms lose or gain electrons to become ions, electrons are lost or added from its LAST energy level. Ex. Na atom: 2-8-1 electron configuration, Na+ ion (lost one electron): 2-8 Ex. S atom: 2-8-6 electron configuration, S2– ion (gained 2 electrons): 2-8-8
  19. Valence Electrons = Electrons in the LAST (or outermost) energy level of an atom
    • IN GENERAL: In all chemical reactions, there is conservation of mass (matter, atoms), energy AND charge BUT…
    • Nuclear reactions are the only exception: mass/matter is converted into a lot of energy.

Unit 6: Nuclear

  1. Atoms that have large nuclei (atomic number greater than 83) and some atoms that have more neutrons than protons (ratio of neutrons to protons greater than 1:1) are radioactive.
  2. Radioactive = unstable elements that break down into more stable ones through releasing radiation and radioactive particles
  3. Radioisotope or nuclide refers to an isotope of an element that is radioactive. A list of selected radioisotopes can be found on table N
  4. Table N also tells you the DECAY MODE for radioisotopes. The decay mode = what particle (alpha, beta or positron) the radioisotope releases as a product as it breaks down (decay)
  5. 3 types of NUCLEAR Reactions
    • Transmutation: Atoms are being transformed into a different new element. “Mutation” Examples:
    • Fission: Nucleus of one atom is being split apart into 2 or more smaller nuclei. A neutron may bombard (or crash into) a big nucleus to do this. Example:
    • Fusion: Two or more smaller nuclear are FUSING (combining) to form a larger nucleus. Example:
  6. Nuclear reactions convert MASS into a LOT of ENERGY!! (Breaks the law of conservation of mass!)
  7. Table O shows different symbols and particles used in nuclear chemistry. Use the notation in filling in missing particles in nuclear reactions. Remember: mass is on the top left of the notation, atomic number or charge is on the bottom left.
  8. Strongest to weakest PENETRATING Power: Gamma, beta, alpha
  9. Strongest to weakest IONIZING Power: Alpha, beta, gamma (opposite order of penetrating power)
  10. When determining missing reactants or products:
    • Mass numbers (top left corner of notation) on the reactant side add up to equal the mass numbers on the product side
    • Atomic numbers/charge (bottom left corner of notation) on the reactant side add up to equal the atomic numbers/charge of the product side *If a and b are true, then the nuclear reaction is correctly balanced
    • Use the atomic number of the missing particle to look up what element symbol it is on the periodic table
    • Use table N to look up any decay modes needed for a radioisotope (alpha, beta, positron)
    • Use table O to look up any symbol notations for any of the decay mode particles (alpha, beta, positron)
  11. Benefits and Risks of Uses of Radioisotopes
    • Carbon-14 (14C^{14}C): used to date fossils (how old previously living organisms are).
    • Uranium-238 (238U^{238}U): used to date rocks and other geological formations
    • P-31 and C-14: used as tracers (tracks a path of a material in a system)
    • Iodine-131 (131I^{131}I): used to diagnose and treat thyroid cancer and conditions
    • Technetium-99 (99Tc^{99}Tc): used to treat and diagnose cancers.
    • Cobalt-60 (60Co^{60}Co): used to treat (kill) cancer cells
    • Cobalt-60 (60Co^{60}Co) & Cesium-137 (137Cs^{137}Cs): used to kill bacteria
    • Uranium is also used in nuclear power plants
    • Biological Exposure (from radiation and toxic waste products) can cause damage and mutation to healthy human cells
    • Disposal and storage of nuclear waste products is also a challenge
  12. Half Life = amount of time it takes for half the atoms of a given sample to decay (how long does it take for half of a radioactive material to break down into stable isotopes). The half life for radioisotopes are listed in Table N (be careful of units given for time). *Sometimes the half life is also given within the problem for a Regents Question. Read carefully!
    • Tips for solving half-life problems
      • To solve for: The fraction remaining of a radioactive isotope
        • You will need: A half life, A length of time
        • Steps to follow: Find how many half lives happened (total time passed ÷ half-life), Divide 1 by 2 until you reach how many half lives happened.
      • To solve for: The amount remaining of a radioactive isotope
        • You will need: A half life, A length of time, The original mass
        • Steps to follow: Find how many half lives happened (total time passed ÷ half-life), Divide original amount by 2 until you reach how many half lives happened.
      • To solve for: The original amount of a radioactive isotope
        • You will need: A half life, A length of time, The final mass
        • Steps to follow: Find how many half lives happened (total time passed ÷ half-life), Multiply final amount by 2 until you reach how many half lives happened.
      • To solve for: The half-life of a radioactive isotope
        • You will need: A length of time, The original mass, The final mass
        • Steps to follow: Find how many half lives happened (count how many times you need to divide the original amount by 2 to get the remaining/final amount), Divide the total time it took to decay by how many half-lives happened
      • To solve for: The time required for a radioactive isotope to decay
        • You will need: A half life, The original mass, The final mass
        • Steps to follow: Find how many half lives happened, Multiply the number of half lives happened by the half-life of the radioactive isotope (given in the problem OR on Table N)

Unit 7: Periodic Table

  1. The periodic table is arranged by increasing atomic number, starting with H as atomic number 1. As you go left to right and then down the table, the atomic number increases.
  2. Metals are on the left side of the staircase (except H, hydrogen is a nonmetal). Nonmetals are on the right side of the staircase. Metalloids (a.k.a. semi-metals, have both metal and nonmetal properties) are touching the staircase (except for Al, aluminum is a metal).
    • Metals LOSE electrons → + charged ions. After metals lose electrons, they lose an electron shell/energy level and become SMALLER.
    • Nonmetals GAIN electrons → – charged ions. After nonmetals gain electrons, they become BIGGER because they have more – electrons than + protons in the nucleus. Electrons “overpower” the pull from the + protons in the nucleus.
  3. Horizontal Row (left to right) = Period (there are 7 periods)
    • The period that an element is in also matches how many electron energy level/shells that atoms of that element has. Ex. Na is in period 3 and has 3 energy levels with the following electron configuration: 2-8-1
  4. Vertical Column (up and down) = Groups (there are 18 groups)
    • Elements in the same group (column) have similar chemical properties because they have the same number of valence electrons.
      • Ex. Group 18 (noble gases) all have 8 valence electrons (except He, which has 2). A chemical property that they all share is that they are all stable and nonreactive. This is because they all have a full outermost energy level/shell of electrons.
        • Atoms of other elements often gain or lose electrons to have a full outermost electron energy level/shell, matching the electron configuration of noble gases (more on this in Bonding Unit)
      • Group 1: Alkali Metals, Group 2: Alkaline Earth Metals, Group 3-12: Transition Metals (can form different colors), Group 17: Halogens
  5. There are general trends you can see in the periodic table.
    • Atomic Radius: The length between the center of an atom to its outermost energy level.
      • Going across a period, in order of increasing atomic number (left to right), atomic radius DECREASES.
      • Going down a group, in order of increasing atomic number (top to bottom), atomic radius INCREASES.
    • Electronegativity: Atom’s strength of attraction for electrons in a chemical bond. On a scale of 0-4. Closer to 0 = weaker electron attraction. Closer to 4 = Stronger electron attraction.
      • Going across a period, in order of increasing atomic number (left to right), electronegativity INCREASES.
      • Going down a group, in order of increasing atomic number (top to bottom), electronegativity DECREASES.
    • First Ionization Energy: the amount of energy needed to remove the first outermost electron in an atom
      • Going across a period, in order of increasing atomic number (left to right), first ionization energy INCREASES.
      • Going down a group, in order of increasing atomic number (top to bottom), first ionization energy DECREASES.
      • You do not need to memorize the trends because all of these trends can be figured out by looking up values on Table S. Pick 3 or 4 elements in the same period or group and compare the values found on Table S and ask yourself if they are increasing or decreasing?

Unit 8: Bonding

  1. Valence Electrons are the outermost in an atom (the electrons in the outermost or last electron shell or the highest energy level). Example: Na’s electron configuration is 2-8-1; Na has 1 valence electron.
  2. In general, elements in the same group (column) have similar chemical properties because they have the same number of valence electrons. Example: Mg and Ca have similar chemical properties because they both have 2 valence electrons.
  3. Atoms form bonds so that they can become stable.
    • Stable = having 8 valence electrons (an octet rule)
    • Notice that the Noble Gases in Group 18 all have 8 valence electrons already, and so they are all stable without having to bond
    • EXCEPTION: Hydrogen and Helium only need 2 valence electrons to be stable.
  4. Anion (– ion): An atom (usually a nonmetal) that gained electrons, and become negatively charged
  5. Cation (+ ion): An atom (usually a metal) that lost electrons, and become positively charged
  6. There are 3 types of bonds: ionic, covalent and metallic. The type of bond depends on what type of atoms are involved and how they compete for valence electrons (see below)
  7. Ionic compound = Metal ion + Nonmetal ion bonded together
    • A metal TRANSFERS or DONATES its VALENCE electrons to a nonmetal.
  8. Covalent compound (also known as molecules or molecular compounds) = 2 nonmetals SHARE their valence electrons
    • The nonmetals share their valence electrons because neither are electronegative enough (attract electrons) to take away/gain valence electrons
  9. Metallic bonds = only between atoms of METAL elements
    • Valence electrons are mobile or moving around atoms
    • ^This is why metals are malleable, and good conductors of electricity
  10. Metallic bonding is described as “mobile sea of electrons”
  11. Electronegativity: Measures how strongly an atom attracts electrons
    • Nonmetals are highly electronegative. That is why they receive electrons.
    • Metals are not very electronegative. This is why they give up and donate their valence electrons.
  12. Naming an ionic compound (IUPAC Rules):
    • Step 1: Naming the positive cation
      • If the ion has ONE positive charge listed→ the ion name is the same as the element name
      • Example: Na+ → sodium
      • If the ion has TWO or MORE positive charges listed: the charge is listed after the ion name in parentheses
      • Example: Cu → Copper (I) and Copper (II)
    • Step 2: Naming the nonmetal anion:
      • -First name the element of the nonmetal
      • -Then take off the ending of the nonmetal element name and then add -ide
      • Ex. Cl– = chlorine → chloride
    • Step 3: Put it all together! Metal/cation always comes first. **Check Table E for Polyatomic Ions! The names and chemical formulas (with charge!) are listed on Table E. For example: NH4 + is called ammonium. NO3 – is called nitrate. NH4NO3 = ammonium nitrate NH4Cl = ammonium chloride NaNO3 = sodium nitrate
  13. Writing chemical formulas for ionic compounds:
    • Find the element (periodic table) or polyatomic ion (table E) and their charge
    • CRISS-CROSS CHARGES without the sign (charges are found on periodic table or table E or given to you in roman numerals for metals that have multiple charges)
    • REDUCE subscripts to lowest possible ratio if possible Ex1.