Chemistry Regents Review

Table of Contents Review #1-2: Introduction/Matter and Energy pg 3 Review #3: Atomic Concepts.pg pg 12 6 Review #4: Periodic Table. pg 17 Review #5: Bonding and Shapes...pg pg 19 22 Review #6: Polarity and Intermolecular Forces...pg 26 Review #7: Chemical Reactions.. pg 30 Review #8: The Mole and Stoichiometry...pg 32 Review #9: Equilibrium and Solutions... pg 35 Review #10: Energy and Kinetics.. pg 39 Review #11: Acids and Bases.. pg 48 45 Review #12: Redox...pg 52 49 Review #13: Nuclear Chemistry.pg 57 54 Review #14: Organic Chemistry. pg 59 Review #15: Reference Tables...pg 63 Review #16: Big Ideas.... pg 66 2

Review #1-2 Introduction/matter 403.00 8.100 x 10 5 0.0010 3,500,000 When calculating: Multiplying and dividing: round to the same number of significant digits as the factor with the least number of significant digits Adding and Subtracting: round to the same decimal place as the factor with the least number of decimal places. Metric System: 3

Some important formulas: Matter and Energy Matter: has and The Law of Conservation of Mass: Energy: capacity to do and produce Key Idea: Temperature is a measurement of. All three phases differ by the amount of movement among particles. Solids Liquids Gases 4

Key Idea: Kinetic Molecular Theory is a model that is used to explain the behavior of matter. It describes relationships among pressure, volume, temperature, velocity, frequency, and force of. Kinetic Molecular Theory describes an gas: Gas particles are in,, motion Gas particles are separated by distances relative to size. The volume of particles is considered. Gas particles have no forces between themselves. This occurs best at temperatures and pressures. Collisions between the particles and the of their containers create and are perfectly (transfer of energy). In reality: Combined Gas Law (note: temperature in Kelvin) Avogradro's Law Same volumes have same number of particles Graham's Law Lighter gases diffuse/effuse faster Dalton's Law Total pressure is the sum of all the partial pressures 5

Changes in Phase: Key Idea: A heating curve (which shows an process) demonstrates that during a phase change, the, or temperature, does not change. Melting Vaporization Sublimation Freezing Condensation Deposition Physical Change Chemical Change 6

Element: Approx 118 total Found on the Periodic Table Named after people, places, Latin origins, etc Cannot be broken down physically or chemically Compound: Can be chemically separated Fixed ratio of components Properties of compound are different than individual components Written with formulas Mixture: Homogeneous: Heterogeneous: Not pure substances Physically combined, can be physically separated Variable ratio Individual components retain properties Key Idea: Elements and compounds are considered. 7

Separating Mixtures: Filtration: Used to separate a solid from a liquid. WILL NOT SEPARATE A DISSOLVED SUBSTANCE! Distillation: Separates two liquids based on different boiling points. Crystallization: Separated a dissolved substance (aq) from its solvent Chromatography: Separates based on polarity and solubility Practice Regents Questions 1. 5. 2. 6. 3. 7. 8

4. 8. 9. 13. 10. 14. 11. 15. 12. 16. 9

17. 18. Particle of CO 2: 10

Review #3 Atomic Concepts Define: Atom: the basic unit of a chemical element Atomic Number: number of protons in an atom Mass Number: total mass of the atom (protons + neutrons) Atomic Mass: weighted average of naturally occurring isotopes Isotope: same element (protons), different mass (number of neutrons) 12

Valence Electron: outermost electron(s) Orbital: area of probable electron location Principal Energy Level: shell of electrons - 7 total in atom Excited State: electron has jumped to a higher energy level within atom Ground State: electrons are in lowest energy configuration (closest to nucleus) Bright Line Spectrum: produced when electrons return to ground state Electron Configurations: denote how electrons are distributed in P.E.L.s Models of the Atom Dalton s Atomic Theory: 1. All elements are composed of 2. All atoms of a given element are 3. Atoms of different elements are J.J. Thomson: Rutherford: Cathode ray tube experiments Discovered the electron Model of the atom: Gold Foil Experiment: Shot alpha particles at thin gold foil, most went straight through, some deflected and other particles shot straight back. 1. 2. 13

Bohr: Planetary model of the atom Electrons are in defined around the nucleus. Wave-Mechanical Model: In the wave mechanical model ( ), the electrons are in, which are defined as regions of most electron location ( ). The wave- mechanical model describes the dual nature of the electron as it has properties of both a and a. Subatomic Particles & Mass 3 subatomic particles make-up the atom: 1. located in charge of 2. located in charge of 3. located in charge of 1 amu = = mass of 1 proton or 1 neutron Each atom has its own which equals the number of Key idea: The mass of each proton and each neutron is approximately equal to one atomic mass unit. An electron is much less massive than a proton or neutron. Electrons have a charge. In a neutral atom the number of electrons = the number of. Key idea: The Atomic Number is equal to the number of in the nucleus of an atom. The mass number is calculated by adding the number of +. 14

Key idea: The average atomic mass of an element is the of the masses of its occurring. Electron Configuration & Bright Line Spectra Key Idea: The outermost electrons in an atom are called the electrons. The number of valence electrons affects the chemical properties of an element. Key idea: * Each electron in an atom has its own distinct amount of energy. * When an electron in an atom gains a specific amount of energy, the electron is at a higher energy state known as the. * When an electron returns from a higher energy to a lower energy state ( ) a specific amount of energy is released in the form of. The emitted energy can be used to identify an element ( ) Practice Regents Questions 1. 7. 2. 8. 3. 9. 15

4. 10. 5. 11. 6. 12. 16

Review #4 Periodic Table 17

Directions: 1. On the above table label the following: a. Group, Period, Noble Gases, Halogens, Transition Metals, Alkaline Earth Metals, Alkali Metals. b. Place an X over the diatomic elements c. Circle the elements that are liquids d. List all of the gases: e. Label the ions groups 1, 2, 13-18 form above the group. Define: Atomic Radius: distance from the nucleus to the outermost electron Metallic Character: How easily an element loses an electron (reactivity) Ionization Energy: energy required to remove an electron Electronegativity: attraction/affinity for electrons in a bond Facts and Trends of The Periodic Table 1. The elements on the periodic table are arranged in order of: 2. As you go across a period the atomic radius: 3. As you go down a group the atomic radius: 4. A positive ion has a radius than the original atom 5. A negative ion has a radius than the original atom. 6. As you go across a period the ionization energy: 7. As you go down a group the ionization energy: 18

8. As you go across a period the electronegativity: 9. As you go down a group the electronegativity: 10. As you go across a period the metallic character of the elements: 11. As you go down a group the metallic character of the elements: 12. Name the two groups that are so chemically active that they occur naturally only in compounds. 13. Name the group that contains elements in three phases of matter at room temp. 14. The majority of elements on the periodic table are in what phase at room temperature? 15. The seven semimetals, found on the step line, are Properties of Metals and Nonmetals: Location on table Appearance Physical properties Phases Metals Nonmetals Only liquid example Conductivity Ionization 19

Electronegativity Ions formed Diatomic Elements: Key Idea: The periodic table is arranged by. The properties of elements are, or repeat. Valence electrons are the electrons. Key Idea: Because each member of a has the same number of valence electrons, each member of the group has. Key Idea: An is an atom with a charge. Metals form ions by electrons, while nonmetals form ions by electrons. Practice Regents Questions 1. 8. 2. 9. 20

3. 10. 4. 11. 5. 12. 6. 13. 7. 14. 15. 17. 16. 21

Review #5 Bonding and Shapes Lewis Dot Diagrams: depict valence electrons Li Ca O Br -1 Mg +2 H 2 O NH 3 MgCl 2 CO 2 CH 4 22

Define: Ionic Bond: Example Covalent Bond: Example Octet Rule: Key Idea: Atoms will gain, lose, or share to become. Stability is achieved with valence electrons. (Remember first shell has valence electrons) Naming Ionic Compounds: name the and. (Table E has ions). Add an ending if the anion is not from table E. Add roman numerals, which indicate the, onto the following elements: Naming Covalent Compounds: use the listed below. You can drop the prefix on the atom only. Writing Formulas: ensure that the balance. 23

Criss cross method: Types of Solids: Crystalline Solid: Amorphous Solid: Allotrope: Practice Regents Questions 1. 6. 24

2. 7. 3. 8. 4. 9. 5. 10. 11. 14. 25

12. 15. 13. 16. Review #6 Polarity and Intermolecular Forces 26

Polar Bond: Example Nonpolar Bond: Example Key Idea: The larger the in electronegativity, the more the bond is. Shapes and Polarity: Bent: Pyramidal: Linear: Tetrahedral: Trigonal Planar: Intermolecular Forces: Stronger Weaker 27

Key Idea: Intermolecular forces affect physical properties. The stronger a substance s intermolecular substances, the its boiling point, weaker its pressure, and more tension it experiences. Vapor Pressure: 28

Practice Regents Questions 1. 7. 2. 8. 3. 9. 4. 10. 5. 11. 6. 12. 29

Review #7 Chemical Reactions Synthesis Decomposition Single Replacement Double Replacement Single Replacement Occurs: Double Replacement Occurs: Balancing: remember, add only. Do the following reactions occur? Fe + NaCl Na + FeCl 3 Cu(OH) 3 + NaCl CuCl 3 + NaOH Key Idea: The law of conservation of mass states that the number of atoms on each side of the equation must be the same. 31

Practice Regents Questions 1. 5. 2. 6. 3. 7. 4. 8. 32

Review #8 The Mole and Stoichiometry Gram Formula Mass: Mathematical Calculations: % Composition Empirical/Molecular Formulas The Mole: Two most likely conversions: Key Idea: Converting between any quantity and another is best accomplished with dimensional analysis. Ensure that your units cancel out. 33

Practice Regents Questions 1. 6. 2. 7. 3. 8. 4. 9. 5. 10. 34

Chemical Equilibrium: Review #9 Equilibrium and Solutions Key Idea: When a system is at equilibrium, reaction are equal. Concentrations of reactants and products are, but not necessarily equal. Le Chatelier s Principle: Stresses on a system include a change in: 1. 2. 3. Note: The addition of a catalyst the rate of the forward and reverse reactions. However, there is NO SHIFT IN EQUILIBRIUM. Equilibrium can also be in. The processes of and are opposite reactions; equilibriums also exist. 36

Soluble: ability to dissolve Insoluble: unable to dissolve Solute: is dissolved Solvent: does the dissolving Solution: solute + solvent = homogeneous mixture Unsaturated: solvent can hold more solute Saturated: solvent cannot hold more solute Supersaturated: solvent is holding more solute then it can hold The solubility of a substance is affected by the of the solute/solvent,, and (for gases only). The rate, or how quickly a solute dissolves, is affected by,, and. Like Dissolves Like When ionic substances dissolve, the become surrounded by water molecules. The oxygen of surrounding water molecules are attracted to ions, while the hydrogens of surrounding water molecules are attracted to ions. Na + Cl - Colligative Property: Key Idea: Boiling point is, freezing point is with the addition of solutes. and pressure also changes with the addition of solutes. 37

Practice Regents Questions 1. 6. 2. 7. 3. 8. 4. 9. 5. 10. 38

Review #10 Energy and Kinetics Chemical Kinetics: Mechanics and rates of reactions Particles must collide to react Key Idea: Collision theory states that for a reaction to occur, a must occur between two particles. The particles must have the correct and. Reaction rates are affected by: Energy: Capacity to do work/produce heat Always transferred/change form = conservation of energy warm cold 1. What is the number of joules of energy released when 25g of water are cooled from 20 C to 10 C? 2. What is the total number of kilojoules of heat needed to change 150g of ice to water at 0 C? 40

Exothermic Endothermic Catalyst 41

Heat of reaction: ΔH = Entropy: ΔS = Table I -ΔH = Exothermic 1. created 2. increased 3. more created 4. from solids and liquids Universal Tendencies: 2. If 2.0 moles of nitrogen are consumed in this reaction, how much heat is absorbed? 1. Explain, in terms of S and H, why this reaction is spontaneous. 42

Practice Regents Questions 1. 5. 2. 6. 3. 7. 4. 8. 43

9. 12. 10. 13. 11. 14. 44

Review #11 Acids and Bases Define: Arrhenius Acid: donates H + Arrhenius Base: donates OH - Bronsted-Lowry Acid: donates H + Bronsted-Lowry Base: accepts H + Neutralization: acid + base H 2 O + salt Indicator: changes color at different ph levels Ionize/Dissociate: molecules break into ions ph: measures amount of H + ions Titration: lab technique to measure molarity of acid/base 46

Electrolytes: Key Idea:,, and are all electrolytes due to their. Titration: M A V A = M B V B Remember! MA = molarity of H + ions MB = molarity of OH - ions [H2SO4] =.3 M [H + ] = [Ca(OH)2] = 1.4 M [OH - ] = [H 3PO4] =.15 M [H + ] = [Mg(OH)2] = [OH - ] = 1.5 M Bronsted-Lowry: 47

Practice Regents Questions 1. 8. 2. 9. 3. 10. 4. 11. 5. 12. 6. 13. 48

7. 14. 49

Review #12 Redox GER Oxidation: Reduction: When an atom is oxidized (LOSE ELECTRONS) it acts as a because that electron is GAINED by another atom. When atom is reduced (GAINS ELECTRONS) it acts as an because the electron is LOST by another atom. 50

Voltaic Cell Electrolytic Cell. 1. Cu + 2AgNO3 Cu(NO3)2 + 2Ag Oxidized: Reduced: Oxidizing Agent: Reducing Agent: Oxidation ½ Reaction: Reduction ½ Reaction: 2. 2Mg + Pb(NO3)4 2Mg(NO3)2 + Pb Oxidized: Reduced: Oxidizing Agent: Reducing Agent: Oxidation ½ Reaction: Reduction ½ Reaction: Key Idea: A voltaic cell is a redox reaction. Oxidation occurs at the ( ) and reduction at the ( ) 51

1. Label the anode 2. Label the cathode 3. Draw in the flow of electrons 4. What is the purpose of the salt bridge? 5. Which electrode is gaining in mass? 6. Which electrode is losing mass? 7. Write the half reaction that takes place in half- cell 2: 8. Write the half reaction that takes place in half- cell 1: 9. Write the overall balanced redox reaction: 10. What type of energy conversion takes place in the cell shown? Key Idea: An electrolytic cell is a reaction that requires an energy source to occur. and are examples. 1. Label the anode 2. Label the cathode 3. Draw in the flow of electrons 4. Which electrode is gaining in mass? 5. Which electrode is losing mass? 6. What is the purpose of the battery? 52

Practice Regents Questions 1. 5. 2. 6. 3. 7. 4. 8. 53

Review #13 Nuclear Chemistry Nuclear Reactions: Key Idea: A nucleus that is will spontaneously emit radioactive particles. An atom s nucleus is unstable if: 1. The is unbalanced. 2. The nucleus has more than and is simply too large to hold itself together. 239 Pu 94 90 Sr 38 53 Fe 26 55

Half Life The amount of time it takes for half the sample to decay. If I have a 12.5 gram sample of N-16, how much time has elapsed if the original sample was 100.0 grams? Tips for solving half-life problems: Never divide grams by grams. You can only multiply or divide grams by 2. How many grams of a 60 gram sample of Ra-226 will be left after 3198 years? # of H.L. = total time elapsed Length of H.L. Use Table N for specific times. 56

Artifical Transmutation Fission Fusion Uses & Dangers of Radioisotopes Radioactive isotopes have many beneficial uses. Several different radioactive isotopes are used in medicine and industrial chemistry. Tracing biological processes Detecting diseases Treating cancers/diseases Nuclear power Iodine 131: Cobalt-60 Carbon-14 U-238 to Pb-206 57

Practice Regents Questions 1. 8. 2. 9. 3. 10. 4. 11. 5. 12. 6. 13. 7. 58

Review #14 Organic Chemistry Alkanes: Single bonds, Saturated Draw and name: C3H8 Draw: Butane Name the structure below: Alkenes: Double bonds, Unsaturated Draw and name: C3H6 Draw: 1-butene Name the structure below: Alkynes: Triple bonds, Unsaturated Draw and name: C3H4 Draw: 2-pentyne Name the structure below: 60

Name the following: 2-bromobutane 2,3-butandiol Pentanoic acid Ethyl pentanoate 61

Organic Reactions: Combustion Hydrocarbon + O2 Fermentation Sugar Substitution (remember: involves hydrocarbons) Addition (remember: involves hydrocarbons) Esterfication Saponification Ester + Base Alcohol + Soap Polymerization Addition: Condensation: 62

1. Put a circle around the alkanes, a square around the alkenes, and a triangle around the alkynes. Butane 3-propyne 2-methyl-1- hexene 2-octene 2,3-dimethyl octane C 4 H 6 C 3 H 4 C 7 H 16 C 2 H 4 CH 4 2. Which of the hydrocarbon series are considered saturated? 3. What is wrong with the following diagram? 4. Draw the isomers of hexane try to do it without your notes - there are 5 total 5. Put a circle around the compounds that could be used in an addition reaction, and a square which could be used in a substitution reaction. C 4 H 6 C 3 H 4 C 7 H 16 C 2 H 4 CH 4 63

Review #15 Reference Tables Directions: Find each of the following pieces of information in the Reference Tables. 1. If I start with 100 grams of a sample and after 58.2 years I have 25 grams, what atom did I start with? 2. What is the name of the anion in KMnO4? What is its charge? 3. If a solution turns yellow with the addition of methyl orange and bromthymol blue, what range could the ph be? 4. How much NaNO3 can be dissolved in 200 grams of water at 45 C? 5. If I overhear a discussion about the pascal, what quantity is being discussed? 6. What classification of hydrocarbon would C2H4 be? 7. What liquid boils at 80 C under normal pressure (101.3 kpa)? 8. This element has an electronegativity of 2.6 and a boiling point of 718 K. 9. What is the standard temperature of water? 10. Is NaOH soluble under normal conditions? 11. Of the following compounds, the decomposition of which one is exothermic: KNO3, NH4Cl, NaOH, or NaCl. 12. What is the formula for acetic acid? 13. Can this reaction occur: Ag + CuNO3?? 14. What emitted particle has a mass of 4 amu? 64

15. CH3 CH2 O CH3 would be what class of organic compound? 16. What element has an electron configuration of 2-8-1? 17. 1.5 x 10-9 meters can also be expressed as 1.5 meters. 18. What is the weighted average of all the naturally occurring isotopes of oxygen? 19. What is the heat of fusion of water? Using table T, solve the following problems: 1. Give the parts per million of solute for a solution containing 25g of sodium chloride in 200g of water. 2. If the accepted value for the mass of an object is 10.3g and a student found that the mass was 10.1g, what is the student s percent error? 3. If a peanut is burned in a calorimeter containing 50g of water, and the water temperature changes from 45 0 C to 57 0 C, how many joules of energy were released by the peanut? 4. Convert: 40 grams of NaCl into moles of sodium chloride. 5. What mass of iron has the dimensions 3.20 cm x 6.50 cm x 0.30 cm? 6. Convert: 384 K into C. 65

7. A sample of gas is in a container at 30 C and 100 ml under a pressure of 40 kpa. If the container is put under STP conditions, what is the new volume of the container? 8. If I dissolve 5 moles of NaCl into 2.50 L of water, what is the molarity of the resulting solution? 9. 28 ml of 0.500 M HCl is required to neutralize 50. ml of NaOH. What is the molarity of the sodium hydroxide? a. How would your answer change if the base being neutralized was Ca(OH)2? 10. What is the percentage composition of oxygen in MgSO4? 66

Review #16 Big Ideas 1. Atomic # = # of Protons = # of electrons in an atom 2. Isotopes: Same element (same protons) different number of neutrons 3. How light is produced: An electron gains energy and jumps to a higher principle energy level (excited state). Then the electron falls back to the ground state emitting energy in the form of light. 4. Know Periodic Trends in atomic radius, ionization energy, and electronegativity. 5. Ideal Gas: High Temperature and Low Pressure 6. Heat is a form of energy. Heat travels from hot to cold. 7. When a chemical bond is formed energy is released. 8. When a chemical bond is broken energy is absorbed. 9. Temperature measures Average Kinetic Energy of molecules and atoms. 10. Higher boiling points = STRONGER intermolecular forces 11. If a substance sublimes at room temperature it has WEAK intermolecular forces. 12. Gases have the most entropy 13. A saturated solution is at equilibrium 14. Electrolytes: Ions in an aqueous solution that will conduct ELECTRICITY. 15. Electrolytes include: Acids, Bases, and Aqueous Salt solutions 16. Atoms are made up of mostly empty space. 17. Run away from an added stress in a reaction at equilibrium. 18. The addition of a catalyst to a reaction lowers the energy of activation. 19. Coefficients are used to balance an equation. They also represent the ratio of moles. 20. Standard Temperature Pressure (Table B). 21. Ionic Bond: Between a metal and a nonmetal 22. Covalent Bond: Between two nonmetals 23. Like polarity will dissolve in like polarity. (polar in polar, etc) 24. As temperature increases the solubility of a gas decreases. 25. Pressure affects the solubility of gases NOT solids. 26. Acids donate protons (H + ) 27. Bases accept protons (H + ) 28. Arrhenius Acids: Yield H + ions 29. Arrhenius Bases: Yield OH -1 ions 30. Fission: (Look for NEUTRONS) ( ) 31. Fusion: Combining of Hydrogen to make helium.( 1 H + 3 H 4 He + Energy) 32. All reactions show a conservation of MASS, CHARGE and ENERGY 33. LEO: Lose electrons oxidation (get more +) 34. GER: Gain electrons reduction (get more -) 67

35. An Ox: Anode Oxidation 36. Red Cat: Reduction Cathode 37. Alpha particles have a mass of 4 amu and a positive charge 38. Beta particles have no mass and a negative charge 39. Particle Diagrams: Elements Compounds Heterogeneous Mixture 40. Collision Theory: If you add more reactants there will be more effective collisions between the reactants causing the reaction to shift right and make more products. 41. H2 O2 F2 Br2 I2 N2 Cl2 : Diatomic elements 42. Valence Electrons: Are used in bonding and give elements their unique chemical properties 43. If you see the word TITRATION or NEUTRALIZATION use the formula: MA VA = MB VB 44. Mole Mole problems: set up a proportion using coefficients and the # of moles given 45. Mole Gram or Gram Mole conversions: Moles = Given Mass Gram formula mass 46. Hydrocarbon: Alkane, Alkene, Alkyne 47. Class of organic compounds and functional groups: Table R 48. Organic Reaction: Substitution Two reactants and two products 49. Organic Reaction: Addition Two reactants and one product 50. Organic Reaction: Esterification Organic Acid + Alcohol 51. Heating Curve for Water q = mc T q = mhf q = mhv 68