Chemistry Review. No work = No credit. Entire document must be complete for credit to be awarded.

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1 1 Unit 1 1. States of matter: Phase 1. Solid 2. Liquid 3. Gas Chemistry Review No work = No credit. Entire document must be complete for credit to be awarded. Definite Shape (Yes or No?) Definite Volume (Yes or No?) 2. Explain how can a color change be classified as a physical change? 3. Chemical change Signs of Chemical Change Examples Unit 2 4. How do you know you have produced a certain gas? (Hint: Laboratory test) a. Hydrogen b. Carbon dioxide 5. Classification of matter a. Pure Substance : element (cannot be decomposed) vs. compound (can be decomposed by chemical means) b. Mixtures homogenous (solutions, uniformed compositions, all the same) vs. heterogeneous (not uniformed) 1. Metric conversion ( K ing H enry D ied U nusually D rinking C hocolate M ilk) 2. Volume (amount of space occupied by an object) a. Graduated Cylinder (for liquids) 3. Density = mass/volume (remember the triangle) a. What is the density of a substance that has a mass of 22.5g and a volume of 5.0mL? What is this substance? (look on your reference sheet) b. What is the density of a 3.03cm x 10.cm x 2.5cm substance that has a mass of 50.0g? What is this substance? (look on your reference sheet)

2 2 Unit 3 1. Know the locations, charges, and relative mass of each of the particles found in the atom Proton Electron Neutron Charge Relative Mass Location 2. Atomic number = 3. Atomic mass = 238 U 4. Know how to identify the atomic number, mass number, number of electrons, protons, and neutrons of various atoms Electrons Protons Mass Number Atomic Number Neutrons 16 O 2 23 Na Isotopes Same element!!! Different Masses!!! Different Neutrons!!! They differ in the number of not (which always must stay the same as they indicate the atomic number!) Ex. A Z X Z means A means Ex. 20 F Atomic Number Mass Number 6. Atomic Mass Unit (amu) relative scale based on 7. Average Atomic Mass 8. Molar Mass conversions a. What is the molar mass of CO 2? b. How many grams are in 3.0 moles of H 2 SO 4? c. How many molecules are in 64 grams of O 2? d. How many moles are in 84.2 grams of CO 2?

3 3 e. How many moles are in 3.04x10 23 molecules of H 2? f. How many atoms are in 3.46 moles of carbon? g. How many grams are in 4.59x10 25 particles of NaCl? Unit 4 1. Electromagnetic radiation definition: 2. How are wavelength, frequency, and energy related? Longest wavelength Shortest frequency Least Energy Shortest wavelength Highest frequency Most Energy 3. Quanta:. 4. Each energy level corresponds to a on the periodic table. ** Highest occupied energy level corresponds to. Ex. Br 2 O 2 5. Ground State vs. Excited State 6. Hydrogen Line Emission Spectrum How is light emitted in reference to electrons? (Flame Test) 7. Questions using Bohr Model of the atom: What is the energy which is emitted when the electron falls from the 6 th energy level to the 3 rd? 8. Where are the s, p, d, and f blocks located on the periodic table? 9. Orbital Notation for bromine (this is arrows) 10. Electron configuration for bromine 11. Noble gas configuration for bromine 12. How many valence electrons do these have? What element is it? How many electrons will it gain/lose? a. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 b. [Ne] 3s 2 3p 4 ***To be stable the p and d orbitals must be ½ filled or complete filled.

4 4 13. Number of orbitals = n 2 Ex. How many orbitals can the 4 th energy level have? 14. Number of electrons = 2n 2 Ex. How many electrons can the 6 th energy level hold? Unit 5 1. Periods ( ) vs. Groups ( ) 2. The periodic law: elements in the same have similar, the same number of, and the same. 3. Cations are charged ions. They are formed by electrons. 4. Anions are charged ions. They are formed by electrons. 5. Draw the hill of oxidation numbers: 6. The periodic table is arranged by. 7. The periodic table is mostly. 8. Metals: a. Where? b. Gain or lose electrons? c. Positive or negative ions? d. Properties? 9. Non metals: a. Where? b. Gain or lose electrons? c. Positive or negative ions? d. Properties? 10. Metalloids: a. Where? b. Which elements? c. Properties? 11. Representative Elements (Main Group Elements) a. Where?

5 5 12. Alkali Metals a. Where? b. Gain or lose electrons? (How many?) c. Charge/Oxidation number? 13. Alkaline Earth Metals a. Where? b. Gain or lose electrons? (How many?) c. Charge/Oxidation number? 14. Transition Metals a. Where? b. Gain or lose electrons? (How many?) c. Charge/Oxidation number? d. What do you use when naming? 15. Halogens a. Where? b. Gain or lose electrons? (How many?) c. Charge? 16. Noble Gases a. Where? b. Gain or lose electrons? (How many?) 17. Reactivity: a. Which group of metals are the most reactive? b. Group that is not reactive or inert? 18. Valence Electrons correspond to the of the periodic table. 19. Atomic Radius trend a. Definition b. Left to Right: Increase or Decrease Reasoning: c. Top to Bottom: Increase or Decrease? Reasoning: d. Put the following in order from largest to smallest atomic radius Ca, K, Cu, Se?

6 6 20. Ionization Energy a. Definition b. Left to Right: Increase or Decrease Reasoning: c. Top to Bottom: Increase or Decrease? Reasoning: d. Put the following in order from lowest to highest ionization energy N, Bi, P, Sb? 21. Electronegativity a. Definition b. Left to Right: Increase or Decrease Reasoning: c. Top to Bottom: Increase or Decrease? Reasoning: d. Put the following in order from greatest to lowest electronegativity Mg, Ra, Be, Ca? 22. Ionic Radii a. Definition b. Cations: Smaller or Larger than uncharged atom? Reasoning: c. Anions: Smaller or Larger than uncharged atom? Reasoning: Unit 6 1. Valance electrons are responsible for the element s. 2. Why do atoms bond?. 3. Metallic bonds involve only. They are characterized by also called sea of. 4. Ionic compounds involve and. They are characterized by the of electrons. 5. Covalent molecules involve. They are characterized by the of electrons.

7 7 6. Type of bonding and electronegativity differences. Type Electronegative Difference Example Polar NonPolar Ionic 7. Diatomic molecules (magic 7 HOFBrINCl) a. Make sure you know the seven. b. All have bonding. 8. Octet rule states all atoms. 9. Electron Dot Notation (valence electrons shown as dots only!) a. C b. Cl 10. Lewis structure: Draw the following molecules. What shape are they? Are they polar/nonpolar? a. CCl 4 b. CO 2 c. NH Place these in order of increasing strength: Please these in order of increasing bond length: Hint: Draw the molecules. a. Single Bond example F 2 b. Double Bond example O 2 c. Triple Bond example N Molecular Geometry and VSEPR main shapes Shape # of Atoms Attached Example a. Linear b. Bent c. Trigonal pyramidal d. Tetrahedral

8 8 13. How do you know if the molecule is polar or nonpolar? a. Polar and lone pairs around center. b. Nonpolar and lone pairs around center. 14. Intermolecular Forces: Which elements does hydrogen need to be attached to in order to form a Hydrogen Bond? 15. Properties of Bonds Ionic Compounds Covalent Molecules Metallic Bonds Unit 7 If the formula is given: Write the name. If the name is given: Write the formula 1. Binary Compound: elements. 2 nd element ends in. 2. Types of Naming a. Ionic Bonding between and i. Write the formula for Sodium Chloride ii. Name Mg 3 N 2 b. Covalent Bonding between and (use ) i. Name CCl 4 ii. Write the formula for Dinitrogen pentoxide c. Transition metals between and (use ) i. Write the formula for Iron (II) oxide ii. Write the formula for Manganese (III) oxide d. Polyatomic Naming involves at least one polyatomic you have more than two capital letters look up all polyatomics in reference table!!) e. Polyatomic Ions end in or. Two exceptions on your reference table that end in ide are and. is the only positive polyatomic ion on your reference sheet and it is written first in formulas. i. Write the formula for Calcium hydroxide ii. Name FeSO 4 f. Acids Hydrogen is in front Still cross charges!!! i. Write the formula for Nitric Acid ii. iii. iv. Write the formula for Acetic Acid Write the formula for Hydrochloric Acid Write the formula for Sulfuric Acid

9 9 3. Percent composition by mass Part/Whole x100 a. Calculate the percent by mass of water in silver sulfate pentahydrate 4. Empirical Formula ( ) vs. Molecular Formula ( ) a. A compound contains 32.38% Na, 22.65% S, and 44.99% O. What is the empirical formula? b. A molecule has an empirical formula of CH 4 and a molar mass of 48, what is the molecular formula? Unit 8 1. Balance equations follow the Law of Conservation of Mass, which states. 2. Ionic equations are written to show what actually happens. a. Spectator Ions are ions which. b. Ionic equations show everything but. 3. Predict the products, balance and label what type of reaction each of the following are a. C 3 H 8 + O 2 type: b. NaCl type: c. H 2 + O 2 type: d. NaF(aq) + Pb(NO 3 ) 2 (aq) type: e. Ag + Ca(NO 3 ) 2 type: f. HgO type: g. SO 2 + H 2 O type: h. NaI(aq) + Pb(C 2 H 3 O 2 ) 2 (aq) type: Unit 9 1. What do the following symbols mean? a. (s) b. (l) c. (g)

10 10 d. (aq) 2. What is a mole to mole ratio? 3. Use the reaction to answer the following questions N 2 (g) + H 2 (g) NH 3 (g) a. How many moles of nitrogen gas are required to produce 20 moles of ammonia (NH 3 )? b. How many grams of hydrogen gas (H 2 ) are required to fully react with 10 moles of N 2? c. How many liters of nitrogen are used to form 3.45g of ammonia? d. How many grams of ammonia are formed when 5.67L of hydrogen react with excess nitrogen? e. What is the total number of moles of H 2 used to produce 34 liters of NH 3? Unit Solubility of gases when temperature increases and when pressure increases. 2. Diffusion ( ) vs. Effusion ( ) 3. Gas Law Practice a. A 400mL sample of gas at a pressure of 400torr is reduced to 150torr at a constant temperature. What is the new volume of the gas? b. A 5.0L container of gas at 100 o C is raised to a temperature of 200 o C. What is the new temperature? c. A tire was at a pressure of 4.0atm and a temperature of 25 o C. If the temperature is raised to 30 o C, what is the new pressure?

11 11 d. A 10.L container of nitrogen gas is at a temperature of 100 o C, 0.80atm. If the pressure is decreased to 0.68atm and volume is increased to 15L, what is the new temperature? 4. Ideal Gas Law How do I know when to use the Ideal Gas law formula? a. Formula: b. How many grams of nitrogen gas are present in a 25.0L container at a pressure of 0.75atm and a temperature of 80 o C? 5. Molar Volume: L = 1 mole at a. How many moles are in 6.6L of nitrogen at STP? b. What volume is occupied by 2.0 moles of nitrogen gas at STP? 6. Stoichiometry of gases (Liters to Liters) a. Given the equation C 5 H 12 + O 2 CO 2 + H 2 O. How many liters of water are produced when 5.0L of C 5 H 12 is burned? 7. Dalton s Law of Partial Pressure a. Formula b. A sample of nitrogen gas is collected over water. The pressure of the water is 21.1mmHg. What is the pressure of the nitrogen gas if the atmospheric pressure is 785mmHg? Unit Know the differences between the states of matter. Gases Liquids Solids Particles (Draw) Does it flow? Compressibility

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13 13 2. Energy and changes of state: For each of the following label: endothermic and exothermic. State the Phase Change Endothermic/Exothermic a. Freezing Liquid to Solid Exothermic b. Melting c. Condensation d. Evaporation e. Sublimation f. Deposition 3. Remember a ph ase change is a ph ysical change! 4. Be able to read the phase diagram, identify regions, phases, and phase changes. 5. Molar Heat of Fusion: The energy required to change one mole of a to a. 6. Molar Heat of Vaporization: The energy required to change one mole of a to a. a. How much heat energy is absorbed when 72g of ice melts at STP? b. How many grams of water are used to release 5000J when turned into vapor?

14 14 7. Freezing and Cooling Curves. Know the different regions, phases, and phase changes. 8. Phase changes can occur with a change of and/or. 9. What is latent energy? Unit Solutions: Solvent, Solute, Electrolytes a. is the dissolving substance b. is the dissolving medium c. can conduct electricity 2. Solution Equilibrium a. Unsaturated Definition: amount is the line b. Saturated Definition: amount is the line 3. Solubility know how to read the solubility curve and how to use solubility rules located in your reference tables. a. How many grams of sodium nitrate can be dissolved in 100g of water at 40 o C? b. At what temperature can 22 grams of potassium chlorate be dissolved in 100g of water? c. What salt is the most soluble at 70 o C? d. What salt is the least soluble at 30 o C? e. What do the dotted lines mean?

15 15 Unit Equilibrium: 2. If K is larger than one are favored 3. If K is smaller than one are favored 4. Write the equilibrium constant for the following: a. Ammonia gas, NH 3, decomposes into nitrogen gas and hydrogen gas. b. CO(g) + H 2 (g) CH 4 (g) + H 2 O(l) 5. For the reaction N 2 (g) + 3Cl 2 (g) 2NCl 3 (g), an analysis of an equilibrium mixture is preformed at a certain temperature. It is found that [NCl 3 ] = 0.19M, [N 2 ] = M, and [Cl 2 ] = 4.3x10 3 M. Calculate K for the reaction. Unit Electrolyte a. Definition b. Things that are electrolytes: 2. Nonelectrolyte a. Definition b. Things that are nonelectrolytes: 3. Make sure you know the different properties of acids and bases. Acid ph Proton Produce in water Taste Phenolphthalein Indicator Litmus Paper Indicator Example Reacts with metal to form: Base

16 16 4. ph scale/ph/poh a. ph of a solution is the negative of the common logarithm of the hydronium ion concentration b. poh of a solution is the negative of the common logarithm of the hydroxide ion concentration c. Formulas (remember these are in your reference tables) i. ph = log[h 3 O + ] ii. poh = log[oh ] iii. K w = [H 3 O + ] [OH ] = 1.0x10 14 iv. ph + poh =14 v. [H 3 O + ] = 10 ph vi. [OH ] = 10 poh d. What is the ph of a substance that has a [H 3 O + ] of 1.0x10 2? e. What is the ph of a solution whose HCl concentration is 1.0x10 1? f. What is the ph of a substance that has a [H + ] concentration of 1.0x10 4? g. What is the poh of a substance that has an [OH ] of 1.0x10 3? h. What is the ph of a solution whose NaOH concentration is 1.0x10 2? i. What is the poh of a substance that has a [OH ] of 1.0x10 3? Is this solution acidic or basic? j. What is the [H 3 O + ], [OH ], poh of a solution which has a ph of 4? 5. As the ph of a solution increase the concentration of [H + ] increases by. a. The concentration of [H + ] increases by when the ph increases from 3.0 to Neutralization a. Acid + Base + Unit 17 b. HCl + NaOH + i. Net Ionic Equation: c. HBr + Ca(OH) 2 + i. Net Ionic Equation: 1. Exothermic reaction (lose/gain) energy and energy is shown on the (Product/Reactant) side. 2. Endothermic reaction (lose/gain) energy and energy is shown on the (Product/Reactant) side.

17 17 1. Specific Heat a. Definition b. Formula c. A 5.0g sample was heated from 200K to 500K and was found to have absorbed 40J of heat. What is the specific heat capacity of the sample? d. How much heat is needed to raise the temperature of 5.0g of gold by 25 o C? e. Joule is defined as 2. Heat ( ) vs. Temperature ( ) 3. List the three parts of the Collision Theory: 4. Reaction rate is proportional to the number of effect collisions a. Increase temperature = number of collisions = rate b. Increase concentration = number of collisions = rate c. Increase pressure = number of collisions = rate d. Increase surface area = number of collisions = rate 5. What are the four ways to effect rate? a. b. c. d.

18 Potential Energy Diagram a. Activation Energy = b. Enthalpy Change = c. Potential Energy of products = d. Potential Energy of reactants = e. Exothermic or Endothermic = Unit Know the properties of the three different nuclear decay particles: Alpha Beta Gamma Symbol Charge Mass Strength 2. Nuclear Fission: of a nucleus into 3. Nuclear Fusion: of a nucleus into 4. Balancing Nuclear Reactions (remember Law of Conservation of Mass) a Cl H b Zn Cu n c U Sr + d. 6 3 Li n 4 2 He + 5. Half life is a. Element 106(Seaborgium) has a half life of 0.90 seconds. If one million atoms of it were prepared, how many atoms would remain after 4.5 seconds? b. Iron 59 is used in medicine to diagnose blood circulation disorders. The half life of iron 59 is 44.5 days. How much of a 2.000mg sample will remain after days?