General Chemistry: Semester 1 Final Review

Transcription

1 General Chemistry: Semester 1 Final Review Name Per Unit 1: Classification of Matter Define: chemistry, matter, element, compound, pure substance, homogeneous mixture, heterogeneous mixture, chemical change, and physical change Describe characteristics, indicators, and examples of elements, compounds, pure substances, homogeneous mixtures, heterogeneous mixture, chemical changes, and physical changes Distinguish between: elements and compounds, and pure substances and mixtures. Classify matter as an element or compound and as a pure substance or mixture. Distinguish between mixtures and classify them as homogeneous or heterogeneous mixtures. Unit 2 & 3: The Atom & Periodic Table Define: atomic number, mass number, isotope, ion, and average atomic mass. Describe the three subatomic particles of the atom (location, weight, and charge). Calculate the number of protons, neutrons, electrons, mass number, and charge of an atom, isotope, or ion. Draw a diagram of an atom that shows the correct location and number of neutrons, protons, and electrons. Identify an element when given its atomic number. Draw the atomic symbol (nuclear symbol) for an element when given its name, mass number, atomic number, number of neutrons, number of protons, number of electrons, and/or its charge. Calculate mass number, average atomic mass, Provide an example of an isotope and an ion. Distinguish between mass number and average atomic mass. Distinguish between atoms and ions. Identify and label every group on the periodic table. Describe the unique characteristics for each group on the periodic table. Distinguish between metals and non-metals. Distinguish between groups and periods. Identify an element given its period number and group name or number. Explain how the periodic table is organized. Unit 4: Naming Define ionic compound, covalent (molecular) compound, and polyatomic ion. Distinguish between ionic and covalent (molecular) compounds. Use correct prefixes and suffixes for naming covalent compounds. Determine the chemical formula of a covalent compound. Determine the chemical name of an ionic compound. Use the periodic table to determine the chemical formula of an ionic compound. Convert polyatomic ion names into their chemical formulas (and vice versa). Determine the name or chemical formula of an ionic compound that contains a polyatomic ion. Determine the charge of a transition metal when given a roman numeral or by its chemical formula. Determine the name or chemical formula of an ionic compound that contains a transition metal. Unit 5: The Mole Calculate the molar mass of a compound given its name or chemical formula. Define the term mole. Explain Avogadro s number and the possible units that can be used with it. Convert moles into grams (mass), grams (mass) into moles, moles into particles, particles into moles, grams (mass) into particles, and particles into grams (mass). Calculate the percent composition (mass percent) of a chemical compound. Distinguish between empirical and molecular formulas. Determine the empirical formula of a compound. Determine the molecular formula of a compound. Unit 6: Chemical Reactions Explain the law of conservation of matter and how it relates to balanced chemical equations. Balance chemical equations. Define and classify the following types of reactions: synthesis, decomposition, single replacement, double replacement, and combustion. Convert written chemical reactions into balanced chemical equations using chemical compounds (including states of matter symbols). Predict the possible products of chemical reactions (including states of matter symbols). Write balanced chemical equations using diatomic elements properly (Br I N Cl H O F).

2 Unit 7: Stoichiometry Balance and write chemical equations. Identify the mole ratio between two compounds in a balanced chemical equation. Use the mole ratio in a balanced chemical equation to determine the amount of product produced in a chemical reaction. Use the mole ratio in a balanced chemical equation to determine the amount of reactant needed in a chemical reaction. Write out all the steps of stoichiometry. Use a balanced chemical equation to convert moles of one compound into moles of another compound. Use a balanced chemical equation to convert moles of one compound into grams of another compound. Use a balanced chemical equation to convert grams of one compound into moles of another compound. Use a balanced chemical equation to convert grams of one compound into grams of another compound. Use stoichiometry to determine the volume (liters or ml) of a compound reacted or produced in a chemical reaction. Unit 8: Atomic Theory Explain how John Dalton contributed to our current model of the atom. What was his model of the atom? Explain how J.J. Thomson contributed to our current model of the atom. What did he discover? Describe his experiment. What model of the atom did he develop? Explain how Ernest Rutherford contributed to our current model of the atom. What did he discover? Describe his experiment. What model of the atom did he develop? Explain how Niels Bohr contributed to our current model of the atom. What did he discover? Describe his experiment. What model of the atom did he develop? Explain the properties of electromagnetic radiation. List the colors of visible light in order of increasing energy. Distinguish between the bright line spectrum and continuous spectrum. Explain how electromagnetic radiation can be absorbed and released by atoms. How do energy levels account for the existence of a unique line spectrum for each element? Describe the current quantum model of the atom. Unit 9: Quantum Model Describe the current quantum model of the atom including the term probability density. Explain what energy levels are and how they relate to periods on the periodic table. Label them on the periodic table. Explain what sublevels (s, p, d, and f) are and state how many electrons each one can hold. Label the S block, P block, D block, and F block on the periodic table. Explain what an atomic orbital is and how many electrons each one can hold. State the number of atomic orbitals that are in the s, p, d, and f sublevels. Describe the 3-D shapes of the s, p, and d sublevels. Explain how electron configuration models the location of all electrons in an atom. Explain how electron configuration models the location of all electrons in an atom. Determine and draw the electron configuration for any atom or ion. Draw the orbital diagram for any atom or ion. Use electron configuration to identify an element. Unit 10: Periodic Trends Identify and label every group on the periodic table. Describe the unique characteristics for each group on the periodic table. Distinguish between groups and periods. Define electronegativity. Define ionization energy. Define atomic radius. Explain how electronegativity changes across a period and down a group. Explain how ionization energy changes across a period and down a group. Explain how atomic radius changes across a period and down a group. Explain how ionization energy and electronegativity affect atomic radius. Explain the properties of non-metals in terms of electronegativity and ionization energy. Explain the properties of metals in terms of electronegativity and ionization energy. Describe how metals tend to react. Describe how non-metals tend to react. Define valence electron. Identify elements based on their reactivity, size, group location, energy level, and number of valence electrons.

3 Unit 11: Bonding Describe a non-polar covalent bond and how it is formed. Describe a polar covalent bond and how it is formed. Describe an ionic bond and how it is formed. Describe a metallic bond and how it is formed. Identify and draw a non-polar covalent, polar covalent, ionic, and metallic bond. Distinguish between covalent and ionic bonding when given a chemical compound formula. Explain the octet rule. Determine the number of valence electrons a specific atom has. Draw a Lewis dot structure. Draw a Lewis structure. Define lone pair. Define polar bond Define resonance. Determine if a Lewis structure is polar or if it has resonance. Use valence shell electron pair repulsion (VSEPR) to determine the geometry (3-D shape) of a Lewis structure compound. Define London forces, hydrogen bonds, and dipole-dipole interactions (intermolecular bonds) Identify the possible intermolecular bonds a compound might have. Unit 12: States of Matter and Gas Laws Define the terms solid, liquid, and gas. Describe characteristics for each of the states of matter. Define freeze, melt, boil, condense, sublime, and deposition. Define triple point and critical point. Define pressure, temperature, and volume. Draw a properly labeled phase diagram of water. Define absolute zero. Convert absolute zero into Celsius and Kelvin. Outline the kinetic molecular theory. Describe an ideal gas. Calculate pressure, temperature, or volume using the combined gas law. Calculate pressure, temperature, volume, moles (grams), or pressure using the ideal gas law. Describe the relationship between volume and pressure. Describe the relationship between volume and temperature. Determine which gases will diffuse fastest. Describe Dalton s law of partial pressure and use it mathematically.

4 PRACTICE QUESTIONS Unit 1 1. Define the following terms: matter, pure substance, homogeneous mixture, heterogeneous mixture, element, atom, and compound. 2. What are the two types of mixtures? Describe each and say how we can tell them apart. 3. Classify each of the following as an element, compound, heterogeneous mixture, or homogeneous mixture. Element, Compound, Mixture?(Y/N) Homogeneous, Heterogeneous? a. Gold (Au) b. Kool-Aid completely dissolved in water c. Sodium bicarbonate (H 2 CO 3 ) d. Sulfur (S) e. NaCl f. Salsa 4. Classify each of the following changes as either chemical or physical and EXPLAIN YOUR REASONING. a. Cutting a sheet of aluminum foil into 4 pieces b. Burning of coal c. Cooling a liquid until it freezes d. A white solid and sulfuric acid are mixed and an orange gas is produced e. Dissolving sugar in tea f. A piece of copper is hammered into a thin sheet

5 5. Give three examples of physical changes. 6. Give four examples of SIGNS of chemical changes. Units 2 & 3 1. Define (or write the equation that defines them): atomic number, mass number, average atomic mass, atom, ion, and isotope. 2. Provide an example of an ion. 3. Provide an example of two or more isotopes: 4. Fill in the following table: Subatomic Particle Mass (amu) Charge Location Proton Neutron Electron 5. What two subatomic particles contribute to an atom s mass? 6. What two subatomic particles contribute to an atom s charge?

6 7. Fill in the following table: Atomic Symbol Atomic # # of protons # of neutrons # of electrons Mass # **Put a star next to all the ions in the table above** Fill in the atomic symbol for the following table: Atomic # of protons # of neutrons # of electrons Symbol Using the table above, which elements are isotopes of each other (2 pairs)? 10. Write the atomic symbol for an ion with 17 protons, 19 neutrons, and 18 electrons. 11. Write the atomic symbol for an atom with 6 protons, 8 neutrons, and 6 electrons. 12. Write the atomic symbol for an ion with 3 protons, 3 neutrons, and 2 electrons. Draw a diagram of this ion. 13. Calculate the average atomic mass of a sample that contains 80% Iodine-127, 17% Iodine-126, and 3% Iodine-128 (the number after the element is the mass number). 14. Calculate the average atomic mass of a sample that contains 50% Gold-197 and 50% Gold-198 (the number after the element is the mass number)

7 15. Fill in the following statements with the appropriate elemental symbol: a. is the 6th period alkaline earth metal b. is the 2nd period noble gas c. is the 7th period alkali metal d. is the 1 st period halogen e. is the transition metal with the lowest atomic number f. is the 8 th Lanthanide element g. is the only non-metal in the boron family h. is the 1 st Rare Earth metal i. is the element in the nitrogen family that is in the 6 th period j. is the element in the 3 rd period and 2 nd family Use the blank periodic table to complete the following questions. 16. Label the chart with the following groups: Hydrogen Family, Transition Metals, Halogens, Alkaline Earth Metals, Alkali Metals, Metalloids, Noble Gases, Rare Earth Metals, Boron Family, Carbon Family, Nitrogen Family, and Oxygen Family. 17. Label metals and nonmetals on the chart below. 18. Label periods and groups on the periodic table (write their numbers in). 19. For each of the groups labeled on the table above, give TWO properties all elements in that group possess.

8 Unit 4 1. The two types of compounds are ionic and covalent (molecular). How are these compounds different? 2. Briefly outline the rules of naming covalent compounds: 3. Briefly outline the rules of naming ionic compounds. Be sure to include the extra steps for transition metals and polyatomic ions: 4. What does the Roman numeral with a transition metal represent? 5. Define polyatomic ion. Provide an example. 6. Identify each compound as either ionic or covalent. Then name it based on the proper set of rules. Compound Ionic/Covalent Name CuCO 3 Mg(NO 3 ) 2 SiF 8 Cu(OH) 2 SO 3 Pb(NO 3 ) 3 Ca 3 N 2 C 2 I 3

9 BrF MgSO 4 FeO N 7 O 4 CaSO 4 7. Identify each compound as either ionic or covalent. Then write a correct formula. Compound Ionic/Covalent Formula nitrogen tribromide lead (V) phosphate ammonium sulfide trichlorine monoxide sodium fluoride aluminum phosphide calcium iodide manganese (II) phosphate potassium phosphide iron(iii) oxide iron(ii) oxide barium chloride dihydrogen monoxide

10 Unit 5 1. Calculate the molar mass of Fe 2 (SO 4 ) 3, Na 2 O, NH 4 NO 3 and C 5 H 6 2. How many moles of NO 2 are present in 7.61 x molecules of CO 2? 3. How many grams of magnesium chloride are present in 4.3 moles of magnesium chloride (**remember naming rules and use charges**)? 4. How many moles of bismuth are present in 88 grams of bismuth? 5. How many atoms of sulfur are present in 15.2 moles of sulfur? 6. How many grams of sodium bromide are present in 97.8 x formula units of sodium bromide? 7. How many atoms of silver are present in 9.8 grams of silver? 8. Calculate the percent composition for copper (II) nitrate, Cu(NO 3 ) Calculate the percent composition for sodium sulfide. 10. Determine the empirical formula of a compound containing 20% carbon and 80% oxygen. 11. Determine the empirical formula of a compound containing 27.3% nitrogen and 72.7% oxygen.

11 12. Determine the molecular formula of a compound if the empirical formula is CH 3 and molar mass of molecular formula is 45g/mol. 13. Determine the molecular formula of a compound if the empirical formula is NO 2 and molar mass of molecular formula is 276 g/mol. 14. Determine the molecular formula of a compound containing 82.4% nitrogen and 17.6% hydrogen when the molecular formula is 68 g/mol. Unit 6 1. What law is met by balancing equations? What does this law state? 2. Balance the equation below. LABEL the following: subscript, coefficient, reactant, product. Al + Fe 2 O 3 Al 2 O 3 + Fe 3. Balance the following equations: a. Al(NO 3 ) 3 + NaOH Al(OH) 3 + NaNO 3 b. KClO 3 KCl + O 2 c. Mg (s) + H 2 O (l) Mg(OH) 2 (s) + H 2 (g) d. NH 3 (g) + O 2 (g) NO (g) + H 2 O (g) e. Na 2 CO 3 (s) + HCl (aq) NaCl (aq) + CO 2 (g) + H 2 O (g) f. Mg(OH) 2 + (NH 4 ) 3 PO 4 Mg 3 (PO 4 ) 2 + NH 3 + H 2 O 4. The five types of reactions are listed below. Define each type. a. Single replacement

12 b. Double replacement c. Synthesis d. Decomposition e. Combustion 5. Write the following chemical equations with chemical FORMULAS (ex. H 2 O). Remember BrINClHOF! a. Sodium metal is burned in oxygen gas to form solid sodium oxide. b. Aqueous sodium chloride is combined with aqueous magnesium iodide to form aqueous sodium iodide and solid magnesium chloride. c. Methane (CH 4 ) is burning in oxygen gas to form carbon dioxide and water. d. Solid zinc iodide is heated to create zinc metal and iodine gas. e. Chlorine gas is mixed with solid sodium oxide to produce solid sodium chloride and oxygen gas. 6. For the reactions in #7, state which type of reaction is occurring. a. d. b. e. c. 7. Determine the type of reaction for each listed below, and then write in the NAMES of the products. a. Reaction Type: Na + O 2 b. Reaction Type: C 3 H 6 + O 2 c. Reaction Type: AlCl 3 d. Reaction Type: Cu + ZnO e. Reaction Type: NaCl + BaS 8. Write out a balanced chemical equation for the following and classify the reaction type. a. An aqueous solution of barium hydroxide is mixed with aqueous iron(iii) sulfate.

13 b. Magnesium metal is burned in nitrogen gas. c. Aluminum foil is immersed in aqueous silver nitrate solution. d. Ethanol (C 2 H 5 OH) is burned completely in air. e. sodium nitride is heated to decomposition. Unit 7 1. What is the difference between a mole problem and stoichiometry? 2. 2KClO 3 2KCl + 3O 2 a) What is the mole ratio for KClO 3 to KCl? b) What is the mole ratio for KCl to O 2? c) What is the mole ratio for O 2 to KClO 3? 3. Solve: Tin(II) fluoride is formed by: Sn (s) + 2HF (g) SnF 2(s) + H 2(g) a. How many moles of HF are needed to produce 9.40 moles H 2? b. How many grams of H 2 are produced by the reaction of tin with 20.0 moles HF? c. How many grams of SnF 2 can be made by reacting 7.42 Liters of HF with tin? 4) What is the percent yield if 4.65g if copper us produced when 1.87g of aluminum reacts with an excess of copper(ii) sulfate? 2Al (s) + 3CuSO 4(aq) Al 2 (SO 4 ) 3(aq) + 3Cu (s)

14 Unit 8 1) How many orbitals are in the following sublevels? a. 3p sublevel b. 2s sublevel c. 4f sublevel d. 4p sublevel e. 3d sublevel 2) Write the complete electron configuration for each atom: a. lithium b. chlorine c. vanadium d. krypton 3) Define ground state 4) Arrange the following sublevels in order of decreasing energy: 2p, 4s, 3s, 3d, 3p 5) What percentage of the time would you find an electron within the boundaries of the electron cloud? 6) How many electrons are in the 2 nd energy level of the following: a. chlorine b. phosphorus c. potassium 8) List the colors of the visible spectrum in order of increasing wavelength. 9) Draw the shapes for the s, p, and d orbitals? 10) How many electrons are in the outermost principal energy level of phosphorus? 11) What is released when an electron moves from a high to low energy level? Unit 9 & 10 1) What is the general electron configuration for noble gases? 2) Explain the following trends as you move across and down the periodic table, and describe why the trend occurs: a. atomic radius

15 b. ionic size c. ionization energy d. electronegativity 3) Where would you find the highest and lowest element for the above trends? Unit 11 1) Define: a. valence electron b. octet rule c. metallic bond d. ionic bond 2) How many valence electrons do the following have: a. alkali metals b. alkali earth metals c. noble gases d. halogens 3) How many electrons would the following need to gain to achieve noble gas configuration? a. H b. Li c. Kr d. P e. Mg f. Al g. O h. N 4) Which of the following pairs is most likely to form an ionic compound? sodium and potassium phosphorus and oxygen lithium and hydrogen nitrogen and calcium

16 1) Define: a. unshared pairs b. covalent bond c. resonance structures d. dipolar e. polar bond f. nonpolar bond 2) How many covalent bonds do each of the following diatomic contain? a. F b. O c. P d. Br e. N 3) Rank the following in order from lowest strength to highest strength: hydrogen bonds dipole interactions London dispersion forces 4) Draw and name the VSEPR configuration for the following. What type of bonds do each have? What is the hybridization and polarity for each? Star any with resonance. a) NH 3 b) H 2 c) H 2 O d) HF e) H 2 S

17 f) N 2 g) CO 2 h) H 3 P i) CH 4 6. Why do elements form chemical bonds? 7. An ionic bond is characterized as the of electrons between an element with electronegativity and an element with electronegativity. 8. Covalent bonding is defined as the of electrons between an element with electronegativity and another element with electronegativity. 9. List the three types of intermolecular forces in order of weakest to strongest. van der Waals, dipole-dipole, H- bonds Unit Write the kinetic molecular theory 2. What is the lowest temperature possible called? 3. What is the lowest temperature in Kelvin and Celsius? 4. Draw a phase diagram. Label all parts 5. How are temperature and volume related? 6. How are pressure and volume related? 7. How are temperature and pressure related? 8. How are pressure and number of particles related? 9. What is STP? What are the values for temperature and pressure at STP? 10. If a gas has a temp of 15 C and heats up to 45 C, what will happen to the 2.2atm pressure?

18 11. If 4.5g of Helium are in a 2.3L balloon and the temperature of the room is 23 C, what pressure will the balloon have? 12. Why does a bag of potato chips that is unopened get bigger or even burst when you drive to the top of Pikes Peak? 13. If I have He, H 2, and O 2, which will move fastest? Why?