Honors Chemistry Review for Semester One Final Exam

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1 Honors Chemistry Review for Semester One Final Exam Chapter 1 What is chemistry? Branches of chemistry Scientific method 1. What is chemistry? 2. Which measurement depends on gravitational force mass or weight? 3. Which branch of chemistry studies the composition of substances? 4. How does qualitative data differ from quantitative data? 5. What is the function of the control in an experiment? Chapter 2 Units of measurement SI units for time, length, mass, density, volume, temperature Formulas for volume and density Metric conversions Scientific notation Dimensional analysis Accuracy and Precision Percent error Significant figures Interpreting graphs How many significant figures do the following numbers have? 1) ) ) 890 4) ) ) ) ) 3.4 x ) 9.0 x ) x 10-2 Solve the following mathematical problems such that the answers have the correct number of significant figures: 1) grams grams 2) 34.1 grams / 1.1 ml 3) 2.11 x 10 3 joules / 34 seconds 4) meters 0.11 m 5) 349 cm cm cm 6) 450 meters / 114 seconds 7) kilograms kilograms 8) 84 m/s x s Convert the following numbers into scientific notation: 1) 3,400 2) ) 101,000 4) ) ) 1,000,000 7) ) 4.50 Convert the following numbers into standard notation: 9) 2.30 x ) 1.76 x ) x ) 8.65 x ) 9.11 x ) 5.40 x ) 1.76 x 10 0

2 16) 7.4 x 10-5 Unit Conversions 1) How many inches are there in 45.6 cm? (There are 2.54 cm in 1 inch) 2) How many centimeters are there in 1.23 x 10-6 kilometers? 3) How many hours are there in 34.5 years? 4) How many inches are there in 355 millimeters? 5) How many milliliters are in a cubic meter? (There are 1,000 L in 1 m 3 ) 6) How many miles are there in 3.44 x 10 8 inches? There are 0.61 miles in 1 km). Density 1. A 5-ml sample of water has a mass of 5 g. What is the density of water? 2. The density of aluminum is 2.7 g/ml. What is the volume of 8.1g? Percent error 1. The accepted density for copper is 8.96 g/ml. Calculate the percent error for each of the following measurements. a g/ml b g/ml c g/ml d g/ml Chapter 3 Properties of matter Physical properties Chemical properties Physical changes Chemical changes Law of conservation of mass Mixtures Heterogeneous Homogeneous Methods to separate mixtures Elements Compounds Law of definite proportions Percent by mass Law of multiple proportions Conservation of mass 1. A 28.0 g sample of nitrogen gas combines completely with 6.0 g of hydrogen gas to form ammonia. What is the mass of ammonia formed? 2. A 13.0 g sample of X combines with a 34.0 g sample of Y to form the compound XY. What is the mass of the reactants? Law of definite proportions 1. A 25.3 g sample of an unknown compound contains 0.8 g of oxygen. What is the percent by mass of oxygen in the compound? 2. Magnesium combines with oxygen to form magnesium oxide. If g of magnesium reacts completely with 6.96 g of oxygen, what is the percent by mass of oxygen in the magnesium oxide? Law of multiple proportions 1. Carbon reacts with oxygen to form two different compounds. Compound I contains 4.82 g carbon for every 6.44 g of oxygen. Compound II contains g carbon for every 53.7 g of oxygen. What is the ration of carbon to a fixed mass of oxygen for the two compounds. Chapter 4 Structure of the atom Development of the modern model of the atom Atomic number Mass number Symbol and hyphen notation Isotopes Calculate average atomic mass Nuclear reactions Types of radiation Fill in the blanks in the following worksheet. Please keep in mind that the isotope represented by each space may NOT be the most common isotope or the one closest in atomic mass to the value on the periodic table.

3 Atomic symbol Atomic number Protons Neutrons Electrons Atomic mass B Pb Mo No 261 Yb Using your knowledge of nuclear chemistry, write the equations for the following processes: 1) The alpha decay of iridium-174 2) The beta decay of platinum-199 3) Write the symbols for an alpha particle, beta particle, gamma ray Average atomic mass 1. Silver has two isotopes silver-107has a mass of amu (52.00%) and silver-109 has a mass of amu (48.00%). What is the atomic mass of silver? 2. Boron-10 and boron-11 are the naturally occurring isotopes of elemental boron. If boron has an atomic mass of amu, which isotope occurs in greater abundance? Chapter 5 Wavelength, frequency, speed, and energy Bohr model of the atom Principle energy levels Energy sublevels Atomic orbitals Electron configurations Orbital diagrams Valence electrons Electron dot structures Draw the bohr model, write the electron configurations and draw the orbital diagrams of the following elements: 1) sodium 2) iron 3) bromine 4) barium 5) neptunium Determine what elements are denoted by the following electron configurations: 1. 1s 2 2s 2 2p 6 3s 2 3p s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 3. [Kr] 5s 2 4d 10 5p 3 4. [Xe] 6s 2 4f 14 5d 6 5. [Rn] 7s 2 5f 11 Wavelength, frequency, speed, and energy 1. The laser in a compact disc (CD) player uses light with a wavelength of 780 nm. What is the frequency of this light? 2. What is the energy of a photon of red light having a frequency of 4.48 x Hz? 3. What is the energy of an ultraviolet photon having a wavelength of 1.18 x J? Electron dot structures 1. Draw the electron-dot structures for each of the following atoms: carbon, arsenic, polonium, potassium, barium Chapter 6 and chapter 7 The periodic table Groups

4 Periods Transition metals Metalloids Non-metals Lanthanides Actinides Alkali metals Alkaline earth metals Halogens Noble gases Trends in the table Chapter 8 Ionic bonds Ions (cations and anions) Monoatomic polyatomic Forming an ionic bond Naming ionic compounds Writing formulas for ionic compounds Metallic bonds Name the following ionic compounds and draw the lewis dot diagram: 1) NH 4Cl 2) Fe(NO 3) 3 3) TiBr 3 4) Cu 3P 5) SnSe 2 6) GaAs 7) Pb(SO 4) 2 8) Be(HCO 3) 2 9) Mn 2(SO 3) 3 10) Al(CN) 3 Write the formulas for the following compounds: 11) chromium (VI) phosphate 12) vanadium (IV) carbonate 13) tin (II) nitrite 14) cobalt (III) oxide 15) titanium (II) acetate 16) vanadium (V) sulfide 17) chromium (III) hydroxide 18) lithium iodide 19) lead (II) nitride 20 silver bromide Chapter 9 Covalent bonds Forming covalent bonds Lewis structures Sigma bonds Pi bonds Naming covalent compounds Naming acids Molecular structures Molecular shapes Polarity Write the names of the following chemical compounds: 1) BBr 3 2) C 2Br 6 3) IO 2 4) CH 4 5) N 2O 3

5 Write the formulas of the following chemical compounds: 6) tetraphosphorus triselenide 7) disilicon hexabromide 8) diselenium diiodide 9) tetrasulfur dinitride 10) silicon dioxide 11) diboron tetrabromide Polarity For each of the following pairs of molecules, determine which is most polar and explain your reason for making this choice: 1) carbon disulfide OR sulfur difluoride 2) nitrogen trichloride OR oxygen dichloride 3) boron trihydride OR ammonia 4) chlorine OR phosphorus trichloride 5) silicon dioxide OR carbon dioxide 6) methane OR CH 2Cl 2 7) silicon tetrabromide OR HCN 8) nitrogen trifluoride OR phosphorus trifluoride Lewis structures Draw the Lewis structures for the following compounds: 1) PBr 3 2) N 2H 2 3) CH 3OH 4) -1 NO 2 5) C 2H 4 Chapter 10 Chemical reactions Reactants Products Word equations Skeleton equations Balanced chemical equations Types of reactions Predicting products Complete ionic and net ionic equations Balancing equations Balance the equations below: 1) N 2 + H 2 NH 3 2) KClO 3 KCl + O 2 3) NaCl + F 2 NaF + Cl 2 4) H 2 + O 2 H 2O 5) Pb(OH) 2 + HCl H 2O + PbCl 2 6) AlBr 3 + K 2SO 4 KBr + Al 2(SO 4) 3 7) CH 4 + O 2 CO 2 + H 2O 8) C 3H 8 + O 2 CO 2 + H 2O 9) C 8H 18 + O 2 CO 2 + H 2O 10) FeCl 3 + NaOH Fe(OH) 3 + NaCl 11) P + O 2 P 2O 5 12) Na + H 2O NaOH + H 2 13) Ag 2O Ag + O 2 14) S 8 + O 2 SO 3 15) CO 2 + H 2O C 6H 12O 6 + O 2 16) K + MgBr KBr + Mg 17) HCl + CaCO 3 CaCl 2 + H 2O + CO 2 18) HNO 3 + NaHCO 3 NaNO 3 + H 2O + CO 2 19) H 2O + O 2 H 2O 2 20) NaBr + CaF 2 NaF + CaBr 2 21) H 2SO 4 + NaNO 2 HNO 2 + Na 2SO 4 Word equations Write the word equations below as chemical equations and balance: 1) Zinc and lead (II) nitrate react to form zinc nitrate and lead. 2) Aluminum bromide and chlorine gas react to form aluminum chloride and bromine gas.

6 3) Sodium phosphate and calcium chloride react to form calcium phosphate and sodium chloride. 4) Potassium metal and chlorine gas combine to form potassium chloride. 5) Aluminum and hydrochloric acid react to form aluminum chloride and hydrogen gas. 6) Calcium hydroxide and phosphoric acid react to form calcium phosphate and water. 7) Copper and sulfuric acid react to form copper (II) sulfate and water and sulfur dioxide. 8) Hydrogen gas and nitrogen monoxide react to form water and nitrogen gas. Writing equations Write equations for the following reactions: 1) The reaction of ammonia with iodine to form nitrogen triiodide (NI 3) and hydrogen gas. 2) The combustion of propane (C 3H 8). 3) The reaction of nitric acid with potassium hydroxide. 4) The reaction of copper (II) oxide with hydrogen to form copper metal and water. 5) The reaction of iron metal with oxygen to form iron (III) oxide. 6) The reaction of AlBr 3 with Mg(OH) 2 7) The decomposition of hydrogen peroxide to form water and oxygen. 8) The reaction of ammonia with sulfuric acid. Predict products and classify type of reaction 1) NaOH + CaBr 2 2) Pb(NO 3) 2 + HCl 3) AgNO 3 + CuSO 4 4) AgF + NiCl 2 5) Ca(OH) 2 + HF 6) Pb(NO 3) 2 + K 2CrO 4 7) NaC 2H 3O 2 + H 2SO 4 8) Cu(OH) 2 + H 3PO 4 9) AgNO 3 + Na 2CO 3 10) Zn + H 2CO 3 Types of reactions Balance the following equations and indicate the type of reaction taking place: 1) NaBr + H 3PO 4 Na 3PO 4 + HBr 2) Ca(OH) 2 + Al 2(SO 4) 3 CaSO 4 + Al(OH) 3 3) Mg + Fe 2O 3 Fe + MgO 4) C 2H 4 + O 2 CO 2 + H 2O 5) PbSO 4 PbSO 3 + O 2 6) NH 3 + I 2 N 2I 6 + H 2 7) H 2O + SO 3 H 2SO 4 Chapter 25 Types of radiation Properties of alpha, beta, and gamma radiation Types of radioactive decay Alpha decay Beta decay Electron capture Positron emission Writing balanced nuclear equations Transmutation Calculating half life Nuclear fission and Nuclear Fusion ***You can also look at the chapter assessments on the following pages: 22, 50-52, 82-84, , , , , , , , ***You can also look at the practice problems on the following pages: , 886

SAINT VIATOR HIGH SCHOOL CHEMISTRY 722 MRS. COYLE

STUDY GUIDE FOR SEMESTER 1 EXAM 2013-2014 Material to know Safety in the Lab Rules Review those safety rules. Chapter 2 SAINT VIATOR HIGH SCHOOL CHEMISTRY 722 MRS. COYLE 1. What is matter? 2. What are