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Chapter 1 Notes Scientific Notation Convert the following numbers to scientific notation 17600.0 4.76 0.00135 -0.1544 10.2 301.0 -67.30 -0.000130 How to expand scientific notation 1) If the exponent is ___________, move the decimal to the _______ to make the number _________. 2) If the exponent is ___________, move the decimal to the _______ to make the number _________. Samples: 1) Expand the scientific notation of 8.02 x 10-4 to regular notation. 2) Expand the scientific notation of -9.77 x 105 to regular notation. Expand the following numbers to regular notation 4.96 x 10-2 7.01 x 100 5.50 x 104 5.61 x 10-3 -9.3 x 10-3 4.92 x 102 -8.37 x 101 -9.23 x 10-1 No Calculator Practice To multiply numbers in scientific notation without a calculator 1) Multiply the number in front 2) Add the exponents 3) Put the answer in correct scientific notation if necessary To divide numbers in scientific notation without a calculator 1) Divide the number in front 2) Subtract the exponents 3) Put the answer in correct scientific notation if necessary (2.3 x 103)(1 x 102) = (6 x 10-7) / (3 x 10-8) (4.0 x 105) / (2 x 103) = (2x103)(4x10-2) / (1x10-4) 27 x (3.4 x 102) (9x10-7) / (1x105)(3x10-1) Metric Prefixes Prefix Abbreviation Scientific notation equivalent Prefix Tera Centi Giga Milli Mega Micro Kilo Nano deci Pico Abbreviation Convert 4.98 pL to L Convert 87 µm to Mm Convert 56 mm to meters Convert 0.006 kL to nL Scientific notation equivalent Uncertainty of Measurement Accuracy = How close to the ___________________ you are. Precision = How close your measurements are _______________________________. Rules for counting significant figures 1. Nonzero integers always count Ex 456 _____ sig figs 2. Leading zeros never count. Ex 0.0025 _____ sig figs 3. Captive zeros always count. Ex 1001 _____ sig figs 4. Trailing zeros sometimes count Ex 1020 _____ sig figs 1.00 _____sig figs 5. Exact numbers have an infinite number (because it’s not a measurement) Ex 5 students How many significant figures 20600 0.00960 10 apples 3.90 x 106 10.90 Rules for math functions 1. Multiplication or division The answer should have the same number of _______________________ as the _________ precise measurement. Ex. 1.34 x 0.04 = ______ 2. Addition or subtraction The answer should have the same number of ______________________ as the _________ precise measurement. Ex. 1.34 + 2.3 = _______ Practice: The solution to 9.99/22.41 x (18.465 + 0.464) is _________. Density __________ The denser object will be on the _______________. Practice: 1) Water has a density of 1 g/mL. If you have 150mL, how many mg does it weigh? 2) You have an unknown object that weighs 60g. You place it in a graduated cylinder which has 30mL of water in it. When you place the object in the water the final volume of the graduated cylinder of reads 75mL. What is the density of this unknown object? Temperature Conversions Temperature in Celsius + 273 = Kelvin Temperature o C Kelvin o C 56oC Kelvin 190 K 780 K -93oC Classification of Matter Element – matter composed of __________________________ of atom. Compound – pure substances that are ______________ joined that consist of ___________________ of atom. Homogeneous Mixture – a ___________ mixture of two or more substances which are __________________ bonded. Heterogeneous Mixture – mixture that is __________________ in composition. Element Compound Copper Chocolate Chip Cookie Mercury Coke Sodium chloride Vegetable soup Rocky Road Ice Cream Homogeneous Mixture Heterogeneous Mixture Oil and water Iron Rust Tap water Milk Ways to separate a mixture 1) ______________= separating a solution due to differences in boiling points 2) ______________ = Separates a solid and liquid mixture, or a mixture of two liquids at are immiscible. 3) ______________ = Used to separate mixtures by their polarity Chapter 2 Notes Atomic Structure Protons are located in the ___________ and have a __________ charge. Electrons are located in the ______________ and have a _____________ charge. Neutrons are located in the ____________ and have a ____________ charge. Atomic Number = ______________________________________ Atomic Mass = ________________________________________ Chemical Symbol 23 11 23 11 Write the chemical symbol for Magnesium-23 Isotope = Atoms that contain the same number of __________________ but a different number of ________________. 12 _______ _______ Ex: C-12 and C-14 14 _______ _____ Atoms are neutral due to equal _______________ and _____________________. Ions have charges. Cations = ____________________ Anions = __________________ Practice: 1) How many electrons does Ca2+ have? __________ 2) How many protons does iron have? __________ 3) How many electrons does F-1 have? __________ 4) How many total protons are found in two molecules of C20H30O? __________ Bonding and Naming ______________ = One loses electrons and one gains electrons. Occurs between metals and nonmetals. ______________ = Sharing of electrons. Occurs between two non-metals. Use the naming tree to review your naming rules. How to name acids if you’re given their formulas: In summary for acids: Examples: 1) HNO3 __________________________ ____ide hydro ____ic acid ____ate _____ ic acid 2) HF __________________________ 3) H2C2O4 __________________________ 4) Hydrosulfuric acid ______________ 5) Chromic acid ______________ ____ite _____ ous acid Prefixes: 1- mono 2- di 3- tri 4- tetra 5- penta 6- hexa 7- hepta 8- octa 9- nona 10- deca Remember that polyatomic ion names are not changed in any way. (do not change ending to –ide) Practice Naming: 1) NaF ____________________ 4) Li3PO4 ____________________ ____________________ 5) MgO ____________________ 3) H2SO4 ____________________ 6) CuO ____________________ 2) NO2 Practice Formula Writing: A. Iron (III) chloride _____________ D. Hydrofluoric acid _____________ B. Calcium sulfite _____________ E. Zinc chloride _____________ C. Sulfur hexafluoride _____________ Chapter 3 Notes Mass spectrometer = tells you which _________________________ are in your sample. Average Atomic Mass = is found on the periodic table, it is the average of all of an element’s ______________. = (mass isotope 1) (% abundance isotope 1) + (mass isotope 2) (% abundance 2) + …… Practice: 1) Assume that element Uus is synthesized and that it has the following stable isotopes: 284 Uus (283.4 amu) 34.60% 285 Uus (284.7 amu) 21.20% 288 Uus (287.8 amu) 44.20% What would the average atomic mass be? 2) What is the average atomic mass of an element with two isotopes, one isotope has a mass of 64 and is 20% abundant. The other isotope has a mass of 55 and is 80% abundant. Moles 1 mole = molar mass 1 mole = 6.022 x 1023 atoms/ molecules 1 mole = 22.4 liters Practice: 1) Calculate the molar mass of calcium chloride 2) How many moles are in 496g of CH4? 3) What is the molar mass of iron (III) chloride 4) How many liters are in 15 moles of oxygen gas 5) How many molecules are in 54.9g of NaCl? 6) How many liters would 100g of CaCl2 take up? % composition Part/whole x100 Practice: 1) Calculate %H in water 2) Calculate %N in ammonia 3) What is the % composition of N in nitrogen dioxide? Empirical Formula = The simplest whole number ratio of atoms in a compound. Steps to Determine the Empirical Formula 1) Convert % to decimal form for each element 2) Divide decimal form by molar mass 3) Divide each answer by the smallest # 4) Multiply all by a # (if necessary) to get all elements to a whole. Example: 1) Determine the empirical formula of a compound that is 71.65% Cl, 24.27% C, and 4.07% H. 2) NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of Nutrasweet. 3) Suppose a substance has been prepared that is composed of carbon, hydrogen, and nitrogen. When 0.1156g of this compound is reacted with oxygen, 0.1638g of CO2 and 0.1676g of water is collected. Assume all Carbon is converted to CO2. I started with how much carbon? What is the empirical formula for the compound? 4) What is the empirical formula of the compound N6O10? Molecular Formula = the exact formula of a molecule _______________ of the empirical formula Divide the empirical formula’s molar mass by the molar mass of the compound. Practice: 1) Calculate the molecular formula of a substance if the empirical formula is ClCH2 and has a molar mass of 98.96 g/mol. 2) A white powder is analyzed and found to contain 43.64% P and 56.36% O by mass. The compound has a molar mass of 283.88 g/mol. What are the compound’s empirical and molecular formulas? 3) Determine the molecular formula of a compound that contains 26.7% P, 12.1%N, 61.2% Cl and a molecular weight of 580. Balancing Equations = you can only add coefficients to make the atoms on the reactants equal the atoms on the products side. _____Ca(OH)2 + _____H3PO4 _____H2O + _____Ca3(PO4)2 _____Al(OH)3 + _____HCl _____AlCl3 + _____H2O _____Au2S3 + _____H2 _____Au + _____H2S Converting one substance to another = Stoichiometry 1) Balance equation 2) Convert known mass to moles 3) Use appropriate mole ratios to convert moles of given substance to moles of other substance. 4) Convert back to grams if required _____CH4 + _____O2 _____CO2 + _____H2O If you want to produce 30g of water how much methane (CH4) should you start with? _____H2 + _____O2 _____H2O If you have 5g of oxygen, how much water can you produce if you have excess hydrogen? _____C7H6O3 + _____C4H6O3 _____C9H8O4 + _____HC2H3O2 What mass of C4H6O3 is needed to completely consume 1.0 x 102g C7H6O3? What is the maximum amount of C9H8O4 that could be produced if you started with 1.0 x 102g C7H6O3? Limiting reactant (Reagent) = the reactant that is consumed _______ so it limits the amount of products that can be made. Solve the problem twice, once with each reactant, the smaller answer will be the correct answer. N2 + 3H2 2NH3 Calculate how much NH3 you would produce if you started with 15g of nitrogen and 25g of hydrogen. _____H2 + _____O2 _____H2O Determine how much water you would produce if you started with 6g of hydrogen and oxygen gas. Hg + Br2 HgBr2 What mass of HgBr2 can be produced from the reaction of 10g Hg and 9g Br2? What mass of which reactant is left over? COCl2 + 2NaOH 2NaCl + H2O + CO2 Identify the limiting reagent, and the number of grams left over in excess, respectively, for the above equation, if 1.9g of each phosgene (COCl2) and sodium hydroxide are combined. % _______________ !100 Actual yield = _______________________ Theoretical yield = ___________________ 2H2 + CO CH3OH Suppose 68.5 kg CO is reacted with 8.6kg H2. Calculate the theoretical yield of methanol (CH3OH) if 3.57 x 104g CH3OH is actually produced, what is the % yield. Consider the following unbalanced reaction: _____P4 + ____F2 _____PF3 How many grams of fluorine are needed to produce 120g of PF3 if the reaction has a 78.1% yield?