Summary of Basic Chemistry Significant Tell you about the precision of a All nonzero #s and zeros between non-zero #s are significant. 708 has 3 sig figs Figures measurement If a # is less than one, count all the #s after the first non-zero #. EX: 0.0009870 has 4 sig figs. Accuracy = correct If a # is greater than 1 and no decimal point is written, count only the non- Precision = consistent zero #s and zeros in-between non-zero #s. Ex: 408000 has 3 sig figs. If a # is greater than one and a decimal point is written , count all #s. Ex: 9487.000 has 7 sig figs Adding and Subtracting Line up the decimal point. Add or subtract the numbers. Round your answer (round the answer to the to the place of the least decimal places given (the one farthest to the left) LEAST number of decimals Example: 15.00 cm 5100 cm from the givens) + 4.352 cm + 4.3 cm 19.352 = 19.35 cm 5104.3 cm = 5100 cm Multiplying and Dividing Multiply or divide the numbers. Count the total number of significant figures (Round the answer to the in each number. Your final answer should have no more significant figures LEAST # of significant figures than the lowest number of significant figures you started with of all of your givens) Example: 6.7 moles * 1.101 g/mol = 7.3767 = 7.4 g Prefixes kilo (1 km = 1000m), deci (10dm = 1m), centi (100cm = 1m), milli (1000mm = 1m) these prefixes work for any unit (ex: 1kPa = 1000Pa, 1kL = 1000L etc…) Metric Volumes : 1L = 1 dm 3 1mL = 1cm 3 Density = mass values Temperatures : 0 °C = 273K (all gas laws need to be in Kelvin!) volume Pressures : 1atm = 760 mmHg = 101.3kPa 23 1 Mole = 22.4L = 6.022 X 10 molecules = molar mass in grams (MM) density of water = 1g/mL # moles = coefficient # in a balanced equation Subatomic Subatomic Particle Charge Mass Location Formula Particles Proton (defines the type of atom) +1 1 Nucleus = atomic number Neutron 0 1 Nucleus = atomic mass – atomic number Electron -1 0 Orbits around the nucleus = atomic number – charge Average Weighted sum of all isotopes : Example: Atomic Mass There relative abundances of Carbon-12 and carbon-13 are 98.9% and 1.19 % Σ (% Abundance) (mass) respectively. The average atomic mass is: 12*0.989 + 13*0.0119 = 12.02 amu or (do not round the final answer, units are AMU = atomic mass units) (% * mass) + (% * mass) + …. % Error Measure of accuracy (how % Error = | actual – experimental | X 100 correct your data is) actual Molar mass Synonyms: Add up the atomic masses of each atom. Pay attention to subscripts and Formula mass parenthesis. Formula weight Example: Cu(C 2H3O2)2 = Cu + 4C + 6H + 4O molecular mass = 63.546 + (4 * 12.011) + (6 * 1.0079) + (4 * 15.999) = 181.633 g/mol molecular weight Empirical (empirical formulas are the EX: 79.8% C, 20.2% H Formula LOWEST ratio of atoms) Ste p 1for C: 79.8 g (1 mol/12.011 g) = 6.64 mol C Step 1 for H: 20.2 g (1 mol/ 1.008 g) = 20.0 mol H (no diatomics!) Step 1: Calc. moles of each atom Step 2 for C: 6.64 mol C /6.64 mol = 1 C Step 2: Divide by smallest # Step 2 for H: 20.0 mol H /6.64 mol = 3 H moles to find a mole ratio So the empirical formula is CH 3 Molecular (molecular formulas are the The molecular mass of a compound with an empirical formula of CH 3 i s 45 g. Formula UNREDUCED ratio of atoms) What is the molecular formula? Step1: Molar mass = 45 Molar mass = multiplying Empirical mass is C + 3H = (12.011 + 3*1.008) = 15 Emprical massFM factor Step 1: 45 ÷ 15 = 3 Step 2: Step 2: 3 (CH 3) = C3H9 Multiply the empirical formula by the multiplying factor % Mass of part X 100 EX: % composition of Na in Na 3P Composition Mass of total %Na = 3(22.98) X 100 = 69.0% 99.91 +1 -2 -2 -1 Polyatomic Ammonium (NH 4 ) Sulfate (SO 4 ) Carbonate (CO 3 ) Chlorate (ClO 3 ) -1 -1 -3 ions Hydroxide (OH ) Nitrate (NO 3 ) phosphate (PO 4 ) ate always has 1 more oxygen than ite Mole G moles Use the compound's Molar Mass (from the periodic table) Conversions Example: 1 mole of Ca 3(PO 4)2 = 3(40) + 2(31) + 8(16) = 310grams/mol How much would 0.8moles weigh? 0.8 moles x 310 g = 248grams 1 mole Particles moles Use Avogadro's Number: (atoms, molecules, or formula units) Example: 1 mole = 6.022 x 10 23 molecules, atoms, or formula units for ANY substance How many moles of Fe would you have if you had 2.8 x 10 23 atoms 2.8 x 10 23 atoms x 1 mole = 0.46 moles 6.022 x 10 23 atoms Liters Moles GASES ONLY, Use 1 mole = 22.4 L for any gas at STP (STP = 0 °C and 1 atm) Example: How much space would 2.5moles of diatomic chlorine (Cl 2) occupy? 2.5 moles x 22.4 L = 56L 1 mole Multi-step Examples: Convert 10g H2O to molecules: 1 mo1 H O 6.02 X 10 23 molec OH 23 10g H 2O 2 = 3.34 x 10 molecules H 2O 18.015 g H 2O 1 mol 2OH Convert 2.5 L H 2O to molecules: 1 mo1 OH 6.02 X 10 23 molec OH 22 2.5 L H 2O 2 = 6.7 x 10 molecules of H 2O 22.4 L 2OH 1 mol 2OH Remember moles and molecules are NOT the same thing! Electron Rules & Laws: Alkali metal ns 1 2 Config. Alkaline Earth ns Transition = nd ? Aufbau Principle: electrons are added one at a time 5 in order of lowest energy 1 st Halogens = np Noble gases = np 6 Pauli Exclusion Principle: orbitals have a maximum of 2 e - (must be opposite spin ↑↓) Read electron configurations from the Periodic table Hund’s Rule: (pairing rule) electrons fill sublevels p, N = 7 e - = 1s 22s 22p 3 (note that the super scripts = e -) d and f, they must fill each sublevel with one ↑ e - before any gets a paired electron ↑↓) Formula Ionic formulas Ionic Compounds have no charge, so add subscripts so that the total + & - Writing (also called salts ) charges are equal. Made of positive cations TRICK: Switch the charges for subscripts (leave off charge and make and negative anions sure the subscripts are not divisible by anything but 1) 4+ 2- Example: Tin (IV) oxide = Sn O , so you have Sn 2O4 Since these subscripts are divisible by 2… simplify! SnO 2 Molecular formulas Use prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca) to (covalent) determine subscripts . Note: mono is not used for the 1st element Made of two nonmetals Examples: Carbon tetrachloride = CF 4 Tri sulfur hexoxide = S 3O6 (DO NOT REDUCE covalent!) +1 -2 -2 -1 Polyatomic Ammonium (NH 4 ) Sulfate (SO 4 ) Carbonate (CO 3 ) Chlorate (ClO 3 ) -1 -1 -3 ions Hydroxide (OH ) Nitrate (NO 3 ) phosphate (PO 4 ) ate always has 1 more oxygen than ite Naming Ionic Compounds Name the cation. Name the anion (change ending to –ide if an element… if a Compounds (contains a metal and/or a polyatomic ion do not change the ending –ite or –ate) polyatomic ion) (If the cation can have more than 1 charge, determine charge and include as a roman numeral) Examples: Ca 3N2 = calcium nitr ide Ca(NO 3)2 = calcium nitr ate Fe 2O3 = Iron (III) Oxide Molecular Compounds Use prefixes to denote subscripts (mono,di,tri,tetra,penta,hexa,hepta, (two nonmetals) octa,nona,deca)(do not start the 1 st word with mono-) Examples: CO = carbon monoxide ; N 4O8 = tetranitrogen octoxide Naming ( H+ with and anion) Anion ending: -ide = hydro(stem)ic acid EX: HCl = hydrochloric acid Acids -ite = (stem)ous acid EX: HNO 2 = nitrous acid -ate = (stem)ic acid EX: H 3PO 4 = phosphoric acid Balancing Sn(ClO 3)4 → SnCl 4 + O 2 You can't change the compounds at all you can only Add coefficients to Equations Sn(ClO 3)4 → SnCl 4 + 6O2 get the same number of atoms of each element on each side of the Look for common MULTIPLES Types of Synthesis (combination) A + B AB EX: H 2 + O 2 H 2O Equations Decomposition AB A + B EX: CaO Ca + O 2 Single Replacement A + BC B + AC EX: Li + CaSO 4 Ca + Li 2SO 4 Double Replacement AB + CD CB + AD EX: CaCl 2 + NH 4NO 3 Ca(NO 3)2 + NH 4Cl Combustion Carbon compound (C xHy) + O 2 CO 2 + H 2O this part never changes Neutralization Acid + base salt + water (type of double replacement) H2SO 4 + NaOH Na 2SO 4 + HOH a Molarity M = mol & M 1V1 = M 2V2 Molarity = Moles of solute concentration L (dilution) Liters of solution Don’t forget to change grams to moles before using the formula for molatrity Gas Laws The Ideal Gas Law PV=nRT (R=0.0821 L*atm/mol*K or 8.31 L* kPa/Mol*K or 62.4 L*mmHg/mol*K) STP = Combined Gas Law P1V1/T 1 = P 2V2/T 2 Kinetic Molecular Theory 0°C and 1 atm Gases can be compressed or expanded Charles’ Law V1/T 1 = V 2/T 2 (direct proportion) V ↑T↑ Gases fill up the volume of the container All gas laws Move faster with higher temperature require Boyle’s Law P1V1 = P 2V2 (indirect proportion) P ↓V↑ Constantly moving (straight paths until they KELVIN!!! bounce off – elastic collision) Avogadro’s Law V1/n 1 = V 2/n 2 (direct proportion) V ↑n↑ NO attractive forces Less dense than solids and liquids Gay-Lussac’s Law P1/T 1 = P 2/T 2 (direct proportion) P ↑T↑ Dalton’s Law P1+ P 2 +P 3 + P 4 + ….