Final Review C H E M I S T RY 1 4 0 5 Method of Course Evaluation • Lab Report Average • Quiz Average • Lab Practical 50% 20% 30% • Lab is 20% of your final CHEM 1405 grade. • Each student’s lowest lab report grade will be dropped. Final Practical • All personal belongings (e.g. backpacks, purses) will • • • • • • be placed on the printer table or lab bench Final exam will be require only written responses Hands-on lab exercises will not be required A periodic table will be provided A calculator may be used Smartphones, tablets, or computers are not permitted Long pants and close-toed shoes are needed to comply with lab policies Final Practical • Format may contain any of the following: • Fill in the blank • Multiple choice • Math problems • Approximately 60 questions • Average time for completion is 1 hour SDS • Safety Data Sheets • SDS may be found in each lab or on the internet. • They contain valuable information about chemicals you will use in your experiments. Density • Density is a derived quantity. = / Reading a Meniscus Physical Change • A change in matter that does NOT change its identity. • Example: Ice (solid) melts into water (liquid). The compound remains H 2 O despite the change in physical state. Chemical Change • Changes that produce a NEW kind of matter with DIFFERENT properties. • Example: Adding baking soda to vinegar results in products of a salt, water, and carbon dioxide gas. Evidence of Chemical Reaction • Energy change (e.g. heat, light, electricity) • Gas produced (e.g. evolution of carbon dioxide) • Precipitate formed (e.g. sodium chloride solution + silver nitrate solution forms insoluble silver chloride precipitate) • Color change Element • Simplest form of matter. • Cannot be decomposed by chemical means. • Each element can be found on the periodic table. • Example: K – potassium, Fe – iron, S – sulfur, C - carbon Compound • Can be decomposed into simpler substances by chemical means. • Compounds are composed of 2 or more elements found on the periodic table. • Example: H2O – water, CO2 – carbon dioxide, NaHCO3 – baking soda (sodium bicarbonate) Graphing • All graphs need: • • • Title Labeled independent x-axis with units Labeled dependent y-axis with units Electromagnetic Spectrum Spectral Lines • When atoms are excited by an energy source (e.g. electric current, heat), their electrons absorb energy. • This causes the electrons to “jump” from their ground state to an excited state (higher energy). • When electrons fall back (relax) to their ground state they emit energy in the form of a quantum of light. • The light quantum has a characteristic frequency and wavelength which produces a characteristic spectral line. Hydrogen Bright Line Spectrum Spectral Lines • Each element produces its own unique line spectrum. • Line spectra are often used to identify an unknown. Nomenclature • Is the chemical ionic? • Is the chemical covalent? • Is the chemical an acid? • Memorize polyatomic names and charges • Review handout Balancing Equations • Review lab handout • Knowledge of molar mass calculations • Mastery of mole concept Symbols Used in Equations • Reactants → Products • → Yields or produces • (s) solid • (l) liquid • (g) gas • (aq) aqueous: a substance dissolved in water • ↔ A reversible reaction • ∆ Heat Classifying Chemical Reactions • Combination reaction (or synthesis) • Decomposition • Single replacement reaction • Double replacement reaction Combination Reactions • Two elements or compounds → One compound CO2 + H2O → H2CO3 MgO + H2O → Mg(OH)2 3 H2 + N2 → 2 NH3 Decomposition Reactions • One compound → Two or more elements or compounds • Usually requires energy – heat, electricity 2 NaCl → 2 Na + Cl2 2 Ag2O → 4 Ag + O2 2 KClO3 → 2 KCl + 3 O2 Single Replacement Reactions • One element + compound → Different element + compound • An element that is higher in the activity series will replace one that is lower. Mg + CuSO4 → MgSO4 + Cu Cu + 2 AgNO3 → Cu(NO3)2 + 2 Ag 2 KBr + F2 → 2 KF + Br2 Double Replacement Reactions • The cation (+) in one compound reacts with the anion (-) in the other compound. • The ions switch partners. K2CO3 + CaCl2 → CaCO3 + 2 KCl 2 Li3PO4 + 3 BaCl2 → 6 LiCl + Ba3(PO4)2 K2SO4 + Sr(NO3)2 → 2 KNO3 + SrSO4 Diatomic Elements • Hydrogen – H2 • Nitrogen – N2 • Oxygen – O2 • Fluorine – F2 • Chlorine – Cl2 • Bromine – Br2 • Iodine – I2 Molecular Geometry Linear Trigonal planar Trigonal bypiramidal Tetrahedral Octahedral Predicting Molecular Geometry A = Central atom X = Substituent E = Electron Pair Linear Examples: CS2; HCN; BeF2 Trigonal planar Examples: Examples: SO3, BF3, NO3-, CO32- SO2, O3, PbCl2, SnBr2 Tetrahedral Examples: CH4, SiCl4, SO42-, ClO4- Examples: NH3, PF3 ClO3, H3O+ Examples: H2O, SCl2 Acids and Bases 2 H2O ↔ H3O+ + OHpH = -log [H3O+] pH range is 1 - 14 • Acid – A chemical species with a relatively high amount of [H3O+] hydronium ion. pH range of 1 – 6. • Base – A chemical species with a relatively high amount of [OH-] hydroxide ion. pH range of 8 – 14. • Neutral – pH of 7. Buffers and Equilibrium • Buffer – Solutions that resist pH changes when small amounts of strong acid or strong base are added to it. • Buffer – A weak acid and its conjugate base OR a weak base and its conjugate acid. • HA ↔ H+ + A• Since this weak acid is only partially dissociated, small amounts of acid or base will shift the chemical reaction only slightly due to Le Chatelier’s Principle.