Chemistry 103 1st Edition Final Exam Study Guide Chapter 1 Matter Measurement Problem Solving Units of measurement Significant figures o For addition subtraction Use smallest number of decimal points o For multiplication division Use smallest number of significant digits Dimensional analysis unit conversions including density and volume o d m v Chapter 2 Atoms and Elements The structure of an atom protons neutrons electrons o Protons Positive 1 charge 1 amu Located in nucleus o Neutrons 0 charge 1 amu Located in nucleus o Electrons Negative 1 charge Negligible mass Floats around the nucleus Isotopes atomic numbers mass numbers o Isotopes Elements can have differing numbers of neutrons Each different version of the element is called an isotope o Atomic number Number of protons in the nucleus of an element o Mass number Number of protons and neutrons in the nucleus Can vary Atomic masses molar masses Moles converting mass moles of particles Chapter 3 Molecules Compounds Chemical Equations Molecules molecular compounds Ions ionic compounds Naming molecular ionic compounds and acids o See packet handout Balancing chemical equations o Be sure there is an equal number of each element on each side of the equation Molar masses of compounds percent composition from formula o Molar masses of compounds Add the molar mass of each element in the compound Ex NaCl 22 99 35 45 58 44 o Percent composition of formula Divide the element s molar mass by the compound s molar mass and multiply by 100 Ex Na in NaCl 22 99 58 44 100 39 3 Empirical formulas Lowest whole number ratio of elements in a compound Ex The empirical formula for H O is HO 2 2 Chapter 4 Chemical Quantities and Aqueous Reactions Quantitative information from balanced equations stoichiometry Limiting reactants theoretical yields percent yields o Limiting reactant The reactant that runs out first o Theoretical yield Using mole ratios finding how much product should be formed o Percent yield Actual yield Theoretical yield X 100 Solution concentration solution stoichiometry titration o Solution concentration Molarity Number of moles of chemical in 1 L of solution Electrolytes Precipitation reactions solubility double displacement reactions net ionic equations Acid base gas evolution reactions o See handout Oxidation reduction reactions oxidation numbers the activity series o Activity series Elements higher up on the activity series are more likely to become oxidized Activity series will be provided for the exam Chapter 7 The Quantum Mechanical Model of the Atom Wave particle duality of light c and E h Bohr Model of the atom de Broglie wavelength of matter Atomic orbitals and quantum numbers Chapter 8 Periodic Properties of the Elements Electron configurations Hund s Rule the Pauli Exclusion Principle condensed electron configurations o Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers Consequently a maximum of two electrons can occupy a given orbital and if two electrons occupy the same orbital they have opposite spins o Know and understand that an electron configuration shows the number of electrons that occupy particular orbitals in atoms and is the basis for chemical reactivity o Know that the spin quantum number ms can have values of 1 2 and 1 2 o Understand that the sublevels within a given principal energy level are from lowest energy to highest s p d f o Understand the general principles of electron shielding and orbital penetration o Write electron configurations Effective nuclear charge General trends in o Atomic ionic sizes o Ionization energy o Electron affinity o Metallic character Chapter 9 Chemical Bonding I Lewis Theory Ionic bonds electron transfer Covalent bonds electron sharing Lewis symbols the octet rule Bond polarity and electronegativity dipole moments o Know and understand that a pair of electrons does not have to be shared equally between two atoms Unequal sharing results in a polar covalent bond o Define electronegativity and know its periodic trends o Understand that bonds can range from a nonpolar covalent bond to a polar covalent bond to an ionic bond depending on the difference in electronegativity between the two atoms o Define dipole moment and percent ionic character Drawing Lewis dot structures formal charge resonance structures exceptions to the octet rule Chapter 10 Chemical Bonding II Molecular Shapes Valence Bond Theory Molecular shapes VSEPR model effect of lone pairs multiple bonds Molecular polarity related to shape electron geometry vs molecular geometry Covalent bonding and orbital overlap hybrid orbitals multiple bonds Chapter 5 Gases The individual gas laws Charles s Boyle s Avogadro s o Boyle s Law P V k o Charles s Law V T k o Avogadro s Law V mol k The ideal gas law o P V mol T k Densities of gases amounts of gases in chemical reactions stoichiometry o d M P R T M molar mass P pressure R gas constant T Temperature in Kelvin Gas mixtures partial pressures mole fractions Kinetic molecular theory of gases Chapter 11 Liquids Solids and Intermolecular Forces Intermolecular forces In order of increasing relative strength o London Dispersion with nonpolar covalent bonds o Dipole dipole forces with a polar bond o Hydrogen bonding when a hydrogen is bonded to O F or N o Ion dipole A dipole with a metal ion Vapor pressure boiling point Chapter 12 Solutions Factors that affect solubility solute solvent interactions pressure effects temperature effects Concentration units o ppm mass of solute mass of solution X 10 o mass mass of solute mass of solution X 100 o ppb mass of solute mass of solution X 10 o Mole Fraction X moles of solute moles of solute moles of solvent o Molarity M moles of solute liters of solution o Molality m moles of solute kg of solvent Colligative Properties 6 9 solute Chapter 19 Nuclear Chemistry Handout from class Radioactivity nuclear equations types of radioactive decay Nuclear stability neutron to proton ratio magic numbers Rates of radioactive decay half life