Final Exam Study Guide
Final Exam Study Guide 111
Popular in General Chemistry 1
Popular in Chemistry
This 8 page Study Guide was uploaded by Morgan Deal on Friday December 4, 2015. The Study Guide belongs to 111 at University of South Carolina taught by Dr. Thomas Vogt in Summer 2015. Since its upload, it has received 259 views. For similar materials see General Chemistry 1 in Chemistry at University of South Carolina.
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Date Created: 12/04/15
FINAL EXAM STUDY GUIDE Chapter 2: Atoms, Molecules and Ions Dalton’s Atomic Theory o Matter is composed of atoms o An element is composed only of one type of atom o A compound contains atoms of different elements; the ratio of atoms from each element is the same throughout a compound o Atoms do not change identity, the way they are connected changes Law of Constant Composition- all samples of a pure substance contain the same proportion of elements Law of Conservation of Mass- atoms do not experience any change in identity or mass in a chemical reaction Atoms- the smallest unit of an element with all of its properties o Electron- negatively charged subatomic particle o Proton- positively charged subatomic particle; mass of 1 o Neutron- subatomic particle with no charge; mass of 1 o Nucleus- center of the atom containing protons and neutrons; very small but contains most of the atom’s mass Describing atoms and ions o Atomic number (Z)- # protons in the nucleus o Mass number (A)- # of protons and neutrons in an atom o Isotopes- atoms of an element that have different numbers of neutrons AX Z = symbol for an isotope o Ion- atom with a charge Anion- negatively charged ion; atom has gained electrons Cation- positively charged ion; atom has lost electrons o Atomic mass- average of the masses of all isotopes of an element = (mass of isotope A)(fraction of isotope A)+(mass of B) (fraction of B)… o Polyatomic ion- group of atoms that behaves as a single particle (be familiar with chart on p. 36 of lab manual!) Ionic compounds- hard, brittle solids; high melting points o Dissociate (dissolve) o Electrolyte- forms ions when dissolved in water and can conduct electricity Molecular compounds- low melting points; often liquids or gases o Usually nonconducting Chapter 3: Equations, the mole and Chemical Formulas All chemical reactions must be balanced Important types of reactions o Neutralization- acid and base react to form a salt and water o Combustion reaction- something burns in oxygen to produce carbon dioxide and water Mole- standard measurement of atoms and compounds 6.022∗1023 o Avogadro’s number = Number of entities (atoms or molecules) in one mole of anything o Molar mass- mass of one mole of a substance Mass percentage- proportion of mass of one element in a compound o Divide the molar mass of one element in the compound by the total molar mass of the compound o If given mass percentage- multiply mass percentage (in decimal form) by the total mass of the compound Elemental analysis by combustion Calculating empirical formula- find the number of moles of each substance in the formula; divide each by the lowest number of moles; if resulting values are not integers, multiply by the denominator to get whole number values Calculating molecular formula- must be a whole number value based off of the empirical formula; use experimental mass to determine Stoichiometry- studies ratios between compounds in a chemical reaction o Guidelines BALANCE EQUATIONS Calculate moles based off of given mass Convert to moles and then grams of desired substance o Theoretical yield- quantity of product assuming all reactants are unlimited o Limiting reactant- reactant that is completely consumed in a reaction Actual yield- mass of a product that is produced based on the limiting reactant Chapter 4: Chemical Reactions in Solution Electrolytes o Strong electrolyte- separates completely into ions when dissolved in water o Weak electrolyte- partially ionizes when dissolved in water Soluble ionic compounds o Group 1A cations and ammonium o Nitrates o Perchlorates o Acetates o Chlorides, bromides, iodides (except with Ag, Hg, Pb) o Sulfates (except with Sr, Hg, Pb, Ba) Insoluble ionic compounds o Carbonates (except with group 1A cations and ammonium) o Phosphates (except with group 1A cations and ammonium) o Hydroxides (except with group 1A cations, ammonium, Sr, Ba) Precipitation reaction- a precipitate (solid) is formed Molarity (M)- number of moles of solute in 1 L of solution molsolute o M= Lsolution M V =M V Dilution- c c D D Chapter 5: Thermochemistry Study of the relationship between heat and chemical reactions Kinetic energy- energy possessed by objects in motion Law of conservation of energy- energy cannot be created or destroyed Exothermic- heat is transferred to surroundings o ΔH is negative Endothermic- heat is absorbed o ΔH is positive Specific heat- heat needed to raise 1g of a substance 1 degree C o q=mC∆T Hess’s Law- when multiple thermochemical equations are added, the enthalpy change is the sum of the added equations Standard enthalpy of formation o ∆ Hrxn ∑ n∆ H f[oducts ] ∑ m∆H [feactants] Chapter 6: The Gaseous State Gases have no fixed volume of definite shape Boyle’s Law- increasing pressure decreases volume of a gas P1V 1P V2 2 o Charles’s Law- increasing temperature increases the volume of a gas V 1= V2 o T T 1 2 Avogadro’s Law- equal volumes of gases at the same temperature and pressure contain the same number of particles V1 V2 o = n1 n2 Ideal gas law- PV = nRT o STP = 0 degrees C and 1 atm Molar mass and density- example Partial pressure- total pressure of combined gases is the sum of their partial pressures PT=P AP B o o Mole fraction- (χ) moles of one component divided by total number of moles Partial pressure = PA=χ AP T Collecting gases over water Kinetic molecular theory of gases- increasing temperature of gas increases the speed o Speed and molar mass have inverse relationship- as molar mass increases, speed decreases 3RT o Speed = urms √ molar mass Chapter 7: Electronic Structure Quantum numbers o Principle quantum number (n)- tells the energy and distance for the electron from the nucleus o Angular momentum quantum number (l)- tells the shape of the orbital Must be < n o Magnetic quantum number (m)- orientation of the orbital m = -l to l Shape of p orbitals- two lobes of electron density on opposite sides Shape of d orbitals- 4 lobes of high electron density Effective nuclear charge- average of nuclear charge affecting an electron o Electron shielding- influence of inner electrons on effective nuclear charge Increasing energy order of orbitals- p. 125 of lab manual o 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p Pauli exclusion principle- each electron in an atom must have a unique quantum number Aufbau principle- electrons are added to atoms one at a time and occupy the lowest energy orbital available Anomalous electron configurations Chapter 8: The Periodic Table and Trends The periodic table is organized into 4 sections based on the highest energy subshells of elements Valence orbitals- hold valence electrons (electrons with the highest quantum number residing in any unfilled subshell) Isoelectronic series- group of atoms and ions containing the same number of electrons Atomic/ionic radius- distance between the nucleus and the next element in the molecule o Cations are always smaller than their atoms; cations with higher positive charges are smaller o Anions are always larger than their atoms; anions with higher negative charges are larger Atomic size- atom size is determined by the effective nuclear charge on the outermost electrons in an atom o Increase in effective nuclear charge decrease in atomic size Ionization energy- energy required to remove an electron o Increase in effective nuclear charge increase in ionization energy o In an isoelectronic series, the atom/ion with the largest nuclear charge will have the largest ionization energy o 2 , 3 , … ionization energies dramatically increase Electron affinity- energy required to add an electron Chapter 9: Chemical Bonds Ionic bonding- bond resulting from the transfer of electrons between anions and cations o Lattice energy- energy required to separate an ionic solid into its ions Most strongly influenced by the charges on each ion in the compound- increase the charge, increase the ionization energy Also influenced by distance between the ions- decrease distance, increase ionization energy Covalent bonding- bond resulting from the sharing of electrons between atoms Lewis structures- represent covalent bonds showing how valence electrons are shared o Octet rule- each atom in a molecule must have a total of 8 electrons around it Bond polarity o Dipole movement- unequal sharing of electrons in a molecule One atom is more electronegative and possesses the electrons more than the other o Nonpolar- both atoms share the electrons equally Formal charge- charge assigned to an atom in a Lewis structure assuming shared electrons are divided equally between bonded atoms Resonance structures- multiple Lewis structures for the same molecule differing in how valence electrons are distributed o Dominant resonance structure is one in which the octet rule is satisfied for all atoms and there are few formal charges Bond energies- energy needed to break bonds in a molecule o ALWAYS endothermic o Enthalpy of reaction- ∆ Hrxn ∑ bondenergiesof bondsbroken∑ bondenergiesof bonds formed Chapter 10: Molecular Structure and Bonding Theories VSEPR- theory that predicts molecular shape will minimize repulsions between electron pairs o Electron pairs are as far apart as possible o Steric number = (# lone pairs on central atom) + (# atoms bonded to central atom) Determines bonded-atom lone-pair arrangement from VSEPR Molecular shape- arrangement of atoms in a molecule o Different from bonded-atom lone-pair arrangement o Steric numbers can have multiple possible molecular shapes Be familiar with charts on page 176 of lab manual! Polarity of molecules o Molecules with dipole (polar, unequal sharing of electrons) movement are polar o Molecules without dipole movement (equal sharing of electrons) are nonpolar Nonpolar molecules also experience equal vectors of electron sharing in each direction, which cancels out the unequal sharing of electrons Hybrid orbitals- orbitals resulting from mixing multiple atomic orbitals on one central atom o Be familiar with chart on p. 186 of lab manual Multiple bonds o Sigma bonds- single bond Directly between the nuclei of two atoms o Pi bond- found in double and triple bonds Found above and below the joining of the bonded atoms A double bond contains one sigma bond and one pi bond A triple bond contains one sigma bond and two pi bonds Homonuclear diatomic molecules o Bonding orbital- electron density is concentrated between atoms o Antibondig orbital- electron density is not between bonded atoms o Bond order = ½ (# electrons in bonding orbital) – (# electrons in antibonding orbital) o Electron configuration - (σ2s) (σ∗2s) (π2 p) (σ2 p) (π∗2 p) (σ∗2 p) 2 Chapter 11: Liquids and Solids Intermolecular attractions o Dipole-dipole attractions- intermolecular attractions between polar molecules o London Dispersion forces- attractions between ALL molecules due to induced and instantaneous dipoles o Hydrogen Bonding- occurs between a H atom in one molecule and a lone pair on an N,O, or F atom in another molecule VERY strong attractions o London dispersion forces < dipole-dipole < Hydrogen Bonding
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