Significant Figures (pg 36-45)   Math Review (pg 1213-1220)   Dimensional Analysis (pg 44-47)

Historical Experiments leading to the Atomic Theory (2.1-2.2) Heisenberg Uncertainty Principle Tells us we can't know both the location and velocity of an electron at the same time Atomic Structure and Composition (2.3) Particle Mass measured in kg and atomic mass units (amu) 1/12 mass of atomic Carbon unit (carbon atom) Charge Measured in Coloumbs and electronic charge Atomic Volume mostly comprised of electron cloud comprised of Electrons Atomic Mass mostly comprised of nucleus comprised of Protons and Neutrons https://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=2ahUKEwicjIaYwYTfAhVlmeAKHZ4nCH8QjRx6BAgBEAU&url=http%3A%2F%2Fguillerexploren.blogspot.com%2F2013%2F02%2Fmodelos-atomicos_27.html&psig=AOvVaw1fHv3pisRnJU0qOrbN14G9&ust=1543955650034867 Chemical Symbols and Sub-atomic Particles (2.3) The chemical symbol, as pictured in the image below shows the mass number and atomic number as well as the charge and atomic symbol https://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=2ahUKEwjitabEwoTfAhXIT98KHSBdBUgQjRx6BAgBEAU&url=http%3A%2F%2Fdata.allenai.org%2Ftqa%2Finside_the_atom_L_0772%2F&psig=AOvVaw0qCay2NvRbd1GC9TLuqGAx&ust=1543956021104034 atomic number is the number of protons ions (charged atoms) do not have equal numbers of protons and electrons they have the same amount of protons as the base atom but a different amount of neutrons and electrons cations - positively charged ions anions - negatively charged ions Isotopes (2.3) Isotope abundance and average mass average mass is equal to the sum of the quantity of the fractional abundance multiplied by the isotopic mass represented as mu mass to find the abundance, use the average mass formula and replace fractional abundance with x and 1-x Empirical/Molecular Formulas and Molecular Weight (2.4) molecules - combinations of different atoms bonded together each non-hydrogen seeks to have 8 valence electrons by sharing two electrons with another atom in a bond double bonds share 4 electrons triple bonds share 6 electrons Mass v. moles moles compare values of mass and/or volume for different molecules Avogadro's number - 6.02x10^23 "things"/mole   Review of grams-to-moles-to-(molecules or atoms) problems

Features of Waves (3.1) Electromagnetic radiation a wave that travels at the speed of light the electric field is the sine wave and is measured in voltage magnetic field is the cosine wave quantized waves fixed ends of stationary wave restrict number and types of waveforms only discrete values of a continuous set are allowed node - region where amplitude is zero no motion Energy, wavelength, frequency of a photon (3.1) wavelength goes in the opposite direction of frequency and energy frequency is represented by v wavelength is represented by lambda the visible light spectrum is 400-800nm blue is 400nm and red is 800nm https://www.cyberphysics.co.uk/graphics/diagrams/waves/EMSpectrumcolor.png one photon is equal to one light unit Energy is measured in joules Photoelectric effect (3.1) minimum threshold to eject electrons is related to the frequency of light classical physics suggested the amplitude of a wave influences the energy but experimentally, it was discovered that it is the frequency there is a certain amount of energy required to hold the atom together in order to eject an electron, you must overcome the threshold energy Orbit energy of e-, delta E, Bohr Model (3.1, 3.2) line spectra elements and molecules have emission spectra that correspond to a series of line each line is a single wavelength Rydberg equation (included in sheet) n refers to which orbit an electron might be in  1 over lambda https://i.stack.imgur.com/oDIES.jpg wave particle duality light behaves both wavelike and particle-like wavelike property - interference patterns obtained when light passed through a double slit particle like property - photoelectric effect Bohr's atomic theory Bohr used the results of Planck and Einstein's experiments to develop a better atomic theory, making the following assumptions about an atom's electrons electrons travel in circular orbit around the nucleus energy proportional to distance orbits are quantized equation is included on the sheet  E sub n Delta E Z is the nuclear charge energy is absorbed/emitted when orbits change Orbit transitions an electron moves to a higher level (n) when a specific quantity of energy is absorbed an electron moves to a lower level (n) when a specific quantity of energy is emitted the Bohr model fails with more than one electron, but was right about the following e- attracted to nucleus by the opposing charges of the particles e- energies are quantized  e- energies increase as the distance from the nucleus increases discrete energies in spectra of elements result from quantized energy transitions of e- between levels DeBroglie Wavelength (3.3) wavelength is equal to h over linear momentum linear momentum is equal to mass multiplied by velocity quantization of angular momentum, L, is equal to n multiplied by h wavelike properties are only important for small particles (electrons) Electron location and momentum Heisenberg Uncertainty Principle We can only know either the position or momentum of a particle at any given instant we can not know both in one instant consequence of wave-particle duality there are limitations governing the motions of quantum particles schrodinger expanded on de Broglie's work and developed a wave equation that described and electron as a wave function, represented by psi psi are not physical waves but complex probability amplitudes psi squared is equal to the probability distribution (e- distribution) Quantum Theory, Quantum Numbers (3.3) a series of four quantum numbers describe the energies, locations and other properties of an electron in a given atom  principle quantum number (n) discrete energy level energies increase as the distance from the nucleus increase atomic orbital which orbital corresponds to the row on the periodic table angular momentum quantum number (l) shape integers from 0 to n-1 s, p, d, and f orbitals are 0, 1, 2, and 3, respectively when l is not equal to 0, the orbitals have directionality about the z axis described by the magnetic quantum number magnetic quantum number (m sub l) orientation integers from -l to l number of degenerate orbitals in subshell spin quantum number (m sub s) electronic spin overall one magnitude, two quantized states  +1/2, -1/2 spin states degenerate unless in external magnetic field Orbitals and Wave functions (3.3) pauli exclusion principle no two electrons can have the same two sets of electrons they have different quantum numbers in some way or another radial nodes there are certain distances from the nucleus at which the probability density of finding an electron located at a particular orbital is zero the number of nodes is equal to the energy level minus l minus 1 Electron Configurations (3.4) the total number of electrons in an atom populate these orbitals following a set of rules the end result is an electron configuration or orbital diagram hund's rule electrons spread out within a sub shell and have the same spin filled and half filled subshells are especially stable  aufbau principle build up from lowest energy to highest valence electrons the electrons in the outermost principle quantum level of an atom (largest value of n) the 3de- are not counted as part of the valence shell full d and f shells are not considered part of the valence shell because they are not involved in reactions if they are full therefore, determination of the number of valence e- for any main group atom is simply the group number HOWEVER, atoms in the transition series do count the de- toward the valence shell Periodic Trends (3.5, 3.6) periodic law - the properties of the elements are periodic functions of their atomic numbers nuclear charge many of the periodic trends involving e- can be qualitatively explained using the effective nuclear charge (Z sub eff) of the atom this is the pull exerted on a specific electron by the nucleus, taking into account any electron-electron repulsions Z sub off is equal to Z minus the shielding  shielding is the e- repulsions radii covalent radius - one half the distance between the nuclei of two identical atoms when they are joined by a covalent bond ionic radius - size of an ion increases as you go down and to the left ionization energy - Energy required to remove the most loosely bound e- in formation of cation as size increases, IE decreases it becomes easier to remove electrons increases as you go up and to the right  watch for filled and half filled subshells electron affinity - energy change for adding an electron to gaseous atom in formation of an anion increases going up and to the right unit prefixes and conversions pico 10^-12 nano 10^-9 micro 10^-6 milli 10^-3 centi 10^-2 kilo 10^3 Ionic and Covalent Bonds, including polyatomic (3.7) monatomic ions the majority of chemical reactions do not change the nucleus of an atom  electrons can be transferred or shared groups alkali metals alkaline earth metals chalcogens halogens noble gases polyatomic ions usually have a negative charge binary bonding ionic metal and nonmetal groups of atoms held together by electrostatic interactions  transfer of electrons molecular/covalent nonmetal and nonmetal groups of atoms held together by sharing electrons

Ionic Compounds ionic bonding forms ionic compounds network lattice electrons are transferred between atoms of different elements to form ions NaCl covalent bonding forms molecules atoms make chemical bonds by sharing electrons between each other HCl a region of e- density between the nuclei arises from this sharing bond formation results in a release of energy Types of covalent bonds Pure/Non-polar covalent (equal sharing) EN 0-0.4 Polar covalent (unequal sharing) EN 0.4-1.8 ionic  EN >1.8 Bond Dipole Moments un-even sharing of electrons in a bond leads to polar bond and a dipole moment  Naming Compounds Ionic bonds cation + anion numbers of each ion implied by typical ion charges Multiple charged ionic (transition metals) cation(charge) + anion Molecular requires a different system due to enormous variation in the ratios in which atoms are combined more metallic + less metallic use prefixes to specify the numbers of atoms of each element Binary Acids H + nonmetal hydro + nonmetal + ic acid Oxyacids H + polyatomic anion anion + ic acid Drawing Lewis Structures Lewis structures are used to indicate the connectivity of atoms in a molecular compound shows the valence electrons Octet rule atoms like to form bonds to achieve 8 valence e- the number of bonds for main group atoms can be predicted from the number of valence e- required to achieve an octet Steps Determine the total number of valence electrons for cations, subtract one electron for each positive charge for anions, add one electron for each negative charge Draw a skeleton structure of the molecule/ion, with the least electronegative element in the center. Connect each atom to the central atom with a single bond (2 electrons) Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom place all remaining electrons on the central atom Minimize Formal Charge by rearranging electrons of the outer atoms to make multiple bonds with the central atom formal charge is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms number of valence electrons - number of non-bonding electrons - number of bonding electrons/2 have to satisfy the octet rule first Exceptions Odd e- molecules  free radicals usually groups of 3 or 5 Electron deficient molecules Groups 2 and 12 Hypervalent molecules n > 3 exceeds normal valence Resonance sometimes a single Lewis structure is not adequate to describe a molecule, with the other possibilities only differing by the location of e- (atoms connected the same) resonance structures when a bond can be represented multiple ways the delocalization of e- in Lewis structures each discrete Lewis structure for the same molecule that only differs by the locations of e- is called a resonance form VSEPR Theory and Molecular Shape Assumes that electron pairs in the valence shell of a central atom repel one another, and can be used to predict molecular shape e- pairs include bonding e- and lone pair e- https://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=2ahUKEwi2tZW4zJjfAhXxm-AKHYQ4CeYQjRx6BAgBEAU&url=https%3A%2F%2Fsites.google.com%2Fa%2Fcoe.edu%2Fcourtney-s-chemistry%2Fhome%2Fvsepr-theory&psig=AOvVaw2B50NRkJPUgLrwNTJtRnts&ust=1544645871316887 electron pairs repel one another Bond polarity and dipole moment polar covalent bonds have unequal distribution of e- in a bond, giving a bond dipole moment, represented by mu mu = Qr Q = magnitude of partial charge r = distance between charges mu units = D = Debye 3D shape must be considered to predict the molecular dipole moment sum up the vectors (direction and magnitude) for each bond dipole Sigma and Pi Bonds Valence Bond Theory (VBT) describes a covalent bond as the overlap of half-filled valence atomic orbitals that yield a pair of electrons shared between the two bonded atoms Bond Types sigma bonds single bond head on orbital overlap e- density concentrated along internuclear axis pi bonds double bonds side on orbital overlap e- density concentrated orthogonal to internuclear axis with p orbitals internuclear axis is a node double bonds composed of one sigma bond and one pi bond results from one p orbital and one hybridized orbital triple bonds composed of one sigma bond and two pi bonds results from two p orbitals and one hybridized orbital Hybridization with VSEPR Theory Hybridization of standard AOs creates a new hybrid orbitals that achieve a better picture of the experimental bonding Used in VBT to predict accurate bond angles and geometry sp3 hybridization results from average of one s and three p orbitals forms 109.5 degree bond angles (tetrahedral) identified by 4 things around central atom sp2 hybridization results from average of one s and two p orbitals forms 120 degree bond angles (trigonal planar) identified by 3 things around central atom sp hybridization results from average of one s and one p orbital forms 180 degree bond angles (linear) identified by 2 things around central atom   Hybridization with SPE

Molar Mass and the Mole molar mass = formula mass = molecular weight sum of the average atomic mass for each atom in chemical formula  Grams to Moles divide by molar mass Moles to Molecules to Atoms N from chemical formula Percent Composition the percent by mass of a particular atom in the chemical formula by mass mass A/mass compound by molar mass (atoms B x molar mass B)/(molar mass compound) Empirical, Molecular Formula A chemical formula gives the molar ratios of each atom in the molecule empirical formula mass to moles  divide by lowest amount of moles convert ratio to lowest whole numbers Molecular Formula empirical formula molar mass/empirical formula mass Molarity, Volume to Moles for solutions solutions are the most common form of a mixture defined by the quantity of a pure substance per volume of solution concentration = moles of solute/volume solution Dilutions M1V1 = M2V2

Balancing Reactions Ionic reactions and equations ionic compounds are held together as solids electrostatic interactions ionic compounds dissociate in water solvated ions move around independently in solution Solubility, Net ionic Reactions molecular equation displays discrete compunds complete ionic equations displays all individual ions in equation net ionic equation displays all individual ions in equation spectator ions do not directly participate in reactions precipitation reactions occurs when a non-soluble solid crashes out of solution steps identify reactant ions identify possible products ions that form solids have a net neutral charge most ionic materials only contain two types of ions one anion and one cation Solubility rules: https://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=2ahUKEwjiy7ij3JjfAhUNJt8KHYEwCIYQjRx6BAgBEAU&url=https%3A%2F%2Fsdsuchem200.com%2F2018%2F08%2F26%2Fsolubility-rules%2F&psig=AOvVaw0YesRjI4Cuk3SyHGZAfAVW&ust=1544650117466775 Acid/Base Reactions occur when a proton is transferred from one species to another Acids dissolved in water form hydronium strong acids completely form hydronium weak acids partially form hydronium Bases dissolved in water form hydroxide strong bases completely form hydroxide weak bases partially form hydroxide neutralization reaction specific acid-base reaction in which an acid reacts with a base REDOX Reactions and Balancing an abundance of reactions involve the transfer of electrons that are categorized as oxidation-reduction reactions oxidation - loss of electrons reduction - gain of electrons do not necessarily involve ions or ionic species oxidation states the charge its atoms would possess if the compound was ionic  Steps Oxidation state of an atom in an elemental substance is zero the oxidation state of a monatomic ion is equal to its charge Oxidation states for common non-metals are as follows: Hydrogen: +1 when combined with non-metals and -1 when combined with metals Oxygen: -2 in most compounds, sometimes -1 (peroxides), and -1 when combined with metals Halogens: -1 for F always, -1 for other halogens unless combined with O or other more EN halogens the sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion Characteristics of Redox reactions transfer or electrons sometimes forms ionic substances, other times it occurs in a formal sense oxidant/oxidizing agent = species that is reduced reductant/reducing agent = species that is oxidized half reactions separate the oxidation from the reduction for analysis purposes Balancing Redox Reactions Acidic conditions Steps Write separate equations for the oxidation and reduction half reactions For each half-reaction balance all elements except H and O Balance O using water Balance H using H+ Balance charge using e- Balance half-reaction charge, if necessary. Multiply one or both half-reactions by integers to equalize the number of e- transferred in the two half reactions Combine both half-reactions into a complete equation, then cancel identical species Finally, check to be sure the element and charges balance  Basic conditions same as acidic conditions with a few extra steps add the same amount of OH- as H+ on both sides Reaction Stoichiometry involves a quantitative assessment of the relationships between the amounts of substances consumed and produced by the reaction must use a balanced equation coefficients are used to derive stoichiometric factors https://www.google.com/url?sa=i&rct=j&q=&esrc=s&source=images&cd=&cad=rja&uact=8&ved=2ahUKEwj_96Kj4JjfAhVmUd8KHf9rBDoQjRx6BAgBEAU&url=https%3A%2F%2Fwww.onlinemathlearning.com%2Fstoichiometry.html&psig=AOvVaw2xY9g_KIEzOj53CaihngTs&ust=1544651199649109 Limiting Reagent the quantities of the products are limited by the reagent present in the least quantity Reaction Yield in Percent the limiting reagent for a given chemical reaction determines the maximum quantity of a product that is possible to form the reaction yield is an indication of the reaction's efficiency percent yield = actual yield/theoretical yield

Pressure Unit Conversions   Idea Gas Calculations and Ratios   Gas Density Problems   Partial Pressures of Gas Mixtures   Stoichiometry with Gases