CHEM 1030 FINAL STUDY GUIDE
CHEM 1030 FINAL STUDY GUIDE Chemistry 1030
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This 18 page Study Guide was uploaded by Emma Shoupe on Thursday April 28, 2016. The Study Guide belongs to Chemistry 1030 at Auburn University taught by Dr. Livia Streit in Spring 2016. Since its upload, it has received 132 views. For similar materials see General Chemistry 1 in Chemistry at Auburn University.
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Chemistry 1030 FINAL STUDY GUIDE Dr. Streit Exam 1 Chapter 1 Chemistry – study of matter and the changes that matter undergoes Matter- anything that has mass and occupies space Scientific method- o gather data via observations and experiment o identify patterns or trends in collected data o summarize findings with a law o formulate a hypothesis o with time a hypothesis might evolve into a theory substance- form of matter with definite composition and distinct properties mixture- physical combination of 2 or more substances o homogenous- uniform (ocean) o heterogenous- not uniform (ex- trail mix) states of matter- o solid o liquid o gas Quantitative – measured and derived with a number Qualitative – do not require measurement and are based on observation Physical property- one that can be observed and measure without changing the identity (ex. Color, boiling point) Chemical property- what a substance exhibits when it interacts with another substance Physical change- the state of matter changes, identity of the matter does not change (ex. Melting, freezing) Chemical change- a change in composition Atomic mass unit – used to express the masses of atoms and other similar sized objects o 1 amu = 1.6605378x 10−24 g mass – a measure of the amount of matter in an object or sample Celsius – freezing (0); boiling (100) Kelvin – “Absolute” scale; lowest possible temp. 0 K Know how to convert between Celsius and kelvin, Celsius and Fahrenheit Exact numbers: have defined values Inexact numbers: measured by any method other than counting Significant figures- meaningful digits in a reported number o The last digit in a measured number is referred to as the uncertain digit o RULES Any nonzero digit is significant Zeroes between nonzero digits are significant Zeroes to the left of the first nonzero digit are not significant Zeroes to the right of the last nonzero digit are significant if a decimal is present Zeroes to the right of the last nonzero digit in a number that does not contain a decimal point may or may not be significant o When adding and subtracting, the answer cannot have more digits to the right of the decimal point than any of the original numbers o when multiplying and dividing, the number of sig figs in the final product or quotient is determined by the original number that has the smallest number of sig figs accuracy- how close a measurement is to the true value precision- how close a series of replicate measurements are to one another Chapter 2 atoms- smallest quantity of matter that still retains the properties of matter element- a substance that cannot be broken down into 2 or more simpler substances like charges repel, opposite charges attract electrons- rays of negatively charged particles Types of radiation o Alpha- positively charged particles o Beta- electrons deflected away from negative plate o Gamma- no charge and are unaffected by external electric or magnetic fields Rutherford used alpha particles to prove structure of atoms o Protons account for most of the mass of the atom o Neutrons- neutral particles, slightly larger than protons All atoms can be identified by the number of protons and neutrons they contain Atomic number- number of protons in nucleus Mass number- total number of protons and neutrons Isotopes- atoms with the same atomic number, but different mass numbers Atomic mass- mass of an atom in atomic mass units Average atomic mass- average mass of naturally occurring mixture of isotopes Periodic table- a chart in which elements having similar chemical and physical properties grouped together o Metals- good conductors of heat and electricity o Nonmetals- poor conductors of heat and electricity o Metalloids- intermediate properties Mole- defined as the amount of a substance that contains as many elementary entities as there are atoms in exactly 12g of Carbon -12 (for example) Avagadro’s number- 6.022x10^23 Molar mass- mass in grams of one mole of the substance o In units of grams/mole Chapter 3 Joule (J) – amount of energy possessed by a 2kg mass moving at a speed of 1 m/s Electromagnetic spectrum- visible light is only a small component of the continuum of radiant energy Wavelength- distance between two identical points on successive waves Frequency- the number of waves that pass through a particular point in one second Amplitude- vertical distance from the midline of a wave to the top or bottom Speed of light- c= 3.00x10^8 m/s = frequency x wavelength E= hv; h= constant 6.63 x 10^-34 J Schrodinger equation- particle behavior (mass, m), wave behavior (wave function, w- this is the weird symbol pictured, they don’t have it on word!) Quantum mechanics- defines the region where the electron is most likely to be at a given time Energy states and wave functions are characterized by a set of quantum numbers Quantum numbers and wave functions describe atomic orbitals o Principal quantum number (n) – designates size o Angular quantum number (l) – describes shape o Magnetic quantum number (ml) – specifies orientation Electron configuration- describes how electrons are distributed in various atomic orbitals o Ground state- lower level of energy o Excited state- higher level of energy Pauli exclusion principle- no 2 electrons in an atom can have the same 4 quantum numbers Aufbau principle- states that electrons are added to the lowest energy orbitals first before moving to higher energy orbitals Hund’s rule- most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized Paramagnetic- not all electrons are paired (respond to magnetic field) Diamagnetic- all electrons are paired (does not respond to magnetic field) GENERAL RULES OF ELECTRON CONFIGURATION o Electrons will reside in the available orbitals in the lowest possible energy o Each orbital can accommodate 2 electrons o Electrons will not pair in degenerate orbitals if an empty orbital is available o Orbitals will fill in the order Exam 2 Chapter 4 Effective nuclear charge- actual magnitude of positive charge that is “experienced” by an electron in the atom Left to right- gain electrons Shielding- partially shielded from positive charge by other electrons Atomic radius- distance between nucleus of an atom and valence shell Metallic radius- half distance between nuclei of 2 identical metal atoms Covalent radius- half distance between adjacent nuclei connected Atomic radius increases from top to bottom o Decreases left to right Ionization energy- minimum energy required to remove an electron from an atom in the gas phase Results in ion – chemical species with a net charge Cation – positive charge (means loss of electron) 1 ionization energy – removal of the most loosely held electron Ionization Energy: the amount of energy required to remove an + - electron from the ground state of a gaseous atom or ion.A (g) A +e Bigger going right, smaller going down Exceptions: between Group 2&13, Group 5&6 Electron affinity- energy released when an atom in the gas phase accepts an electron Results in an ion Anion – negative charge (means gain of electron) o Easier to add electron to an s orbital than to a p orbital o Electron affinity increases left to right Easier to add an electron as the positive charge of a nucleus increases Electronegativity: tendency to attract electrons in a covalent bond Increases going up and right Metallic character o Metals Shiny, lustrous, malleable, ductile Good conductors of heat and electricity Low ionization energies (form cations) o Nonmetals Vary in color, not shiny Brittle Poor conductors of heat and electricity High electron affinity (form anions) o Metalloids Elements with properties intermediate in between metals and nonmetals o Metals: tend to form cations, metal oxides are basic o Nonmetals: tend to form anions, nonmetal oxides are acidic, poor conductors of electricity o Metallic character increases down a group, decreases across a period o Alkali metals (1A)—The most reactive metal family, these must be stored under oil because they react violently with water! They dissolve and create an alkaline, or basic, solution, hence their name. o Alkaline earth metals (2A)—These also are reactive metals, but they don’t explode in water; pastes of these are used in batteries. o Halogens (7A)—Known as the “salt formers,” they are used in modern lighting and always exist as diatomic molecules in their elemental form. o Noble gases (8A)—Known for their extremely slow reactivity, these were once thought to never react; neon, one of the noble gases, is used to make bright signs. Isoelectric – species with identical electron configurations to the noble gas to the right High Electron affinity easy to accept electrons (+) Low electron affinity hard to accept electrons (-) Low ionization energy easy to form cations High ionization energy hard to lose electrons Ionic Radius o The radius of a cation or anion o When an atom loses an electron and becomes a cation, the radius DECREASES due in part to a reduction in electron-electron repulsions in the valence shell and when all of an atoms valence electrons are removed Comparing Ionic radius with Atomic radius o When an electron gains 1 or more electrons and becomes an anion, its radius INCREASES due to an increase in electron-electron repulsions o Isoelectric series- a series of 2 or more species that have identical electron configurations, but different nuclear charges Chapter 5 Compound- composed of 2 or more elements combined in a specific ratio and held together by covalent bonds o ex- water and salt (sodium chloride) 1:1 ratio Lewis dot symbols o When atoms form compounds, it is their valence electrons that actually interact o Consists of the element’s symbol with dots o For main group elements, such as Na, the number of dots is the number of electrons that are lost o For non metals in the 2 ndperiod, the number of unpaired dots is the number of bonds the atom can form Ionic Bonding- Electrostatic attraction that holds oppositely charged ions together in an ionic compound (salts) o Transfer of electrons o Generally a metal with a nonmetal o Gain or lose electrons to fill an octet o Form large crystalline solids where each cation is surrounded by an anion, etc o High melting points o Do not conduct electricity in solid state Electrons cannot flow in a rigid structure, and electricity is defined as the flow of charged particles o Brittle o Mostly soluble in water o Measured with Coulomb’s law Energy needed to dissociate ions Lattice energy- measures ionic stability (depends on magnitudes of charge and distance) The greater the product of charges of ions in a compound, the greater the attractive forces, therefore, the higher the melting point The bigger the ions, the greater the distance between ions, therefore, the attractive forces are weaker thus lowering the melting point o Chemical formula Denotes the constituent elements and the ratio in which they combine o Bigger the difference in electronegativity the more ionic the compound Naming Ions Monatomic cation o Named by adding the word ion to the name of the element Monatomic anion o Named by changing the ending of the element’s name to –ide Some metals can form cations of more than one possible charge o Fe (II) or Fe (III) n+¿ X¿ - ous n+1 )+¿ -ic X ¿ X -ine or –ide Formulas Sum of charges on the cation and anion in each formula must be zero 1. Name cation a. Omit ‘ion’ b. Use Roman numeral if cation can have more than one charge 2. Name the anion a. Omit ‘ion’ NaBr sodium bromide OR FeCl2 iron (II) chloride o Covalent Bonding- Sharing of electrons between atoms which results in a more stable electron configuration A filled octet Bonds atoms together to form molecules o Single covalent bond Mutual affinity Electrons equally pulled Sharing of 1 pair of electrons between 2 atoms Molecule o Combination of at least 2 atoms in a specific arrangement held together by chemical bonds Element or compound Law of definite proportions o Ratio of masses of one element that combine with a fixed mass of other elements can be expressed in small whole numbers Diatomic molecules o Contain 2 atoms and may be heteronuclear or homonuclear Polyatomic molecules o More than 2 atoms Molecular Formula Shows the exact number of atoms of each element in a molecule Allotropes o Have 2 or more distinct forms Structural formula o Shows elemental composition and general arrangements Empirical formulas o Whole number ratio of elements Simplest formula Glucose – C 6 O12 6 o 1 C: 2 H: 1 O Naming Molecular Compounds o Binary molecule compounds Substances that consist of two different elements Name 1 element that appears in the formula nd Name 2 element, changing its ending to –ide o HCl – hydrogen chloride Greek prefixes 1 Mono- 6 Hexa- 2 Di- 7 Hepta- 3 Tri- 8 Octa- 4 Tetra- 9 Nona- 5 Penta- 10 Deca- Examples o CO – carbon monoxide o CO - carbon dioxide 2 st mono – omitted for 1 element Compounds containing Hydrogen Do not usually conform to systematic Nomenclature guidelines o Common names (water) or… Acid NH 3 – ammonia H 2 - water PH – phosphine 3 Acid o a substance that produces hydrogen ions (H+) when dissolved in water HCl Remove –gen ending from hydrogen Change the –ide ending on the 2 ndelement to –ic Only works for binary compounds o Ionizable hydrogen atom Compound MUST contain; to be an acid upon dissolving Organic compounds Carbon and hydrogen Hydrocarbons contain only C & H o Simplest = alkanes 1 C – meth 2 C – eth 3 C – prop 4 C – but 5 C – penta functional groups – molecule’s reactivity o polyatomic ions 2 or more atoms neutral PO ¿ 4 2 2−¿=Ca ¿3 2+¿PO ¿ ¿ 4 Ca memorize the polyatomic ions!! Oxoanions o Polyatomic anions with 1 or more oxygen atoms and 1 atom (central) o –ate one or more oxygen atom than –ate ion is called “per” one less oxygen atom than –ate is called “-ite” Oxoacids o Producing hydrogen ions and corresponding oxoanions o –ate = -ic o –ite = ous o per (one more) o hypo (lowest) can be monoprotic (1 ionizable H) or polyprotic (more than 1 ionizable H) Hydrates A compound that has a specific number of water molecules within its solid structure o Example copper(II) sulfate has 5 water molecules copper(II) sulfate pentahydrate CuSo 45H O2 when water is driven off by heat, copper(II) sulfate becomes an anhydrous copper(II) sulfate with water = blue without water = white Molecular Mass the mass in atomic mass units of an individual molecule o multiply atomic mass for each element in the molecule by the number of atoms of that element then total masses average molecular weight = average amu on periodic table for ions o use empirical formula to find formula mass (formula weight) o same process as molecule Percent Composition a list of the percent by mass of each element in a compound empirical formula o mass in amu of one empirical formula Molar Mass mass in grams of 1 mole of substance o numerically equal to its atomic mass (element) compound- sum of molar masses of elements it contains 5.15 example in book Chapter 6 6.1 Octet Rule atoms will lose, gain, or share electrons in order to achieve a noble gas electron configuration only 2 valence electrons contribute to bonding pairs of valence electrons not involved in bonding are lone pairs Lewis structure o Representation of covalent bonding o can be expressed with 2 dots or with a dash ( - ) Single bond o Sharing 1 pair of electrons between 2 atoms Double bond o Sharing 2 pairs of electrons between 2 atoms Triple bond o Sharing 3 pairs of electrons between 2 atoms Bond length o Distance between nuclei of 2 covalently bonded atoms o The stronger the bond, the shorter the bond o Triple bond is always shorter than double bonds o Double bond is always shorter than single bonds We quantify bond strength by measuring the quantity of energy required to break it 6.2 Electronegativity Covalent bonds occur between atoms that share electrons Ionic bonds occur between metal and nonmetal with ions Polar covalent bonds o Electrons shared but not shared equally Electrons “hang out” around the more electronegative element o A polar bond: Must have at least 1 polar bond (F, O, N, Cl) or… Must be asymmetrically shaped a lone pair on the central atom or… all bonding atoms are not the same Electronegativity o Ability of an atom in a compound to draw electrons to itself Varies with atomic numbers Nonpolar o Electronegativity differs by less than 0.5 Polar o Electronegativity differs by 0.5 – 2.0 Ionic o Electronegativity differs by 2.0 or more Dipole moment o Direction of electrons shift o Quantitative measure of polarity Exam 3 Chapter 7 Valence Shell Electron Pair Repulsion (VSEPR) – predicting molecular shape; basic idea is that electrons repel each other; electrons are found in different domains (lone pairs/single bonds/double bonds/triple bonds) o electrons will arrange themselves as far as possible o arrangements minimize repulsive interactions o make sure to know the different types and bond angles of molecular geometry Electron domain geometry- arrangement of electron domains around central atom Bond angle- angle between two adjacent A-B bonds 4 Steps To determine geometry- o Draw lewis structure o Count the number of electron domains around the central atom o Determine electron-domain geometry by applying VSEPR model o Determine molecular geometry by considering the positions of the atoms only Lone pairs take up more space than bonded pairs of electrons Van der Waal’s Forces o London forces – natural attraction between all molecules Increases with molar mass Size of electron cloud determines strength Similar molecular weight = same London force strength Weakest of all forces Only type between nonpolar molecules o Dipole-dipole forces- attraction between oppositely charged portions of 2 or more polar molecules The more polar, the stronger the attraction forces (higher melting and boiling point) Stronger the London forces Bigger the dipole moment, the more polar the molecule Hydrogen bonds – special type of dipole attraction where a hydrogen gets trapped between two highly electronegative elements (F, O, N) o Both molecules MUST have a hydrogen directly attached to a F, O, or N. o Strongest attraction of the 3 intermolecular forces Dispersion forces – result from Coulombic attractions between instantaneous dipoles of non-polar molecules Valence Bond Theory – atoms share electrons when atomic orbitals overlap o the H-H bond in H2 forms when the singly occupied 1s orbitals of the 2 H atoms overlap o A bond forms when single occupied atomic orbitals on 2 atoms overlap o The 2 electrons shared in the region of orbital overlap must be opposite spin o Formation of a bond results in a lower potential energy for the system Hybridization- accounts for observed bond angles in molecules that could not be described by the direct overlap of atomic orbitals o If the s orbital and 3 p orbitals hybridize, then it is sp3 hybridization Example – methane CH4 Steps to determine hybridization- o Draw lewis structure o Count electron domains on the central atom. This will be equal to the number of hybrid orbitals o Draw the ground state orbital diagram for the central atom o Maximize number of unpaired valence electrons by promotion o Combine the necessary number of atomic orbitals to generate required number of hybrid orbitals o Place electrons in hybrid orbitals, putting one electron in each orbital before pairing any electrons Sigma bond – forms when sp2 hybrid orbitals overlap o 1 bond between ANY 2 atoms is ALWAYS a sigma bond Pi bond – when the overlap of the orbitals does NOT lie on a line drawn between the 2 nuclei o Only p onditalsrdan form pi bonds o The 2 or 3 bonds of a double or triple bond are always pi bonds o Pi bonds are not as strong as sigma bonds because the orbitals do not overlap as much in a pi bond Chapter 8 For chemical reactions, each species on the left is a reactant Each species on the right is a product (g) – gas; (l) – liquid; (aq) – aqueous; (s) – solid Equations must be balanced so the law of conservation of mass is obeyed o Achieved by writing stoichiometric coefficients to the left of the chemical formulas o Steps for balancing Change coefficients of compounds before changing the coefficients of elements Treat polyatomic ions that appear on both sides of the equation as units Count atoms or polyatomic ions carefully Combination reaction – 2 or more reactants combine to form a single product Decomposition – 2 or more products form from a single reactant Combustion – substance burns in presence of oxygen o ALWAYS produces CO2 and H2O o Incomplete if CO or C is produced A 1.50 g sample of hydrocarbon undergoes complete combustion to produce 4.40 g of CO 2nd 2.70 g of H O.2What is the empirical formula of this compound? A 0.250 g sample of hydrocarbon undergoes complete combustion to produce 0.845 g of CO a2d 0.173 g of H O. 2hat is the empirical formula of this compound? Know how to use the mole ratios the determine how much product will form from a balanced equation Limiting reactant – the reactant that is used up first o Excess reactants are present in quantities greater than necessary to react with the quantity of the limiting reactant o A 2.00 g sample of ammonia is mixed with 4.00 g of oxygen. Which is the limiting reactant and how much excess reactant remains after the reaction has stopped? o try example 8.7 in chemistry book Theoretical yield – amount of product that forms when all of the limiting reactant reacts to form the desired product Actual yield – amount of product actually determined from the reaction Percent yield – tells what percentage the actual yield is of the theoretical yield Chapter 9 Solution – homogenous mixture of 2 or more substances Solvent – substance present in largest amount Solute – other substances present Electrolyte – a substance that dissolves in water to yield a solution that conducts electricity Dissociation – electrolyte breaks apart into its constituent ions Ionization – a molecular compound forms ions when it dissolves Nonelectrolyte – a substance that dissolves in water to yield a solution that does not conduct electricity Strong electrolyte – dissociates completely o Example- strong acids (HCl), strong bases (NaOH) Weak electrolyte – a compound that produces ions upon dissolving but exists in solution predominantly as molecules that are not ionized Precipitate – insoluble product that separates from a solution Hydration – occurs when water molecules remove the individual ions from an ionic solid surrounding them so the substances dissolves Solubility – maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature Double replacement/metathesis – reactions in which cations in 2 ionic compounds exchange anions Ionic equation – compounds that exist completely as ions in solution are represented as those ions Net ionic equation – an equation that includes only the species that are involved in the reaction o Ions that appear on both sides are called spectator ions and are not included in the overall reaction Oxidation/reduction – chemical reaction in which electrons are transferred from one reactant to another (also called redox) o Oxidation = LOSS o Reduction = GAIN o OIL RIG or LEO goes GER o (oxidation Is loss, reduction is gain); (loss equals oxidation, gain equals reduction) redox – sum of an oxidation half-reaction and a reduction half reaction oxidation number – charge an atom would have if electrons were transferred completely o elements have an oxidation number of zero o steps to determine oxidation number start with the ones you know the total contribution to charge must sum to zero if it is a neutral compound; if not, it must sum to the charge of the compound. Displacement reaction – an atom or an ion in a compound is replaced by an atom of another element Molarity (M) – molar concentration; defined as the number of moles of solute per liter of solution o M = mol/L Dilution – process of preparing a less concentrated solution from a more concentrated one Moles of solute before dilution = moles of solute after dilution McVc = MdVd Chapter 10 System – a part of the universe that is of specific interest Surroundings – rest of the universe outside of the system Thermochemistry – study of heat in chemical reactions Heat – transfer of thermal energy o Either absorbed or released o Si unit is a Joule, J Exothermic process – occurs when heat is transferred from the system to the surroundings Endothermic process – occurs when heat is transferred from the surroundings to the system Thermodynamics – study of the interconversion of heat and other kinds of energy 3 types of systems o open system – can exchange mass and energy with the surroundings o closed system – allows the transfer of energy but not mass o isolated system – does not exchange either mass or energy with its surroundings state functions – properties that are determined by the state of the system, regardless of how that condition was achieved o pressure o volume o energy o temperature First law of thermodynamics – energy can be converted from one form to another, but not created nor destroyed Pressure increases when volume is constant Volume increases when pressure is constant Enthalpy of reaction – the difference between the enthalpies of the products and the enthalpies of reactants Chapter 11 A sample of a gas assumes both the shape and the volume of a container Gases are compressible Densities are much smaller than liquids and solids and are variable depending on the temperature and pressure Gases are always homogenous mixtures with other gases Pressure- force applied per unit o Pressure = Force/Area o 1 Pa = 1 N/m squared o common standard pressure units- 1 atm = 760 mmHg = 760 torr = 101.3 kPa Boyle’s Law- constant temperature o Pressure decreases, volume increases o Inverse relationship o Ex- balloon in a vacuum will expand o V = 1/P o P1 x V1 = P2 x V2 Charles and Gay-Lussac’s Law- states that the volume of a gas maintained at constant pressure is directly proportional to the absolute temperature of the gas o Example- heat a balloon and it expands. Cool a balloon and it shrinks. Avagadro’s Law – states that the volume of a sample of gas is directly proportional to the number of moles in the sample at a constant temperature and pressure Combined gas law- used to solve problems where any or all of the variables changes o look for variables you have and don’t have gas laws can be combined into a general equation that describes the physical behavior of all gases o R is called the gas constant…R = 0.08206 L·atm/mol·K o n= number of moles of gas o Be careful of units! (P= atm, V= Liters, T=Kelvin) PV = nRT
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