Chem 211 EXAM 1 study guide PDF

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Chapter 1 1. Define chemistry, matter, composition and property a. Chemistry = 1) the study of matter and its properties, 2) the changes that matter undergoes, and 3) the energy associated w/the changes b. Matter = anything that has BOTH mass and volume c. Composition = the TYPES and AMOUNTS of simpler substances that make up matter d. Property = characteristics that give a substance uniqueness 2. What are the states of matter and their relation to volume and shape? a. Solid i. Solid shape ii. Solid volume iii. Particles are close and organized b. Liquid i. Varying shape; conforms to shape of container ii. Upper surface iii. Particles are close but disorganized c. Gas i. No fixed shape ii. No fixed volume iii. No surface iv. Particles are far apart and disorganized 3. What’s the difference between a physical and chemical property? Relation between temperature and change of state a. Physical i. Show by itself w/o interaction w/other substances ii. Color, melting/freezing/boiling point, density, attraction/repulsion, opacity, viscosity iii. Composition doesn’t change iv. Changes of state = ex of physical change and are reversible through temperature b. Chemical i. Show w/interaction or transforms into other substances ii. Flammability, corrosiveness, reactivity w/water, pH iii. Composition changes iv. Irreversible 4. What is energy and the diff types? a. Energy = ability to do work i. Potential = energy due to position ii. Kinetic = energy due to movement iii. Total E = P + K b. Lower energy states = more stable and favored c. Energy is neither created nor destroyed; but can be transformed d. Energy is conserved when it’s transformed i. Stable when ball falls, spring relaxed, particles attracted together 5. What are the steps of the scientific method? a. Observations hypothesis experimentModel (theory)Further experiments i. Observation natural phenomena and measured events ii. Hypothesis = Tentative proposal for observations iii. Procedure to test hypothesis; measure 1 variable at time iv. Model theory = set of conceptual assumptions that explains data from experiments; predicts related phenomena v. Tests predictions based on model

6. SI base units a. Mass – kilogram – kg b. Length – meter – m c. Time – second – s d. Temperature – kelvin – K e. Amount of substance – mole – mol f. Electric current – ampere – A g. Luminous intensity – candela – cd 7. Common decimal prefixes a. Tera – T – 1,000,000,000,000 – 1x1012 b. Giga – G – 1,000,000,000 – 1x109 c. Mega – M – 1,000,000 – 1x106 d. Kilo – k – 1000 – 1x103 e. Hecto – h – 100 – 1x102 f. Deka – da – 10 – 1x101 g. ------Base – 1 – 1x101--------h. Deci – d – 0.1 – 1x10-1 i. Centi – c – 0.01 – 1x10-2 j. Mili – m – 0.001 – 1x10-3 k. Micro – μ – 0.000001 – 1x10-6 l. Nano – n – 0.00000001 – 1x10-9 m. Pico – p – 0.000000000001 – 1x10-12 n. Femto – f – 0.000000000000001 – 1x10-15 8. SI – English equivalents a. Length i. 1km = 1000m (103) OR .6214 mile 1. 1mi = 1.609km = 1609m ii. 1m = 100cm (102) OR 1000mm (10-3) OR 1.094yd & 39.37in 1. 1yd = 0.9144m 2. 1ft = 0.3048m iii. 1cm = 0.01m (10-2) OR 0.3937in 1. 1in =2.54cm b. Volume i. 1 cubic meter (m3) = 1,000,000 (106) cubic cm (cm3) OR 1000 dm3 OR 35.31 ft3 1. 1ft3 = 0.02832m3 ii. 1 cubic decimeter (dm3) = 1000 cm3 (AKA 1000mL = 1L) OR 0.2642 gal OR 1.057 qt 1. 1 gal = 3.785 dm3 2. 1 qt = 0.9464 dm3 OR 946.6 cm3 iii. 1cm3 = 1000mm3 = 1mL = 1000μL OR 0.001 dm3 OR .03381 oz 1. 1oz = 29.57 cm3 iv. 1mm3 = 1μL c. Weight i. 1kg = 1000g = 2.205lb 1. 1lb = .4536kg 9. What is a conversion factor? a. Conversion factor = ratio of equivalent quantities used to express a quantity in diff units 10. What is density? a. D = m/V b. m= DV c. V= D/m

11. What are the diff types of temp scales and conversion? a. Kelvin i. Begins at absolute 0 ii. Only positive values iii. No degree sign iv. T (in K) =°C+273.15 b. Celsius i. Based on freezing and boiling points of water ii. Same size degree as Kelvin iii. T (in °C) = K−273.1 iv. T (in °C) = [℉−32](

5 ) 9

c. Fahrenheit i. Diff degree size and 0 points ii. T (in °F) = (

9 )*(°C)+32 5

d. Celsius and Fahrenheit read the same at -40 degrees 12. What are extensive and intensive properties? a. Extensive = dependent on amount of substance present i. Mass, volume b. Intensive = doesn’t depend on amount of substance present i. Density, molar point 13. Sig Fig rules a. Rightmost digit is always estimated in a measurement b. Greater # of sig figs = # certainty c. All digits (1-9) except 0 are significant d. 0’s at end are significant IF there’s a decimal point e. Starting 0’s aren’t significant f. Multiplication and division i. Contains same # of SF as the measurement w/ the fewest SF g. Addition and subtraction i. Contains same # of SF as measurement w/fewest decimal places 14. Rounding rules a. If digit removed… i. More than 5  preceding number inc by 1 ii. Less than 5 – preceding number doesn’t change iii. Is 5, followed by 0’s or no following digits, the preceding number 1. If preceding number is odd  Inc by 1 2. If preceding number is even  don’t change 3. EX: 17.7517.8; 17.65 17.6 iv. Is 5, followed by digits (1-9)  preceding number inc by 1 1. EX: 17.6500  17.6; 17.6513  17.7 15. Exact #’s a. Exact numbers = have no uncertainty associated i. They don’t limit (no effect) the # of SF in calculation b. By definition i. 1000mg = 1g ii. 60min = 1hr iii. 2.54cm = 1in c. By count i. 26 letters in alphabet

16. Precision, Accuracy and error a. Precision = how close measurements are to each other b. Accuracy = how close to the correct value c. Systematic error = makes all values higher/lower than the actual value d. Random error = makes both higher and lower values than the actual value .450 Chapter 2 1. What’s an element, molecule, compound, mixture? a. Element = simplest type of substance i. Unique physical and chemical properties ii. Only 1 type of atom iii. Can’t be broken down by physical/chemical matters b. Molecule = smallest independent unit i. Made from 2 or more atoms that are chemically bonded ii. Behaves as an independent unit c. Compound = 2 or more diff atoms chemically bonded d. Mixture = groups of 2+ elements and/or compounds that are physically intermingled 2. Defining features of compounds a. Elements are present in a fixed mass ratio b. Properties of a compound are diff from its elements c. Can be broken down into simpler substances via chemical change 3. Laws a. Conservation = Total mass before RX = total mass after RX b. Definite composition = no matter the source, a particular compound is composed of the same elements in the same parts (fractions) by mass i. EX: seashells, marble, coral OR tap sea, and rain water c. Multiple proportions = if elements A and B  2 compounds, the diff masses of B that combine w/a fixed mass of A can be expressed as a ratio of small whole numbers i. B mass changes, A mass remains the same ii. EX: 1. Carbon oxide I: 57.1% O 42.9% C  57.1/42.9 =1.33 2. Carbon oxide II: 72.7% O 27.3% C  72.7/27.3 = 2.66 i. 2.66/1.33 =2:1 4. Dalton’s atomic theory postulates a. All matter is made of atoms i. Particles of an element can’t be destroyed/created b. Atoms of an element can’t be converted into atoms of another element c. Atoms of an element are identical in mass and chemical properties but different to other element atoms d. Compounds result from the chemical combination of a specific ratio of atoms of diff elements e. Mass conservation = 1+2 f. Definite composition = 3+4 g. Multiple proportions = 3+4

5. Scientists a. Discovery of the electron i. JJ Thompson ii. Ray bends in magnetic field  charged particles iii. Ray bends to (+) plate  negative particles iv. Ray is identical for any cathode  particles found in all matter b. Measuring electrons charge i. Robert Millikan c. Calculate electron mass i. JJ Thompson 1. m/e = -5.686*10-12 kg/C ii. Millikan 1. e = -1.602*10-19C iii. m = (

m )*e = 9.109*10-31kg = 9.109*10-28g e

d. Discovery of nucleus i. Rutherford 1. Alpha particles carry (+) charge hits gold foil 2. Most pass through = didn’t hit the nucleus 3. Some reflected w/minor angle = hit nucleus slightly 4. Some deflected w/larger angle = hit nucleus a lot ii. Feats of the atom: 1. Electrically neutral 2. Spherical w/(+) nucleus surrounded by (-) e’s 3. Atomic nucleus made of protons (+) and neutrons (0) 6. Atomic symbol, # and Mass a. AZX i. A = mass number (p+n) ii. Z = atomic number (p) iii. X = atomic symbol b. Isotopes = same # of protons (atomic #), diff # of neutrons (mass number) i. 23592U vs 23892U 7. Calculating atomic mass from isotopes a. Mass of isotope (amu)* %abundance b. Add isotopic portions 8. Calculating Average atomic mass of isotopes a. avg atomic mass of element = (atomic mass of isotope #1) (x) 9. the a.

+ (atomic mass of isotope #2) (1-x) Periodic table discovery. Know where the metals, metalloids, transition metals and non-metals are on periodic table Dimitri Mendeleev i. Organized table by increasing atomic mass

ii. Elements w/similar chemical properties fell in the same column iii. Modern table 1. Arranged by increasing atomic number

10. Compound Bonding a. Majority of elements occur in compounds w/other elements i. Only a few occur free in nature 1. Noble gases, N2, O2, S8, Cu, Ag, Au, Pt b. Ionic bonding (e transfer) i. Non metal and metal ii. Lose e (+) iii. Gain e (-) iv. Factors that influence strength 1. Size (inc size [radius]  dec energy 2. Charge (inc charge  inc energy) c. Covalent bonding (e sharing) i. Usually between non-metals ii. EX: H2 = form equal attraction and repulsion 11. What is a molecule? Ion? Elements that occur as molecules? a. Molecule = basic molecular element or covalent compound, i. Consists of 2 or more atoms bonded by sharing e’s (covalent) ii. Elements as molecules 1. Diatomic = H2, N2, O2, F2, Cl2, Br2, I2 2. Tetratomic = P4 3. Octatomic = S8, Se8 b. Ion = A single atom or covalently bonded group of atoms that have an overall electrical charge i. No molecules in an ionic compound ii. Polyatomic ion = 2 or more atoms covalently bonded w/an overall charge 1. EX: Carbonate ion (CO3)212. What’s a cation and anion? Calculate oxidation state? a. Cation (positive charge) i. Group # b. Cation (negative charge) i. Group # – 8 13. What do chemical formulas show? a. The show the type and number of each atom present in the smallest unit of a substance 14. How do you name binary ionic compounds? a. Cation + (anion + -ide) i. Cation stays the metal name ii. Anion adds ide to nonmetal name 15. How do you name binary covalent compounds? a. 2 non-metals b. Lower group # element + (higher group # element + ide) i. Both parts use prefixes to indicate the # of atoms

16. Know the monatomic ions, metal monatomic ions, and common polyatomic ions

17. How do you name oxoanions? a. Oxoanion = Oxygen bonded to another nonmetal b. ClO- = hypochlorite (hypo+root+ite) c. ClO2- = Chlorite (root+ite) d. ClO3- = chlorate (root+ate) ClO4- = perchlorate (per+root+ate)

18.

How do you know acids? a. Dissolved in water acids i. Hydro + anion (nonmetal) name + ic + acid b. Oxoacid acids i. Suffix changes; prefixes stay 1. -ate  ic a. BRO4 = perbromate; HBRO4 = perbromic acid

2. -ite  ous a. IO2 = iodite; HIO2 = iodous acid 19. What are the prefixes for hydrates and binary covalent compounds? a. Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca 20. How do you name alkane chains? a. Alkanes = hydrocarbons = contain Carbon and Hydrogen i. Simplest hydrocarbon ii. –ane suffix iii. Formula = CnH(2n+2) 21. How to calculate molecular mass from chemical formulas (applies to covalent compounds) a. (# of atoms * mass of element X [amu]) + (# of atoms*mass of element Y [amu]) b. 4 sig figs 22. How to calculate formula mass of isotopes? a. (isotope 1 mass* % abundance) + (isotope 2 mass * % abundance) = avg atomic mass of element 23. What’s the diff between a heter/homogenous mixture? a. Hetero = has 1 or more visible boundaries between components b. Homo = no visible boundaries; mixed i. AKA solution OR if in wateraqueous solution c. Mixtures can be separated physically, whereas compounds can’t; they need to be chemically separated 24. What are some filtration techniques? a. Filtration = separate mixture based on particle size i. Crystallization = based on differences in solubility of the mixture components ii. Distillation = based on differences in volatility iii. Extraction = based on differences in solubility in diff solvents iv. Chromatography = based on differences in solubility in a solvent vs a stationary phase

Chapter 3 1. What is a mole? a. Mole = amount of substance that contains the same # of entities as there are atoms in 12g of carbon-12 i. Entities can refer to atoms, ions, molecules, formula units, electrons b. 1 mole = 6.02 x1023 entities AKA Avogadro’s number (N) c. n = amount of substance 2. What is molar mass a. Molar mass (M) = mass per mole of its entities i. Fore Monatomic elements, molar mass is the periodic atomic mass (g/mol) ii. Interconverting moles, mass and # of entities

¿ of grams =g 1 mol 1 mol mass (g) * ( ) = mols ¿ of grams 6.011 x 10 23 # of moles * ( ) = entities 1 mol 1 mol ) = mols (n) # of entities * ( 6.011 x 10 23 entities mass n (mols) = (n =m/M) molar mass

1. # of moles * ( 2. 3. 4. 5.

a. number of entities = n*N i. where N (entities/mol) is 6.011 x1023 3. How do you calculate Mass percent from the chemical formula? a.

b.

atoms of x (¿ formula)(atomic mass of x [ amu] ) *100 molecular mass of compound [ amu] moles of x (¿ formula)( molar mass of x [ g /mol ]) *100 molecular mass of compound [ g /mol ]

4. How to calculate mass fraction and the mass of an element? a.

( mass of compound )∗(mass of element ∈1 mol of compound ) (molecular mass of 1 mol of compund)

5. What are the molecular and empirical formulas? a. Empirical formula = simplest formula that shows lowest whole number of moles and gives relative # of atoms of each element i. AKA ratio of elements b. Molecular formula = shows the actual # of atoms of each element in a molecule of the compound

g ) mol = whole number g ) EF mass ( mol

MF mass( c.

6. What is combustion apparatus and how do you use it to calculate mass of H and C? a. Combustion = RX of substance w/oxygen b. CO2 + H2O are the products 7. What are the steps to reading/balancing an equation? a. Chemical equation show: i. Identities and quantities of substances ii. Uses formula instead of name

iii. Must be balanced – law of conservation b. Steps i. Translate statement into skeleton ii. Balance atoms iii. Adjust coefficients iv. Check 2x v. Add states of matter 8. What is a limiting and excess reactant? a. 1 reactant may limit the amount of product that can form b. Limiting reactant = completely used in RX c. Excess reactant = some left over after RX 9. How do you calculate reaction yield? a.

Actual yield x 100 theoretcal yield

b. Theoretical = is found after doing stoichiometry 10. How do you find the fraction mass of the main product from side product? a. AB = 73% b. BC = 68% c. AC = (.73)(.68)x100=49.6450% Chapter 4 1. Define Solute, solvent, active solvent, electron density a. Solute = being dissolved (smaller amount) b. Solvent = doing the dissolving (larger amount) c. Passive solvent = disperse substances into individual molecules d. Active solvent = strongly interacts and sometimes reacts w/substances e. Electron density = measure of the probability of an electron being present at a specific location i. (low density) +---- (high density) ii. S+ = low density; S- = high density iii. -> = interacts with cations (+) iv. + = interacts with anions (-) 2. Describe the conductivity of solutions a. Electrolyte solution i. Has ions present ii. Conductive b. Non-electrolyte solution i. Distilled water ii. No ions iii. No conduction c. Solid ionic state i. Vibration ii. No free movement of ions  no conduction d. Ions move to electrodes where ReDox RX happen i. Cathode = M+ + e-  M ii. Anode = X- = x + ee. Ionic compounds are electrolytes bc free mobility of ions; unless solid 3. Define concentration and molarity a. Concentration = quantity of solute present in a given quantity of solution b. Molarity =

mass solute L of solution

4. How do you prepare a molar solution?

M=

n V

(n = number of moles of solute)

a. Determine the required concentration b. Divide volume needed c. Calculate the moles needed from which mass is calculated d. Measure and transfer the sample to volumetric flask e. Dilute to the mark 5. How do you convert a concentrated diluted solution? a. M1V1 =M2V2 b. Concentrated solution has more solute per unit volume

6. Writing equations for Aq ionic RX a. AKA precipitation RX b. Steps: i. Write molecular formula ii. Separate into ions (only aq states) iii. Net ionic equation (remove spectator ions/cancel out ions) c. Total ionic equation = shows all soluble ionic substances dissociated into ions (fully written out) d. Spectator ions = no change to chemical change e. Net ionic equation = shows actual change; eliminates spectator ions 7. Predicting precipitate a. Precipitation Rx = 2 soluble ionic compounds  insoluble product AKA precipitate b. Consider all cation-anion possibilities and use solubility rules to determine if combo is INSOLUBLE...