Chem 101 Final Study Guide PDF

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Chapter 1 - Introduction Classifications of Matter – 1.3 • • • • •

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Matter: anything that occupies space and has mass Chemistry: the study of matter and its change Atom: smallest distinctive unit •••• Molecule: two or more atoms together •••• Substance: group of identical atoms or molecules •••• o Element: atoms are identical o Compound: contains different atoms Homogenous Mixture: uniform, seems as one thing Heterogeneous Mixture: non-uniform, able to see the elements

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Pure Substance

Matter Mixtures

Physical and Chemical Properties of Matter – 1.4 •

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Physical Property: measurable, describes system stateCompounds Elements Heterogeneous Homogeno o Melting point, solubility, density, color Mixture Mixture o “Air is a gas,” “boiling point of water is 100 degrees,” “Sulfur crystals are yellow” Chemical Property: chemical reaction, compound cannot be restored without a change o Burning, explosions, decomposition, rusting o i.e. “Burning Hydrogen in Oxygen to form water,” “milk turning sour”

Measurement – 1.5 •

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SI base units of Measurement: o Length – meter (m) o Mass – kilogram (kg) o Time – second (s) o Temperature – Kelvin (K) o Amount of a substance – mole (mol) Temperature Scales: o Kelvin: absolute temperature scale, 273° C = 1 K

Handling Number – 1.6 • • •

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Accuracy: Correctness – How close the results were to where you intended them to be Precision: Exactness – How close the results were in relation to other results / trials Scientific Notation: Expresses a number as a product of a number between 1 and 10 and the appropriate power of 10 o All non-zeros are significant (1-9) o A zero between non-zeros is significant (209) o Final zeros to the right of the decimal are significant (10.0) o Initial zeros are not significant (0.47) o Final zeros in a number with no decimal may or may not be significant (9,500) Adding and subtracting: same number of digits past the decimal point as the number used with the least precision – the least amount of significant figures o If greater than 5, round up Multiplication and Division: least number or significant figures Exact Numbers: have no effect on significant figures – i.e. 60 seconds in a minute Volume (V): Length cubed, mL = 1 cm^3

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Mass (m): Measure of the quantity of matter in an object Density (d): Mass divided by volume, common unit = g/cm^3 o 1 g/cm^3 = 1 g/mL o Density (d) = mass (m) / volume (V) o Density (d) = grams / mL

Dimensional Analysis and Measurement – 1.5 & 1.7 •

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Conversion Factors: used as a multiplier to cancel unwanted units and replace them with new units o 1 ft. = 12 in written as Prefixes Used with SI Units Big Numbers: o Giga (g) = 1 gigameter (gm) = 1 * 10^9 m o Mega (M) = 1 megameter (Mm) = 1 * 10^6 m o Kilo (k) =1 kilometer (km) = 1 * 10^3 m = 1000 m Base Unit: Meter, gram, second Smaller Numbers: o Deci (d) = 1 decimeter (dm) = 0.1 m o Centi (c) = 1 centimeter (cm) = 0.01 m o Milli (m) = 1 millimeter (mm) = 0.001 m, 1000 mm = 1 m o Micro (μ) = 1 micrometer (μm) = 1 * 10^ -6 m o Nano (n) = 1 nanometer (nm) = 1* 10^ -9 m

Chapter 2 – Atoms, molecules and ions Atomic Theory and Understanding the Atom – 2.1 & 2.2 •

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Structure of the atom o Nucleus: dense central core o Protons: positively charged particle (+) o Neutrons: neutral particles o Electrons: negatively charged particles that exist in a large “cloud” around the nucleus (-) Neutral elements are not electrically charged: protons = electrons

Atomic Number, Mass Number, and Isotopes – 2.3 •

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Atomic Number (Z) o Number of protons in a nucleus o Determines element identity o # of electrons = # of protons if it’s neutral atom Mass Number (A): o Protons + neutrons in a nucleus o Determines isotope identity Isotopes: same number of protons and electrons, different neutrons isotopes have similar reactivity and bonding Ions: group of atoms with a net + or – charge o Cation: positive charge caused by the LOSS of an electron o Anion: negatively charge caused by the GAIN of an electron

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The Periodic Table – 2.4 & 2.5 • • • • • • • • •

Group 1A: Alkali Metals, +1 charge Group 2A: Alkaline Earth Metals, +2 charge Group 3A: +3 charge Group 5A: -3 charge Group 6A: -2 charge Group 7A: Halogens, -1 charge Group 8A: Noble Gases Nonmetals: Group 8A, C, N, O, F, P, S, Cl, Se, Br, I Metalloids: B, Si, Ge, As, Sb, Te, Po, At

Chemical Formulas and Naming – 2.6 & 2.7 • • • •

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Empirical: simplest whole number ratio of atoms (HO) Molecular: Exact types and number of atoms (H2O2) Naturally Occurring Diatomic Molecules: 2 atoms bonded together o H2, O2, F2, Br2, I2, N2, Cl2 Ionic Compounds: neutral compound containing cations and anions o Cl- + Na+ ! NaCl o Sodium loses one electron and becomes a sodium cation o Chlorine gains one electron and becomes a chlorine anion Type 1 Compound: Ionic Compound (metal and nonmetal), net charge on ionic compound is zero – so all charges must add up to zero o Subscript of the cation is equal to the charge on the ion o NaCl ! sodium chloride o MgO ! magnesium oxide Polyatomic Ions: lose or gain electrons as a group, charge is spread over 2+ atoms o Acetate – CH3COOHydroxide – OHo Ammonium – NH4+ Mercury (I) – Hg22+ o Bicarbonate – HCO3Mercury (II) – Hg2+ o Carbonate – CO32Nitrate – NO3o Chlorate – CIO3Nitrite – NO2o Cyanide – CNPechlorate – CIO4o Hydronium – H3O+ Permanganate – MnO4 o Phosphate – PO43Sulfate – SO42o Sulfite – SO32Type 2 Ionic Compounds (transition + nonmetal): transition metal becomes positively charged; value of charge is unpredictable o Charge is designated with a Roman Numeral (Carbon has a 4- charge) Type 3 Molecular Compounds: molecular compounds consist of only nonmetallic elements o 1st word: 1st element o 2nd word: stem of element name, change ending to “ide” o Prefixes designated the number of atoms (nonmetals only) ▪ 1 – mono ▪ 2 – di

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▪ 3 – tri ▪ 4 – tetra ▪ 5 – penta ▪ 6 – hexa ▪ 7 – hepta ▪ 8 – octa ▪ 9 – nona ▪ 10 – deca Common Acids Common Bases: o Hydrochloric Acid – HCl Sodium Hydroxide: NaOH o Hydrobromic Acid – HBr Ammonium Hydroxide: NH4OH o Sulfuric Acid: H2SO4 Lithium Hydroxide: LiOH o Carbonic Acid: H2CO3 Potassium Hydroxide: KOH o Nitric Acid: HNO3 o Phosphoric Acid: H3PO4 o Acetic Acid: CH3COOH Type 1 Compound: Ionic Compounds (nonmetal + metal) Type 2 Compound: Ionic Compound (transition metal + metal) Type 3 Compound: Molecular compound (only nonmetals Type 4 Compound: Acid

Chapter 3 – Stoichiometry Atomic Mass, Avogadro’s number, Molar Mass and Molecular Mass – 3.1 & 3.3 • • • • •

Atomic Mass: the mass of one atom of a substance in amu (top number on periodic table) Molecular Mass: the mass of one molecule in amu o Subscript: number of that particular atom in a compound Avogadro’s Number: 6.022 * 10^23 particles = 1 mol particles Moles of Atoms in Compounds: subscripts give mole ratios of atoms in a compound o 1 mole of a compound = molar mass of compound Molar Mass (MM): the mass of one mole of a substance (atoms or molecules) (bottom number on periodic table)

Chemical Reactions and Chemical Equations – 3.7 • • • • • • •

Reactants: compounds / elements being added in equation Products: compound made from reactants + Sign: “reacts with” ! “Yields” Chemical phase: g = gas, s = solid, l = liquid, aq= aqueous Moles reacting: number in front of reactant Rules For Balancing Chemical Equations: o Rule 1: if an element is present in just 1 compound on each side of the equation, balancing that element first o Rule 2: Balance free elements last o Rule 3: Fractions can be cleared at any time by multiplying all coefficients by a common multiplier

o Rule 4: Groupings of atoms (such as in polyatomic ions) may remain unchanged. Balance these groupings as a unit

Stoichiometry: Amounts of Reactants and Products – 3.8 •

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Percent Yield = actual yield / theoretical yield * 100 o Actual Yield: yield recovered during experiment o Theoretical yield: maximum obtainable yield – calculated from limiting reagent Use Mole Ratios to Predict How Much Product(s) o Mass (g) of A ! Moles of A ! Moles of B ! Mass (g) of B o A!B A+ B ! C

Limiting Reagents and Reaction Yield – 3.9 & 3.10 • •

Limiting Reagent: reactant that runs out first Excess Reagent: Reactant that is not consumed during reaction

Percent Composition – 3.5 •

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Percent Composition from Formula: part divided by whole o Step by step: ▪ 1: assume 1 mole of compound, find molar mass of the compound ▪ 2: determine the mass due to the component in the compound solving for ▪ 3. Divide the mass due to the component by the total molar mass compound and multiple by 100 o Tips: ▪ Chemical formula gives molar amounts ▪ Must convert moles to mass in grams ▪ Don’t forget to multiply mass of each atom by the number of occurrences in molecules (part) Empirical and Molecular Formula from Percent Composition: o Mass percent (assume 100 g) divide by molar mass ! moles of each element (divide by the smallest number of moles) ! mole ratio of elements (use mole ratio as integer subscript) ! empirical formula o Step by step: ▪ 1: assume 100 grams compound – convert mass to grams ▪ 2: convert grams to moles ▪ 3: divide by the smallest number of moles ▪ 4: use the resulting numbers as subscripts for an empirical formula ▪ 5: for molecular formula divide the molar mass of the compound by the molar mass of the empirical formula ▪ 6: multiply the empirical formula by this integer to determine molecular formula

Chapter 4 General Properties of Aqueous Solutions 4.1 • • •

Solution: homogenous mixture Solvent: component with largest amount (usually water) Solute: remaining components – smaller amounts

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Dissolve: solvent molecules surround and support solute molecules or ions o Example: bleach in water – bleach: solute water: solvent o Example: Nail polish in acetone – nail polish: solute acetone: solvent Aqueous solution: solution where the solvent is water Solvent is NOT consumed in reaction Strong Electrolyte: conducts electricity when dissolved in water o Fully Dissociates/Ionizes: breaks up into cations and anions ▪ NaCl (aq) -> Na+ (aq) + Cl- (aq) Weak Electrolyte: weak conductor when dissolved in water o Partially dissociates: CH3COOH " ! CH3COO- (aq) + H+ (aq) Nonelectrolyte: do not conduct electricity when dissolved in water o Does not dissociate Conductivity of Electrolytes in Aqueous Solutions o Non-electrolyte: no ionization, i.e. sugar o Weak electrolyte: some ionization, i.e. acetic acid, o Strong electrolyte: 100 % ionization, i.e. NaCl Strong electrolyte: * * * * * * In-between strength electrolyte: ** * * ** Weakest electrolyte: ** ** ** **

Precipitation Reactions (Metathesis Reactions) 4.2 • • • •

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Precipitation: Ions in solution combine to form an insoluble solid salt Precipitate: Solid salt that is formed Soluble: dissolves Molecular Equation: written as though all species exist as molecules o Pb (NO3) 2 (aq) + 2 NaI (aq) ! PbI2 (s) + 2 NaNO3 (aq) Precipitate soluble in H2O Ionic Equation: shows dissolved species as free ions o Pb2+ (aq) + 2NO3- (aq) + 2Na+ (aq) + 2I- (aq) ! PbI2 (s) + 2Na+ (aq) + 2 NO3– (aq) Net Ionic Equation: shows only the species that actually take part in the reaction o Pb2+ (aq) + 2I- (aq) ! PbI2 (s) Spectator Ions: ions that do not react in a solution and remain as ions o NO3- and Na+ Solubility Rules o Nitrates (NO3-) are soluble o Compounds containing Alkali Metals or NH4+ are soluble o Most chlorides, bromides, and iodides are soluble ▪ Exception: halides containing: Ag+, Pb+, and Hg22+ ions o Most sulfates (SO42-) are soluble ▪ Exception: Ca2+, Sr2+, Ba 2+ Ag+, Hg22+ and Pb 2+ o Most hydroxides are not soluble ▪ Exception: Ba(OH) 2 and hydroxides in rule 2 o Most carbonates (CO32-), phosphates (PO43-), sulfites (SO32-), and sulfides (S2-) are sparingly soluble ▪ Exceptions: those whose positive ions are listed in rule 2

Concentration of Solutions and Solution Stoichiometry – 4.5 and 4.6

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Molarity: Molar concentration of a solution o Molarity = Moles solute (n)/ Liters solution (V) Concentration of a solution: amount of solute in a given amount of solution Dilution: Water is added to a small amount of stock solution to make a less concentrated solution o M1V1 = M2V2 Amount of solvent changes – amount of solute DOES NOT Stock solution = Diluted solution Gravimetric Analysis: sample of unknown composition is dissolved in H2O and allowed to react with another substance to form a precipitate. Precipitate is filtered, dried and weighed

Acid – Base Reactions – 4.3 and 4.6 •

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Acids: o Arrhenius: dissociates in water and forms H+ ions (H3O+) and anions o Bronsted: proton donor Base o Arrhenius: dissociates in water to for OH- ions and cations o Bronsted: proton acceptor Strong Acids: o Complete dissociation in water to give H3O+ o HCl, HBr, HI, H2SO4 ,HNO3, HCIO3, HCIO4 o General formula M (OH) n, where M is a Group 1 metal or Barium ▪ Group 1 metal hydroxides: LiOH, NaOH, KOH, RbOH, Cs (OH) ▪ Group 2 metal hydroxides: Ba(OH) 2, Weak acids: rest of acids (HF, CH3COOH, H3PO4) o Partial dissociation in water, most of original compound remains o Other hydroxides (insoluble), Ammonia (NH3) Acid-Base Neutralizations o Reaction between acid & base ▪ Acid + Base ! Salt + Water Acid-Base Titration: solution of known concentration is added to a solution of unknown concentration until moles the base is fully neutralized

Oxidation and Reduction Reactions: Redox – 4.4 • •

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Oxidation-reduction reactions (Redox) are electron transfer reactions o There is a change in oxidation number for both substances Oxidation: o Charge becomes more positive o Electron is lost o Oxidation number increases o Na ! Na+ + 1e0^ +1^ Reduction: o Charge becomes more negative o Electron is gained o Oxidation number is decreased o Cl2 + 2e- ! 2Cl –

0^ -1 ^ Oxidation and reduction always accompany each other – neither can occur alone LEO the lion says GER o Lose Electron Oxidation o Gain Electron Reduction • Oxidation Number: Theoretical Charge on atom o Oxidation numbers have rules – rule earlier in the list always takes priority • Oxidation Number Rules 1. ON = 0 when neutral molecule or free element ▪ All elements at first have a neutral charge because they have the same number of protons and electrons • I.e. Na has an atomic mass of 11; this means 11 protons and 11 electrons. +11-11= 0 so a neutral charge 2. ON = ionic charge for an ion • Ion: molecule with a charge due to loss or gain of electron(s) 3. ON = -2 for oxygen in most compounds • EXCEPT: H2O2 and O224. ON= +1 for hydrogen • ON= -1 for H when bonded to metals in binary compounds (NaH, LiH, CaH2) 5. ON = -1 for halogens (F, Cl, Br, I, At) • If both elements are halogens, then the one higher on the list is -1 • ON is positive when halides are combined with O or ON 6. ON = -2 for 6A elements other than oxygen (O, S, Se, Te, Po) 7. ON does not have to be an integer • Highest possible ON of an element in 1A-7A is its group number • If composed of polyatomic ions: break down the compound into ions before determining the oxidation state of the element o If an ion has two capital letters in it, it’s polyatomic • Redox Half-Reactions: explicitly shows the electrons involved in a redox reaction o Oxidation half-reaction: 2Mg (s) ! 2Mg2+ + 4eo Reduction half-reaction: O2 (g) +4e- ! 2O2o Sum of half reactions: 2Mg (s) + O2 (g) ! 2MgO (s) • Oxidizing agent: o Reactant that helps oxidation o Oxidizing agent is reduced (gains e-) o Characteristic of nonmetals: i.e. fluorine, oxygen • Reducing agent: o Reactant that promotes reduction o Reducing agent is oxidized (loses e-) o Characteristic of an active metal o 2Mg (s) + O2 (g) ! 2MgO(s) 0 0 +2 -2 Mg loses e- Mg is oxidized O2 is the oxidizing agent O2 gains eO2 is reduced Mg is the reducing agent • Combination: 2Al (s) + 3 Br2 (g) ! 2AlBr3 (s) • •

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0 0 +3 -1 Decomposition: 2KClO3 (s) ! 2KCl(s) + 3O2 (g) +1 +5 -2 +1 -1 0 Combustion: CH4 (g) + 2O2 (g) ! CO2 (g) + 2H2O (g) -4 +1 0 +4 +2 +1 -2 Displacement: (halogen, hydrogen, and metal) o Halogen displacement: F2 > Cl2 > Br2 > I Cl2 (g) + 2KBr (s) ! Br2 (g) + 2KCl (s) 0 +1-1 0 +1 -1 o Hydrogen Displacement Reaction: o Activity series: metals are arranged according to their ability to displace H from an acid or water ▪ Reactivity with water: Ba (s) + 2H2O (l) ! H2 (g) +Ba (OH) 2 (aq) ▪ Reactivity with acid: Pb (s) + 2HCl (aq) ! PbCl2 (s) + H2 (g) o Pt (s) +2HCl (aq) ! No reaction o Metal Displacement Reaction: ▪ Reactivity of 2 metals: Higher metal becomes cation. Lower metal will be free metal ▪ Zn (s) CuSO4 (aq) ! ZnSO4 (aq) + Cu (s)

Chapter 5: Gases Properties of Gases – 5.1 and 5.2 •

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Gas: o Farthest apart o Lowest density o Highest kinetic energy o Most motion o Most diffusion o Most compressible o Least intermolecular forces o Not fixed shape or volume Solid: o Closest together o Incompressible o Highest density o Lowest kinetic energy o Slowest diffusion o Least motion o Fixed shape and volume o Most intermolecular forces Properties of gases o Eleven elements that are gaseous under normal conditions (N2, O2, H2, F2, Cl2, He, Ne Ar, Kr, Xe Rn) ▪ Atmospheric gases: N2, O2 ▪ Noble Gases: He, Ne, Ar, Kr, Xe, Rn Pressure: force applied by molecules per unit area

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o Pressure = force / area Atmospheric Pressure: Pressure exerted by the earth’s atmosphere o SI unit: atmosphere (atm) o Standard atmospheric pressure: 1 atm at sea level o 1 atm = 760 mm Hg o 1 atm = 760 torr

Gas Laws – 5.3 & 5.4 •

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Ideal gas: behaves as predicted by the ideal gas law o Molecules do not attract or repel each other o Volumes are negligible o PV=nRT o R= gas constant (0.08206) Standard Temperature and Pressure o 0 degrees Celsius and 1 atm o 1 mole ideal gas = 22.4 L Boyle’s Law: Pressure – Volume Relationship o Volume decreases as pressure increases o PV=k1 o P1V1=P2V2 Charles Law: Temperature – Volume Relationship o Volume increases as temperature increases o (V/T) = k2 o (P1/T1=P2/T2) o (V1/T1)=(nR/P) o (V1/T1)=(V2/T2) Gay Lussac’s Absolute Zero: K = C + 273 Avogadro’s Law: Volume – Mole Relationship o Volume of a gas depends on the number of moles o Volume increases as molecules increases o (V/n) = k4 Law of combining volumes o Equal volumes of different gases contain the same number of moles and molecules

Dalton’s Law of Partial Pressures – 5.5 • •

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Total pressure: sum of pressures each gas would exert alone (same V and T) Partial pressure: pressure exerted by each individual gas in the container o Ptotal=Pa + Pb + Pc o Pa= (na RT/V) o Pt= (ntRT/V) Mole Fraction: of a component in a mixture is the number of moles of the component divided by the total number of moles o Xi = (ni/nt) o Xi = (Pi/Pt) Collection of gases over water o Gas is collected into a container of water, water saturated gas rises ad displaces liquid water

Pt = Pgas + Pwater ▪ o Vapor pressure: partial pressure of the water vapor, must be deducted from the measured pressure of the gas

Kinetic Molecular Theory of Gases – 5.6 •

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Gas pressure: result of collisions between molecules and the walls of their container 1. Ideal gas molec...