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CHEM 121: Final Exam Study Guide
Chapter 1
• Know the Scientific method
• Know the definitions for hypothesis, scientific law, and scientific theory.
Chapter 2
• Length, mass, weight, volume
- Know 1 cm3≡1 mL and 1 dm3≡1 L
• Volume by calculation
- Vrectangular solid = length x width x thickness
• Significant figures and Scientific notation
- Rounding, in addition or subtraction, in
multiplication or division
- Keep as many sig figs until final answer
• Volume by displacement
• Scientific notation
• Metric system
- Know the metric prefixes: kilo (k), deci
(d), centi (c), milli (m), and micro (μ)
- Be able to perform metric-metric
conversions using these prefixes.
• Use the metric-English conversions provided
(1 in. ≡2.54 cm; 1 lb=453.6 g; 1 qt=946 mL)
m
V
- Calculate density, mass, or volume
- Identify what items sink or float given
densities of liquids and solids.
• Density: d =
• Temperature
- Know the formulas for converting ˚F-to-˚C or
˚C-to-˚F and K-to-˚C or ˚C-to-K
• Percentage: ratio of parts per 100 parts
- Given amount of part and whole, calculate %
- Use a given % to solve for part or whole
Chapter 3
• Know that matter is studied at the
macroscopic, microscopic, particulate
(molecular) levels
• Physical states of matter
- Determine physical state of substances
(solids, liquids, gases) given descriptions
of volume, shape, particles moving, etc.
• Classify properties and changes as
physical or chemical
- Know terms for changes of state:
- Melting, freezing, vaporizing,
condensation, sublimation, deposition
CHEM121 Final Exam Study Guide
• Classification of matter
– Given examples, determine which are
elements, compounds, or mixtures
– Given molecular-level images, determine
which are elements, compounds, or
mixtures and solids, liquids, or gases
– Distinguish between homogeneous and
heterogeneous mixtures
• Chemical reaction:
- reactants: starting materials
- products: substances produced in reaction
page 1 of 11
Chapter 3 (Continued)
kinetic energy (KE): energy associated with an
object’s motion
- Faster objects have higher KE.
- Know the relative kinetic energy of solids,
liquids, and gases
potential energy (PE): energy due to its
position or composition (chemical bonds)
heat: energy is transferred from a hotter
substance to a cooler substance
– Identify what lost heat and gained heat given
different scenarios.
Endothermic versus Exothermic changes:
– endothermic: a change requiring energy
– exothermic: a change that releases energy
Endothermic versus Exothermic changes:
– Determine if a physical change or chemical
change is exothermic or endothermic.
• Law of Conservation of Mass
- Solve problems conserving mass.
• Law of Conservation of Energy
- Know 6 forms of energy: heat, chemical,
light, electrical, mechanical, and nuclear
Chapter 4
• Know the main points of John Dalton’s
Model, what was later proven wrong and why
• Atomic notation:
• KNOW people and discoveries associated
with the discovery of protons, neutrons,
electrons
mass # (A): # of protons + # of neutrons
• Know Rutherford’s Alpha Scattering
Experiment and what was determined from
it (atomic nucleus, atom mostly empty space,
relative size of atom relative to its nucleus)
• Determine # of protons, neutrons, and
electrons for any given isotope or ion.
• Subatomic particles
–proton (p+): +1 charge, inside nucleus
–neutron (n): neutral, inside nucleus
–electron (e–): –1 charge, outside nucleus
• Electrostatic force: force resulting from a
charge on particles
- Objects w/ like charges repel one another.
- Objects w/ unlike charges attract each
other.
• Know definitions of isotope, atomic mass
CHEM121 Final Exam Study Guide
mass number = A
E = element symbol
atomic number = Z
atomic # (Z): # of protons=# of electrons
• Know the atomic mass reported on the
Periodic Table is the weighted average of all
naturally occurring isotopes for that element.
• Use the Periodic Table to identify the most
abundant isotope of any element given its
naturally occurring isotopes.
• Use the Periodic Table to identify those
elements whose naturally occurring isotopes
are all radioactive and unstable.
• Determine the number of atoms of an element
in a compound given the chemical formula.
page 2 of 11
Chapter 4 (Continued)
The Periodic Table
column = group, family
row = period, series
Representative Elements: A Group Elements
Group IA: alkali metals (except H)
Group IIA: alkaline earth metals
Group VIIA: halogens
Group VIIIA: noble gases
Transition Metals/Elements: B Group
Elements
Inner Transition Elements:
- Elements in lanthanide and actinide series
• Metals, nonmetals, and semimetals:
- Know properties of metals and nonmetals
- Location on Periodic Table and properties
• Know which elements exist as solids, liquids,
gases at room temperature (25˚C)
• Know the names and symbols for all elements
included in Ch. 4 notes, p. 11, Ti (titanium), Sr
(strontium) , and uranium (U). Spelling counts!
Chapter 9
Wavelength (λ) is inversely related to
frequency (ν) and energy (E):
– As λ↑ → ν↓, Ε↓ or As λ↓ → ν↑, Ε↑
Electromagnetic Spectrum:
– continuum of radiant energy
– gamma (γ) rays to radio waves
Know the people and ideas associated with
the Classical Model of Matter
– Dalton, Thomson, Rutherford, Maxwell
Planck and Quantum Theory
– proposed energy is absorbed and emitted as
bundles = quanta
– single bundle of energy = quantum
Einstein and the Photoelectric Effect
– Experimental evidence for light existing as
particles = photons
Bohr Model of the Atom
– Electrons move in quantized orbits called
“energy levels” around nucleus
– Know if energy is gained or lost for e-s
moving from one energy level to another.
– ground state: e-s in lowest E level(s)
CHEM121 Final Exam Study Guide
Atomic Orbital Shapes
– reflect the “probability density” for an
electron in a given orbital
– As principal energy level (n=1, 2, 3,…)
increases, the orbital size increases.
– Know energy levels and sublevels (s, p, d, f)
– Know the general shapes for s and p orbitals.
Be able to write electron configurations for any
neutral atom for elements #1-20.
• Write using full notation and core notation
(Noble Gas abbreviation)
• Know Representative Elements usually form
ions that are isoelectronic with a Noble Gas
Define atomic radius, ionic radius, and ionization
energy.
Know Periodic Trends for
• Atomic radius and Metallic Character
- Increase down a group
- Decrease left to right across a period
• Ionization Energy:
- Decreases down a group
- Increases left to right across a period
- Opposite trend as atomic radius since
IE ↓ as atomic radius ↑
page 3 of 11
Chapter 10
core electrons: electrons belonging to filled
electron shells
Be able to identify a bond as ionic, metallic,
polar covalent, or nonpolar covalent
valence electrons: outermost electrons
– group number for each element is equal to
the number of valence electrons it has
octet rule: atoms bond such that each has
8 electrons, except H only needs 2 electrons.
Determine the formulas for the ions formed
by Representative (Main Group) Elements using
a Periodic Table.
electronegativity (EN): ability of an atom in a
bond to draw electrons to itself
– Know F is most electronegative, further away
from F, less electronegative an atom.
ionic bond: electrostatic attraction between
cations and anions in an ionic compound
covalent bond: sharing of electrons between
two nonmetal atoms
• polar covalent: unequal sharing of e−s by 2
atoms with different EN values
• nonpolar covalent: equal sharing of electrons
by two atoms with equal EN
Draw dipole arrows to indicate which atom in a
bond is more electronegative
bond length: distance between nuclei of 2
bonded atoms
Know the how bond length is related to bond
strength
– shorter the bond, the stronger the bond
– single bonds are the longest and weakest
– double bonds are shorter and stronger
– triple bonds are the shortest and strongest
Metallic Bond:
- Metals exist as nuclei in a “sea of electrons”
→ special properties of metals resulting from
electrons’ freedom to move around
CHEM121 Final Exam Study Guide
Draw Lewis Structures for Molecules &
Polyatomic Ions
– The central atom will be indicated.
– Know H atoms are bonded to the O atoms in
ternary oxyacids.
Molecular Shapes and Polarity
• Use Lewis structure and to get 3D shape
and bond angles:
- AX2 → linear → 180° ∠
- AX3 → trigonal planar → 120° ∠
- AX4 → tetrahedral → 109.5° ∠
- AX2E → bent → <120° ∠
- AX3E → trigonal pyramidal → <109.5° ∠
- AX2E2 → bent → <109.5° ∠
• Use electronegativity to determine if a
bond is polar or nonpolar covalent.
• Given a molecule, determine if it’s polar or
nonpolar using the 3D shape and dipoles.
“Like dissolves like” rule
- Polar substances will mix and dissolve in one
another, and nonpolar substances will mix
and dissolve in one another.
- Polar substances will NOT mix and dissolve
in nonpolar substances and vice versa.
- Use the polarity of substances to
determine which liquids are miscible and
which solids are soluble in a given solvent.
Hydrocarbon (molecules containing only C and
H atoms) are always nonpolar.
Alcohols (molecules with a hydroxyl –OH
group) are always polar.
page 4 of 11
Chapter 12
Intermolecular Forces (IMF’s): attraction between 2 different molecules in a liquid or solid
• Identify the type of intermolecular force for a molecule as London/dispersion forces, dipoledipole forces, hydrogen bonding, or ion-diple forces
• Know that hydrogen bonds are the strongest type of intermolecular force, dipole-dipole forces
are the next strongest, and London forces are generally the weakest.
– Recognize that London forces increase with more electrons—use size to determine relative
number of electrons for different molecules.
– Know the terms: evaporation, boiling point, vapor pressure, volatile, nonvolatile
– Recognize how IMF’s influence vapor pressure and boiling point.
– Given different substances, determine which has the highest boiling point based on IMF’s.
• Given a bond or intermolecular force, identify it a polar covalent, nonpolar covalent, ionic,
metallic, ion-dipole forces, London/dispersion forces, dipole-dipole forces, or hydrogen bonding.
• Know ionic and covalent bonds are stronger than all intermolecular forces, even H bonds.
• Given its formula, classify a solid as ionic, molecular, metallic or network covalent.
• Know the unique properties of ice resulting from the hydrogen bonds between molecules in the
solid (e.g. density of solid versus liquid, why snowflakes have hexagonal symmetry, etc.).
• Know the regions and features of a Heating/Cooling Curve.
Chapter 5
ionic compound: a compound consisting of
metal cations and anions held together by
ionic bonds
molecule (or molecular compound): a compound
consisting of nonmetal atoms held together by
covalent bonds
Identify a compound as an ionic compound or
molecule given its name or chemical formula.
Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2,
Naming cations:
- Group IA, IIA, Al, Ag, Zn, Cd:
element name + ion
- All other metals - Stock system:
- element name (charge in Roman #s) + ion
CHEM121 Final Exam Study Guide
Naming anions: nonmetal stem + "-ide" + ion
Know the ions formed by the A Group or
Representative elements using the Periodic
Table.
Know the names and formulas of
POLYATOMIC IONS included the Chapter 5
lecture notes!
Naming ionic compounds:
- cation name + anion name
- IA, IIA, Al, Ag, Zn, Cd don't need Roman #s
- All other metals need Roman #s
Nomenclature for Ionic Compounds:
- Given the formula of a compound, determine
its name, or vice versa.
page 5 of 11
Chapter 5 (Continued)
Naming binary molecular compounds:
- Use Greek prefixes when more than one
atom of an element is present.
- Know names for NH3, CH4, and H2O2, or
given the name, write the formula.
Naming binary and ternary acids:
H('s) + "-ide" ion → "hydro____ic acid"
- e.g. Cl- = chloride ion → HCl (aq) = hydrochloric acid
H('s) + "-ate" ion
"_____ic acid"
- e.g. SO4 = sulfate ion → H2SO4 (aq) = sulfuric acid
2-
H('s) + "-ite" ion →"_____ous acid"
- e.g. NO2- = nitrite ion → HNO2 (aq) = nitrous acid
Chapter 6
Avogadro's number = 6.022x1023
Molar Mass - Be able to get molar masses (in
g/mol) for atoms and compounds
Standard temperature & pressure (STP):
T=0°C and P=1.00 atm
Molar volume: 1 mole of any gas at STP
occupies 22.4 L
Carry out Mole calculations using
- Avogadro's Number (N): 6.022 x 1023
- Molar masses of atoms and compounds
- Molar volume at STP: 22.4L/mole for a gas
Percentage composition:
- Find percent composition of all elements in a
compound given its formula or name.
Chapter 7
Writing Chemical Equations: reactants vs.
products
Be able to balance chemical equations.
Change coefficients, NEVER subscripts to get the
same # on both sides of equation.
Classify reaction types & balance equations:
- Combination
- Decomposition
- Single-replacement
- Double-replacement/precipitate
- Double-replacement/Neutralization
- Combustion Reactions
- If a reaction is a redox reaction or not
CHEM121 Final Exam Study Guide
Be able TO PREDICT PRODUCTS given a
set of reactants and Solubility Rules and the
Activity Series for the following TYPES of
REACTIONS:
– Single-replacement reactions
– solid metal + metal solution/acid
– solid metal + H2O(l)
– Combustion reactions
– CxHy + O2 → CO2 (g) + H2O (g)
– CxHyOz + O2 → CO2 (g) + H2O (g)
– Acid-Base Neutralization Reactions
– HX + MOH → water + salt
– HX + MHCO3 → water + CO2 + salt
– HX + MCO3 → water + CO2 + salt
page 6 of 11
Chapter 7 (Continued)
Strong, Weak and Non-Electrolytes
– strong electrolyte: breaks up completely
→ many ions present to conduct electricity
– e.g. strong acids & bases, aqueous salts
– weak electrolyte: breaks up to small degree
→ only few ions present to conduct
electricity
– e.g. weak acids & bases, insoluble salts
– nonelectrolyte: a molecular compound that
forms molecules in water
→ no ions → does not conduct electricity
Acids and Bases as electrolytes
– The list of strong acids will be given.
– All other acids are weak.
– Strong acids and bases break up completely
to produce many H+ or OH- ions.
– Weak acids and bases barely break up,
producing on a few H+ or OH- ions.
Representing Strong and Weak Electrolytes
– Recognize soluble ionic compounds and strong
acids break up into ions with a physical state
of aqueous, (aq).
– All solids, liquids, gases, and weak acids are
shown as compounds.
Chemical Equations & Net Ionic Equations
– Chemical Equation: compounds shown intact
– Complete/Total Ionic Equation:
– shows strong electrolytes as ions
– Spectator Ions: ions that remain
unchanged during a reaction
– Net Ionic Equation: Shows what
substances change in a chemical reaction
Be able to write Total and Net Ionic
Equations.
Chapter 8
Stoichiometry: Use mole-to-mole ratios to relate and calculate the amounts of reactants
and/or products in a chemical reaction
• Calculate the mass or volume of a reactant or product given the amount of another reactant
or product and the balanced chemical equation.
Yields of Reactions
• theoretical yield: amount of product
predicted using the balanced equation when
limiting reagent is used up (can be
calculated)
CHEM121 Final Exam Study Guide
• actual yield: amount of product one actually
obtains (generally given in the problem)
Percent yield =
actual yield
theoretical yield
page 7 of 11
× 100%
Chapter 11: Gases
• Know the properties of gases.
• Know definitions: vacuum, gas pressure,
atmospheric pressure, compressibility
• Recognize that atmospheric pressure
decreases with altitude.
Gas pressure and Atmospheric pressure
• Convert between units of pressure:
1 atm ≡ 760 torr ≡ 760 mmHg = 14.7 psi
• Know how changes in volume, temperature,
and number of particles affect gas
pressure.
• Given 2 sets of conditions, solve problems
P1V1 P2 V2
using
, including canceling
=
T1
T2
variables that stay the same to simplify.
– Recognize that temperatures (T’s) must
be in Kelvins.
Solve for a variety of problems involving gases
• Use ideal gas law (PV=nRT) to solve for P, V,
n, or T. R=
0.0821 L ⋅ atm
mol ⋅ K
will be given.
Standard Temperature & Pressure (STP):
T= 0˚C and P=1.00 atm
• Molar volume of a gas at STP = 22.4L/mol
• Solve for gas density or molar mass at STP.
• Identify an unknown elemental gas given its
gas density at STP by solving for molar mass.
Dalton’s Law of Partial Pressure:
• Use Dalton’s Law (Ptotal = P1 + P2 + P3 + …) to
solve for the total pressure or the partial
pressure of one gas in a mixture of gases.
• Recognize that when a gas is collected over
water, the total pressure is due to water
vapor and the gas.
Chapter 13:
• solution: uniform mixture of two or more
substances as atoms, ions, or molecules
– a solute dissolved in solvent
• Know how temperature affects the
solubility of gases and solids in solution.
• Know how gas pressure affects the
solubility of gases in solution.
• Recognize what occurs at the molecular
level when a solute dissolves in water.
• Recognize what can be done to increase
the rate of dissolving: heating solution,
stirring solution, grinding solute into
smaller particles
• Know the definitions for unsaturated,
saturated, and supersaturated
• Use “Like dissolves like” Rule and the
Solubility Rules to predict what
substances are soluble/insoluble in or
miscible/
immiscible with water or other solvent
Mass Percent Concentraton (M/M%):
M/M%=
mass of solute
mass of solution
Molarity (Molar
Concentration)=
moles of solute
liters of solution
Weight/Volume Percent (% w/v)
% w/v =
CHEM121 Final Exam Study Guide
× 100%
mass of solute
mL of solution
×100%
page 8 of 11
Chapter 13 (Continued)
Dilution Equation: M1 V1 = M2 V2
– Be able to solve for M2 in terms of
molarity, mass percent conc., or percent w/v.
– Account for dilution with DI water for V2
Molarity and Mass Percent Conc.
Calculations
• Solve for amount of solute, solvent, or
solution given mass percent
concentration, molarity, etc.
• Use molarity and volume to solve for moles
• Solve problems involving molarity and mass
percent concentration using unit analysis.
• Know definitions for: diffusion, semi-permeable
membrane, osmosis, isotonic, hypertonic,
hypotonic, and osmotic pressure.
Chapter 14
Know properties of acids and bases
Know the Arrhenius and Brønsted-Lowry
(B-L) definitions for acids and bases
Given an acid-base reaction,
- Classify each reactant as an Arrhenius
and/or a Bronsted-Lowry acid or base
- Indicate the conjugate acid-base pairs.
- Note that conjugate acid-base pairs
differ only by a H+ ion.
Strong and Weak Acids and Bases
– Know strong acids ionize completely to
produce many H+ (or H3O+) ions.
– Know strong bases dissociate completely to
produce many OH− ions.
– Classify a substance as a strong or weak
acid or strong or weak base given its
chemical formula and %ionization or
% dissociation
Recognize water rarely ionizes to form ions
→ It does not conduct electricity.
→ Ion-product or dissociation constant for
water at 25°C, Kw=[H+][OH−] =1.0x10−14
acidic solutions: [H3O+] > [OH–], pH < 7
basic solutions: [OH–] > [H3O+], pH > 7
neutral solutions: [OH–] = [H3O+], pH = 7
CHEM121 Final Exam Study Guide
Use pH to classify a substance as neutral,
strongly or weakly acidic, strongly or weakly
basic.
Calculate pH or pOH using
• [H+]=10–pH and [OH–]=10–pOH
• pH + pOH = 14.00
• Kw=[H+][OH−] =1.0x10−14
• Know definitions for indicator, standard
solution, titration, and endpoint
Buffer Systems
• Know a buffer system consists of a weak
acid and its conjugate base.
• Know how the buffer system (weak acid/
conj. base) can neutralize small amounts of
H+ or OH−, so a buffered solution can
maintain its pH.
• Know how blood is buffered to keep its pH
between 7.35-7.45given its buffer system:
H+(aq) + HCO3-(aq)
H2CO3(aq)
CO2(g) + H2O(l)
– Define acidosis, alkalosis, and
hyperventilation – Use Le Châtelier's Principle to explain how
changes in CO2 levels can cause acidosis or
alkalosis.
page 9 of 11
Chapter 15
• Distinguish between spontaneous and
nonspontaneous processes.
• Know the definitions for reaction rate,
activation energy (Eact), catalyst, enzyme.
– Know how a catalyst lowers the
activation energy for a reaction by
providing an alternative pathway for
bond breaking and bond formation.
• Know Collision Theory, and recognize the
activation energy and collision geometry
requirements for a chemical reaction to occur.
• Be able to identify and explain the three
factors that increase the rate of chemical
reactions.
• At equilibrium, the rates of forward and
reverse reactions are equal.
– Reactant and product concentrations are
constant (not changing) but do not have to
be equal.
Endothermic and Exothermic Reactions
– Know definitions for endothermic and
exothermic reactions.
– Know relative energy of reactants and
products based on the strength of their
bonds.
– Recognize if energy is released/absorbed,
based on temperature changes for the
surroundings following reaction.
Reaction Energy Diagrams
– Recognize the different parts of a reaction
energy diagram, including the energy of
reactants and products, the heat of the
reaction (ΔH), the activation energy (Ea)
for the forward and reverse reactions, and
the effects of a catalyst.
– Use a reaction energy diagram to classify a
reaction as endothermic or exothermic.
Know Le Chatelier's Principle and apply it to
systems at equilibrium to solve problems
involving changes in concentration of
reactants or products or temperature
changes.
Chapter 16
• Be able to determine oxidation numbers for all the elements/atoms/ions in a chemical equation.
• Use oxidation numbers to determine which reactant was oxidized (served as reducing agent)
and which reactant was reduced (served as oxidizing agent)
• Write half-reactions to determine the # of electrons transferred.
CHEM121 Final Exam Study Guide
page 10 of 11
Chapter 17:
• Know the atomic notation for the following: protons, neutrons, electrons, alpha particles, beta
particles, gamma rays, and positrons.
• Complete and balance nuclear equations, including writing complete nuclear equations for alpha,
beta, positron, and gamma emissions.
• Know the terms: parent and daughter nuclide
• Write and balance nuclear equations given the radioactive decay/emission and the parent or
daughter nuclide
• Solve problems involving half-life.
You will be given a CHEM 121 Periodic Table with the Polyatomic Ions, the
Solubility Rules, the Activity Series, and the List of Active Metals.
Part of the exam will be Scantron, so bring a #2 pencil and a basic 100-item
Scantron form (Form no. 882-E) with choices A through E.
CHEM121 Final Exam Study Guide
page 11 of 11