CHEM 139: 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
• Significant figures and Scientific notation
- Rounding, in addition or subtraction, in
multiplication or division
- Keep as many sig figs until final answer
• Scientific notation
• Metric system
- Know all the metric prefixes given on p. 9
in the Chapter 2 lecture notes
- 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)
• Volume by displacement
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
- Calculate weighted averages keeping track of
sig figs and decimal places when appropriate.
• Volume by calculation
- Vrectangular solid = length x width x thickness
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.
• Be able to identify properties and changes
as physical or chemical
- Know terms for changes of state:
- Melting, freezing, vaporizing,
condensation, sublimation, deposition
CHEM139 Final Exam Study Guide
• Classification of matter
– Classify substances as elements,
compounds, or mixtures
– Classify molecular-level images as 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
• Law of Conservation of Mass
- Solve problems conserving mass.
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Chapter 10
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
joule (J): SI unit of energy; 1 kJ=1000 J
calorie (cal): energy needed to raise the
temperature of 1 g of water by 1˚C
Be able to carry out calculations involving
energy in J, cal, Cal, and kilowatt-hours (kW·h).
specific heat (in J/g⋅˚C): amount of heat to
raise temp. of 1 gram of a substance by 1˚C.
- Recognize that the higher the heat capacity
of a substance, the more heat it can absorb
before its temperature rises.
Chapter 4
• 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)
• 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
• Atomic notation:
mass number = A
E = element symbol
atomic number = Z
CHEM139 Final Exam Study Guide
mass # (A): # of protons + # of neutrons
atomic # (Z): # of protons=# of electrons
• Determine # of protons, neutrons, and
electrons for any given isotope or ion.
• 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.
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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), and
Sr (strontium). Spelling counts!
Chapter 11
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)
– excited state: e-s in higher E level(s)
before lower levels are full
CHEM 139 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, p, and d
orbitals and the number of each (e.g. one s
orbital, three p orbitals, five d orbitals)
Be able to write ground state electron
configurations for any neutral atom or ion
• Write using full notation and core notation
(Noble Gas abbreviation)
• Recognize extra stability gained with filled
and half-filled d orbitals (Cr, Mo, Cu, Ag)
• Account for electrons gained or lost for ions
• Representative Elements usually form ions
that are isoelectronic with a Noble Gas
Know definitions for atomic 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: Decrease down a group
- Increase left to right across a period
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Chapter 12
valence electrons: outermost electrons
– group number for each element is equal to
the number of valence electrons it has
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
– 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
octet rule: atoms bond such that each has
8 electrons, except H only needs 2 electrons.
Draw Lewis Structures for Molecules &
Polyatomic Ions
Lewis structures for ternary oxyacids
– Oxygens around the central atom, and one
hydrogen each bonded to an oxygen
resonance structures: two or more
structures representing a single molecule with
delocalized electrons that cannot be
described fully with only one Lewis structure
– Recognize which molecules require
resonance structures.
– Know the relative length and/or strength
of bonds with delocalized electrons.
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.
Be able to identify a bond as ionic, metallic,
polar covalent, or nonpolar covalent
Chapter 14
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
CHEM139 Final Exam Study Guide
page 4 of 9
Chapter 14 (Continued)
• 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 molar mass to determine
relative number of electrons for different molecules.
– Know definitions for gas pressure, atmospheric pressure, and vacuum
– 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.
• Rank different compounds (ionic, molecular, atomic, etc.) in terms of increasing melting point.
• Compare the relative strength of chemical bonds or IMF’s for different substances based on
melting point.
• 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.).
Chapter 5
ionic compound: a compound consisting of
metal cations and anions held together by
ionic bonds
Know the ions formed by the A Group or
Representative elements using the Periodic
Table.
molecule (or molecular compound): a
compound consisting of nonmetal atoms held
together by covalent bonds
Know the names and formulas of
POLYATOMIC IONS included the Chapter 5
lecture notes!
Identify a compound as an ionic compound or
molecule given its name or chemical formula.
Naming ionic compounds:
- cation name + anion name
- IA, IIA, Al, Ag, Zn, Cd don't need Roman #s
- All other metals need Roman #s
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
Nomenclature for Ionic Compounds:
- Given formula of a compound, determine
name.
- Given the name of a compound, give formula.
Naming anions: nonmetal stem + "-ide" + ion
CHEM139 Final Exam Study Guide
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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 8
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.
- Use the formula or name of compound and
its percent composition to determine the
mass of one or more elements in a sample
of the compound.
Determine Empirical and Molecular Formulas
– Know empirical formula is simplest ratio of
atoms/ions present in compound.
– Use the masses of elements present, the
law of conservation of mass, and/or mass
percentage to calculate empirical formula.
– Determine the molecular formula using the
empirical formula and given molar mass.
Chapters 6 and 7
Classify reaction types & balance
equation:
- Combination
- Decomposition
- Single-replacement
- Double-replacement/precipitate
- Double-replacement/Neutralization
- Combustion Reactions
- If a reaction is a redox reaction or not
Be able TO PREDICT PRODUCTS given a
set of reactants and Solubility Rules and
the Activity Series.
CHEM139 Final Exam Study Guide
REACTION TYPES
– Combination/Synthesis reaction:
– metal + nonmetal → ionic compound (s)
– Decomposition reaction: AZ → A + Z
– Single-replacement reactions
– solid metal + metal solution/acid
– solid metal + H2O(l)
– Combustion reaction
– CxHy + O2 → CO2 (g) + H2O (g)
– CxHyOz + O2 → CO2 (g) + H2O (g)
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Chapters 6 and 7
REACTION TYPES (Continued)
– Double-Displacement/Precipitation Reactions
– precipitate (ppt): insoluble ionic compound
– Use Solubility Rules to determine if a
precipitate forms
– Predict products of a reaction given
reactants and Solubility Rules
– Acid-Base Neutralization Reactions
– HX + MOH → water + salt
– HX + MHCO3 → water + CO2 + salt
– HX + MCO3 → water + CO2 + salt
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
– Know strong acids & strong bases in notes!
– All other acids and bases are weak
– Recognize H2SO4(aq) → H+(aq) + HSO4-(aq)
Representing Strong and Weak Electrolytes
– Show soluble ionic compounds and strong
acids broken 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 9
Stoichiometry: Use mole-to-mole ratios to relate and calculate amounts of reactants and/or
products in a chemical reaction
Limiting Reagent Problems
• Calculate the mass or volume of product that can be made using the given amounts of each
reactant and the balanced chemical equation.
– Solve using the comparison-of-moles or comparison-of-mass method (i.e., determine if there is
enough of one reactant to react completely with the other)
• Indicate the limiting reactant (or reagent) and the reactant(s) in excess.
• Calculate the amount of reactant in excess that remains after the reaction.
Yields of Reactions
• theoretical yield: amount of product
predicted using the balanced equation when
limiting reagent is used up (can be
calculated)
CHEM 139 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 9
× 100%
Chapter 18
• Be able to determine oxidation numbers for all the elements/atoms/ions in a chemical equation.
• Use oxidation numbers to determine if a chemical equation represents a redox reaction.
• 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.
Chapter 15:
• solution: uniform mixture of two or more
substances as atoms, ions, or molecules
– a solute dissolved in solvent
Mass Percent Concentraton (M/M%):
M/M%=
mass of solute
mass of solution
× 100%
• Know how temperature affects the solubility
of gases and solids in solution.
Molarity=
• Recognize what occurs at the molecular level
when a solute dissolves in water.
Dilution Equation: M1 V1 = M2 V2
• 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
moles of solute
liters of solution
Molarity, Mass Percent Concentration, and
Solution Stoichiometry Calculations
• Solve for amount of solute, solvent, or
solution given mass percent concentration,
molarity, etc.
• Use molarity and volume to solve for moles
Calculate molarity of ions in solution (based
on number of ions present in a compound)
Chapter 16
Know properties of acids and bases
Know 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.
CHEM 139 Final Exam Study Guide
The strength of an acid is inversely
related to strength of its conjugate base.
– Weak acids have conjugate bases that are
stronger than H2O.
→ They react with water to produce OH−.
F−(aq) + H2O(l)
HF(aq) + OH−(aq)
→ The conjugate base of a weak acid reacts
with H2O to form conjugate acid and OH−.
– The conjugate base of strong acids are
weaker than H2O
→They cannot remove H+ from water, so
no reaction (NR).
page 8 of 9
Chapter 16 (Continued)
Acid-Base Stoichiometry Problems
• Calculate the molarity or molar mass of an
acid/base given the amount of base/acid
required to completely neutralize it.
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
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 (exact)
• Kw=[H+][OH−] =1.0x10−14
• Knowing that because pH is a logarithm,
the number of sig figs for the H+
concentration determines the number of
decimal places for the pH and vice versa.
• Calculate the pH of a solution given the
concentration of a strong acid or a strong
base.
• Calculate the pH of a solution after a
strong acid or a strong base has been
diluted.
You will be given a copy of the CHEM 139 Periodic Table with, the Solubility Rules,
the Activity Series, the List of Active Metals, along with relevant constants and
metric-English conversions.
Part of the exam will be Scantron, so bring a #2 pencil and a basic 100-item
Scantron form with choices A through E.
CHEM 139 Final Exam Study Guide
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