CHM 151: UNIT 1 STUDY GUIDE
OPENSTAX CHEMISTRY: ATOMS FIRST
ESSENTIAL IDEAS: MATTER & MEASUREMENT
CHEMISTRY IN CONTEXT (CH 1.1)



Chemistry
o Central Science
o Study of composition, properties, and interactions of matter
Scientific method
o Observation, hypothesis, experiments, theory or law
Domains of Chemistry
o Macroscopic, microscopic, and symbolic
PHASES AND CLASSIFICATION OF MATTER (CH 1.2)





Mater has both mass and volume
Differences in States: solid, liquid, gas, and plasma
Mass vs Weight
o Mass is the amount of matter
o Weight refers to the force that gravity exerts on an object
o Law of Conservation of Matter states that the amount of matter does not change
during a physical or chemical change
Atoms vs Molecules
o Atoms are the smallest particle of an element that retains the properties of the element
o Molecules consist of 2 or more atoms joined together (N2, H2O, NaCl)
Classification of matter
o Pure substances are composed of a single type of particle and its composition does not
vary
 Elements – simplest form that cannot be broken down
 Compound – composed of 2 or more different elements in a fixed proportion
o Mixtures are composed of 2 or more types of particles in proportions that can vary and
the components can be separated by physical processes
 Homogeneous mixtures are thoroughly mixed together such that it has the same
composition throughout; also called a solution
 Heterogeneous mixtures are not thoroughly mixed such that it has a different
composition in different regions; usually layers are formed
PHYSICAL AND CHEMICAL PROPERTIES (CH 1.3)


Physical Properties and Changes
o Properties measured or observed without changing the chemical composition of the
substance
o Changes in physical state
Chemical Properties and Changes
o Properties measured or observed when changing the chemical composition of the
substance
o Changes in chemical composition
 Intensive vs Extensive Properties
o Intensive properties are independent of amount of substance. such as density, melting
point, and boiling point.
o Extensive properties depend on the amount of substances such as mass, volume, and
moles.
MEASUREMENTS (CH 1.4)


Units
o Metric and SI units
o Know the standard base unit for length (m), mass (g, kg), time (s), temperature (C, K),
amount of substance (mol), and volume (L, m3)
o Prefix multipliers can increase or decrease the size of the base unit and can be used in
front of any metric base unit
 Know prefix multipliers from nano to giga including symbols and values
Density
o Density = mass / volume (g/mL)
o Convert the mass to grams (g) and the volume to milliliters (mL) before dividing
MEASUREMENT UNCERTAINTY, ACCURACY, AND PRECISION (CH 1.5)



Exact vs Measured Numbers
o Exact numbers are counted (ex: 10 hats) or defined relationships within the same unit
of measurement (ex: 1 ft = 12 in)
o Measured numbers are obtained by using a tool (ex: 2.80 cm)
Significant Figures in Measurements
o Estimating a digit while using a measuring tool increases the number of significant
figures (sf) in the measured number
o Be able to determine the number of significant figures in a given number
 Non-zero digits are significant: 359.21 cm (5 sf)
 Zeros in the middle of a number are significant: 400.8 s (4 sf)
 Zeros at the end of a decimal number are significant: 9.200 L (4 sf)
 Zeros at the beginning of a number are not significant: 0.0032 m (2 sf)
 Zeros at the end of a non-decimal number are not significant: 370 g (2 sf)
Significant Figures in Calculations
o Know how to determine the number of significant figures to keep in your answers when
doing calculations
 Multiplication and division answers have the same number of
significant figures as the number in the problem with the fewest
significant figures
2
 Addition and subtraction answers have the same number of
decimal places as the number in the problem with the fewest
decimal places


In problems that contain both types of functions, use the intermediate answers
for each step to determine the final digits to keep in the answer
Precision vs Accuracy
o Accuracy is the closeness to the actual value
o Precision is the closeness of repeated measurement to each other
o Precision of a measuring tool is based on the number of significant figures the tool
provides including the estimated digit
MATHEMATICAL TREATMENT OF MEASUREMENT RESULTS (CH 1.6)

Equalities and Conversion Factors
o Equalities are the relationships between equal quantities in different units of
measurement (ex: 1 L = 1000 mL)
o Conversion factors are the fractional representation of the equality; there are two
𝟏𝑳


possible conversion factors for each equality (ex: 𝟏𝟎𝟎𝟎 𝒎𝑳 𝒐𝒓
𝟏𝟎𝟎𝟎 𝒎𝑳
𝟏𝑳
)
Dimensional Analysis
o Use equalities between different units to convert a quantity from one unit to another
 Know metric to metric relationships (nano to giga) to create conversion factors
 English-Metric and English-English relationships will be provided on the test (see
chart below)
o To solve a problem, determine what you “know” and what you “need” then determine a
plan of attack; gather all the equalities necessary for the plan
o Set up the conversion factors such that initial unit is on the bottom of the first
conversion factor so it will cancel out; continue setting up the remaining conversion
factors in the same manner so units cancel until the final unit is obtained
o When converting units raised to a power, remember to raise both the number and the
unit to the power
o Use density as a conversion factor to convert between mass and volume
Temperature Conversions
o Be able to convert between units of temperature (F, C, K),
 Tc = (TF – 32) / 1.8
 TK = TC + 273.15
ATOMS, MOLECULES, AND IONS
EARLY IDEAS IN ATOMIC THEORY (CH 2.1)

Dalton’s Atomic Theory
o Element are composed of atoms
o An element consists of only one type of atom; all the atoms in the element are the same
o Atoms of one element differ from atoms of another element
3


o Atoms combine in whole number ratios to form compounds
o Atoms of one element cannot change into atoms of another element in a chemical
reaction; they can only be rearranged
Law of Definite Proportions
o All samples of a given compound have the same proportion of the constituent elements
Law of Multiple Proportions
o When different elements can form different compounds with different ratios of the
constituent atoms
EVOLUTION OF ATOMIC THEORY (CH 2.2)

Summarize the important experiments and results
o Thomson’s cathode ray tube experiment determined that electrons are negatively
charged subatomic particle approximately one thousand-times smaller than an atom.
o Millikan’s oil drop experiment determined the charge and mass of an electron.
o Rutherford’s gold foil experiment determined the nuclear structure of an atom and he
further determined the presence of protons.
o Soddy determined that atoms of the same element could have different masses and
these atoms are called isotopes.
o Chadwick determined the presence of neutrons that are neutral and contribute to
different mass of isotopes.
ATOMIC STRUCTURE AND SYMBOLISM (CH 2.3)





Atomic mass unit (amu)
o Mass of 1/12 of a carbon atom
o 1 amu = 1.6605 x 10-24 g
Structure of an Atom
o 3 subatomic particles: protons, neutrons, electrons
 Know charges, approximate size and position in the atom
o Elements are defined by the number of protons: atomic number (Z)
o Mass number (A) = protons + neutrons
o Atoms are neutral so number of protons equals the number of electrons
Ions
o Charged atoms due to gain or loss of electrons
o Representative elements (1A-8A) gain or lose electrons to achieve an octet of 8
valence electrons
o Metals lose e- to become positive ions (cations)
o Nonmetals gain e- to become negative ions (anions)
o Transition metals may form multiple ions
o Atomic charge = number of protons – number of electrons
o Identify the number of protons, neutrons and electrons in an ion
Chemical Symbols
o Know the symbols and names of the first 36 elements
o Be sure to use upper case for first letter and lower case for second letter
Isotopes
4

o Atoms of the same elements with different mass numbers (different # of neutrons)
o Atomic symbols provide the mass number and atomic number of a specific isotope
o Use atomic symbols to identify the number of protons, neutrons and electrons in the
isotope
o Write atomic symbols from information provided about the isotope
Atomic mass
o Weighted average of the mass number of all naturally occurring isotopes
o Average mass = ∑ (fractional abundance x isotope mass)
o Calculate the atomic mass based on information provided
o Mass spectrometry (MS) is used to determine natural abundance of isotopes
CHEMICAL FORMULAS (CH 2.4)


Formulas
o Molecular formulas indicate the type of atoms and how many are present (ex: C6H12O6)
o Empirical formulas indicate the lowest ratios of atoms in the formula (ex: CH2O)
o Structural formulas show how the atoms in the molecular formula are connected to
each other
 Isomers have the same molecular formula but atoms are connected in a different
order
The Mole
o Mole
 An amount of substance that contains a specific number of atoms of an element
of molecules of a compound
o Avogadro’s number
 1 mole of element = 6.022 x 1023 atoms of an element
 1 mole of compound = 6.022 x 1023 molecules of a compound
 Use Avogadro’s number to convert between atoms or molecules and moles
o Molar mass
 Molar mass of 1 mole of an element listed below element on periodic table
 Ex: 1 mole Na = 22.99 g Na *always use 4 sig figs for molar mass
 Formula mass of 1 mole of compound is the sum of the mass of the elements in
the formula of the compound
 Ex: 1 mole MgCl2 = 24.31 g Mg + 2(35.45 g Cl) = 95.21 g MgCl2
 Use molar mass to convert between mass and moles
COMMON EQUALITIES PROVIDED ON THE TEST
ENGLISH-ENGLISH EQUALITIES
ENGLISH-METRIC EQUALITIES
1 ft = 12 in
1 in = 2.54 cm (exact)
1 yd = 3 ft
1 m = 1.0936 yd
1 mi = 5280 ft
1 km = 0.62137 mi
5
1 gal = 4 qt
1 L = 1.0567 qt
1 qt = 32 fl oz
1 qt = 946.34 mL
1 lb = 16 oz
1 kg = 2.2046 lb
1 ton = 2000 lb
1 lb = 453.59 g
6