Chapter 1 Notes
AP CHEMISTRY
UNIT 1
Chapter 1: Matter and Measurements
 I. Introduction
 A. Chemistry is the study of matter and energy
 B. Matter is the physical material that makes up everything in
the universe.
 C. All matter is made up of about 115 elements whose basic
building blocks are called atoms
 D. When atoms combine, compounds are formed. Some
compounds are made up of molecules (which contain
covalent bonds), while others form formula units (which
contain ionic bonds).
A. States of Matter
 II. Classification of Matter
State
Shape
Volume
Relative
Kinetic
Energy
Type of
Motion
Solid
Definite
Definite
Low
Vibrational
Liquid
Variable
Definite
Medium
Vibrational
& Rotational
Gas
Variable
Variable
High
Vibrational,
rotational, &
translational
*Vibrational means that the particles simply "vibrate" in one location, but do not change
position.
 *Rotational means the particles can move or flow around one another, but not move freely.
 *Translational means the particles can move freely to just about any location.

Illustration of States/Phases of Matter
B. Pure Substances
 1. Elements
 a. Definition: matter that is composed of only one type of
atom, containing the same number of protons; cannot be
decomposed.
 b. Element Symbols: abbreviations of element names; 1 or 2
letters long (the first of which must be capitalized); for
example: Ca = calcium; O = oxygen
 2. Compound
 a. Definition: a chemical combination of 2 or more atoms in a
fixed ratio; compounds typically exist as molecules (sometimes
we will refer to formula units for ionic compounds); a given
compound always has the same elemental composition (Law
of Constant Composition or Law of Definite
Proportions) b. Examples: H2O, NaCl, C12H22O11
C. Mixtures
 Mixtures – a physical combination of substances;
NO chemical reaction occurs between substances

1. Types of mixtures
a. Homogeneous- a.k.a. solution; components are evenly
distributed throughout the mixtures; same properties throughout.
 *A typical aqueous solution just looks like a glass of water – you
CANNOT see the dissolved components!
 b. Heterogeneous- components are not evenly distributed;
different parts of the mixtures have different properties
 * Different parts or layers are visible.

Mixtures



2. Colloids/Suspensions – mixtures containing particles larger
than normal solutes but small enough to remain suspended in
the “solvent”; a beam of light is visible when allowed to pass
through (called the Tyndall Effect).
3. Alloy – metallic material made by melting together one or
more elements (one of which is a metal); ex: brass & steel
4. Amalgam – a solution of a metal in mercury
Answer:
 Classify the following as elements, compounds,
mixtures. Sometimes, more than one word is
appropriate.
D. Techniques for Separating Mixtures
Distillation Apparatus
 ** Question: What happens to water molecules in a
boiling pot of water?
Properties of Matter
A. Chemical Properties
 1. Definition: Characteristics of a substance exhibited
when it undergoes a change in the atom ratios within
the particles
 2. Examples: reactivity series, flammability, acidity
B. Physical Properties
 B. Physical Properties
 1. Definition: Characteristics of a substance that can be
observed or measured without changing the ratio of the atoms,
like color, volume, mass, conductivity
 2. Examples:
a. Density = mass/volume (can be used as a conversion factor)
 b. Solubility = the amount of a solute that can be dissolved in a
solvent, often water, often measured as “grams per 100 mL of
water “ (can also be used as a conversion factor)
 c. Melting and Boiling Points

C. Extensive vs. Intensive Properties
 1. Extensive properties are those that depend on the
amount of substance present.

Examples are mass & volume
 2. Intensive properties do not depend on the amount
of substance and are therefore used to identify a
substance.

Examples include density, melting and boiling points.
Chemical and Physical Change
 A. Physical Change
 1. Definition: a change in the physical properties of a
substance, such as size, shape, or physical state; chemical
composition NOT changed so the chemical formula, properties
and identity of the substance is the same
 2. Examples: melting, boiling, sublimation, deposition
Chemical and Physical Change
 B. Chemical Change:
 1. Definition: a change in the ratio of atoms in a substance;
also known as chemical reactions;


2. Four common indicators of a chemical change:


** Chemical composition and properties DO change!
color change; gas formation; precipitation; energy change (heat,
light, sound, etc)
3. Examples of chemical changes:
Digestion
 Combustion

Single & double replacement
Synthesis
Decomposition
Formation of water from its elements
Chemical Change
Measurements and Units
Common Derived Units
Comparison of temperature scales
Common glassware to measure volume
1 Liter
Check for Understanding
 * What is the difference between a chemical and a
physical change?
 * What is the difference between an intensive and
extensive property?
 * Name 3 SI base units & 3 derived units. Then tell
what each one measures.
 * What are 4 pieces of glassware used to measure
volume accurately?
Uncertainty in Measurements; Significant
Figures
 A. Accuracy and Precision
 1. Definitions
Precision is the measurement of how closely individual
measurements agree with one another
 Accuracy refers to how close the individual measurements are to
the “true” value.

Percent Error and Standard Deviation
 A. Percent error is used to quantify accuracy (you
compare the actual and experimental results)
% error = Experimental value x 100
True value
 B. Standard deviation is used to quantify precision
(we will not worry about this calculation)
Counting Significant Figures
 1. Rules
 A. Non-zero numbers are always significant
 B. Zeros between non-zero numbers are always significant.
 Zeros before the first non-zero digit (left side zeros) are never
significant. Example: 0.0003 has one significant figure.
 Zeros at the end of a number (right side zeros) are only
significant if there is a decimal in the numeral. Examples:
12.500 has 5 SF; 12500 has 3 SF; 1250.0 has 5 SF
Examples
Calculating Using Correct Significance
 1. Addition and Subtraction
 A. Rule: Your answer will have as many digits after the
decimal point as the measurement with the fewest past the
decimal point.
 B. Examples: Calculate the answer with correct sig figs.
32.4 g + 0.9345 g + 234.65 g = 267.9845 g = 268.0 g
 5788.1 g + 43.76 g = 6261.86 g = 6260 g
 684 mL - 96.3 mL = 587.7 mL = 588 mL
 540 g - 54.3 g = 485.7 g = 490 g

*** When rounding look to the leftmost number that must be
dropped. If it is less than 5 you may truncate; if it is 5 or greater
round the last remaining digit up by one.
Multiplication and Division
 Multiplication and Division
 A. Rule: your answer will have as many sig figs as the
measurement with the fewest sig figs.
 B. Examples:

1.0008 g/mol x 8.632 mol = 8.6389056 g = 8.639 g

35.0971 g / (34.90 cm x 5.12 cm x 23.20 cm) =
0.008466187 g/cm3 = 0.00847 g/cm3


(342.09 g + 53.100 g) / (23 mL - 12.54 mL) = 395.19 g /10. mL
= 39.519 g/mL = 40. g/mL
Sig Figs and Measuring Devices
 When using a non-digital measuring device, always
round one place beyond the smalles increment given
Other information
 Conversion factors are considered exact numbers
and do not affect the precision of any calculations.

Example: 1 m = 100 cm exactly!
 Numbers that are written out in words are
considered exact and do nut affect the precision of
any calculations.

Example: five meters = 5 m exactly!
Check for Understanding
 • If matter is not uniform throughout, then it is a





______________ ________________.
• If matter is uniform throughout, it is ____________.
• If homogeneous matter can be separated by physical
means, then the matter is a ________________.
• If homogeneous matter cannot be separated by physical
means, then the matter is a _________________.
• If a pure substance can be decomposed into something
else, then the substance is a __________________.
What are some ways you can separate a mixture?