Chapter 1
 How
is chemistry applied to the
matter that makes up the world
around us?
 What
does the science of
chemistry study?
 How does matter change?


Chemistry is the study of matter- its
properties and changes it goes through
Atoms are the most basic
component of matter.

A chemical is any substance that has definite
composition
◦ Natural– water (H2O), carbon dioxide (CO2)
◦ Synthetic– plastics, teflon (tetrafluoroethylene)

They’re everywhere! Some good, some bad



Solid, liquid, gas, plasma
Macroscopic – can see with your eye (macro =
large)
Microscopic – cannot be seen with the unaided
eye (micro = small)
solid
liquid

gas
Figure 2 pg.6 – atoms in each state
Solid
Liquid
Gas
Fixed volume
Fixed volume
Indefinite
volume
Fixed shape
Indefinite
shape
Indefinite
shape
Rigid structure,
Loosely held
organized,
together, slide
vibrate in place past each other
Independent,
weak
attractions

Physical Changes
◦ Identity of matter stays the same.
 Arrangement, shape, or location might
change
◦ Examples:
 Cutting, dissolving, boiling, crushing

Identity of matter changes.
◦ New substances are formed = chemical
reaction
◦ Reactants – starting materials
◦ Products – ending materials
◦ Example:
mercury(II) oxide → mercury + oxygen
1.
2.
3.
4.
Evolution of a gas (bubbles, odor)
Formation of a precipitate (solid formed
from 2 solutions/liquids – looks cloudy)
Release or absorption of energy (change in
temperature or light given off)
Color change
 How
can we measure and describe
matter?


Matter – anything that has mass and takes up
space (has volume)
Volume – a measure of the size of an object
◦ Uses mL, L, cm3 as units (1mL = 1cm3)
◦ Measured using ruler or
graduated cylinder


Mass- measure of the quantity of matter in
an object
◦ Uses: kg, g, mg
Weight – measure of the gravitational force
exerted on an object
◦ Can change with location
 Gravity changes
◦ Units: Newtons (N)



Earth’s gravity is 9.8m/s2
The moon’s gravity is 1.6m/s2
Jupiter’s gravity is 25m/s2

A person weighs 150lb on Earth
The same person weighs 25lb on the moon
And he/she weighs 360lb on Jupiter

These are NOT masses!






Unit – quantity adopted as a standard of
measurement
Quantity – something that has size,
magnitude or amount
SI – Système International (international
system)
Not the same as the metric system!
1960 decided to unify measurement
worldwide
◦ 7 base units
 All others are a combination of these
 Use prefixes to change magnitude
Quantity
Symbol
Unit
Abbreviation
Length
l
Meter
m
Mass
m
Kilogram
kg
Time
t
Second
s
Temperature
T
Kelvin
K
Amount of a substance
n
Mole
mol
King Henry Died Belly-Up Drinking
Chocolate Milk Under (a) Nice Picture.
Prefix
Abbreviation
Power of Ten
Meaning
kilo-
k
103
1000
hecto-
h
102
100
deka-
da
101
10
100
1
Base Unit
deci-
d
10-1
0.1
centi-
c
10-2
0.01
milli-
m
10-3
0.001
micro-

10-6
0.000 001
nano-
n
10-9
0.000 000 001
pico-
p
10-12
0.000 000 000 001

We can form other necessary units from the
seven SI base units.
◦ Examples:
 Speed = distance/time = m/s
 Area = length ∙ width = m2
 Volume = l ∙ w ∙ h = m3
 Use mL in chemistry
 1L = 1000 mL = 1000 cm3
(1mL = 1cm3)
 Why
identify properties?
◦ They help us to learn more about substances.
◦ The more we know, the better we understand
the substance.
 How
are properties identified?
◦ By making observations about substances.

Physical vs. Chemical
◦ Physical Property – a property that can be
measured or observed without changing the
chemical identity of the substance
 Examples: density, color, hardness, transparency (how
well light is able to pass through), melting point
◦ Chemical Property – property that describes a
substance’s ability to chemically interact with
another substance
 Determined by trying to cause a chemical change
 Examples: flammability, reactivity with water, ability
to be oxidized

Density = mass divided by volume = m/V
◦ Units = g/mL or g/cm3
◦ Found by measuring mass and volume
 Volume – use a ruler (l∙w∙h) or water
displacement (graduated cylinder)
◦ Graph results show density by slope of line:
rise = ∆y = mass
run
∆x
volume
 Can be used to identify substances.
 How
do we classify matter?

Atom – smallest unit of an element that maintains
the properties of that element.
◦ Ex: copper, nickel, boron, etc.
◦ Almost all elements are made up of many
individual atoms; but some are made of atoms
that bond together: Br, I, N, Cl, H, O, F

Pure substance – an element or compound that has
definite chemical and physical properties (in other
words has a chemical formula)
◦ Element – cannot be broken down into simpler
substances; contain only 1 kind of atom
◦ Molecule – 2 or more atoms bonded together that
retain the chemical & physical properties of that
element or compound
 Elements: H2, O2, N2, Cl2, Br2, I2, F2 (same type of
atoms bond together)
 Compounds: H2O, CO2, HCl (different types of
atoms bond together)

Allotrope: different form of an element
*Examples: oxygen = O2, ozone = O3, diamond, graphite
diamond
graphite
carbon

Compound – substance made up of atoms of
two or more different elements joined by
chemical bonds.
◦ Can be further classified as ionic or covalent
to be discussed in later chapters)
◦ Examples: salt (NaCl), caffeine (C8H10N4O2)

Mixture – combination of two or more substances
that are NOT chemically combined
◦ Can vary in composition and properties
◦ Homogeneous – uniform throughout,
often called a solution
 Examples: completely dissolved salt water,
tea, stainless steel, maple syrup
◦ Heterogeneous –not evenly mixed
 Examples: not completely dissolved salt water,
orange juice with pulp, chocolate chip cookie,
granite, salad
Matter
Mixture
Homogeneous
separate
physically
Pure
Substance
Heterogeneous
Compound
Element
separate
chemically

Helpful hints for classification:
◦ Homogeneous (solutions): cannot see the
individual substances that make it up (same
uniform appearance & composition
throughout).
◦ Heterogeneous: consist of visibly different
substances or states.
Difference #1:
 Properties of compounds do not reflect
properties of elements making it up.
◦ Properties are different from the elements.
Properties of mixtures do reflect properties of
substances making them up.
Difference #2:
 Compounds always have definite
compositions of elements.
 Mixtures can have varying amounts of
substances making them up.
