Chapter 1 Introduction to Chemistry
 Chemistry is the study of the properties of the substances that make up the
universe and the changes that these substances undergo
 Chemical – of or pertaining to chemistry; substance used in the science of
chemistry
 Matter - anything that occupies space and has mass
 Mass - a measure of the quantity of matter in an object
 Elements – fundamental building blocks of all matter. About 100 or so make
up the majority of materials in the universe. Cannot be broken down into
simpler substances by chemical means.
 Atoms – smallest representative particles of elements. Atoms combine to
form molecules.
 Molecules - aggregates of atoms joined together in a definite arrangement by
chemical forces
Note
difference between mass and weight (F = Mg) -> weight is the
force that gravity exerts on matter
States of matter
 Solid

definite shape/definite volume (rigid)
 Liquid definite volume, but indefinite shape (takes the shape of the
container)
 Gas

indefinite shape and volume (the gas volume depends on the
applied pressure, i.e., gas is compressible)
Note: liquids and solids are not compressible
e.g. water exists as water(s) or ice; water (l); water(g), steam.
However, CO2 only exists as a solid and gas at a pressure of 1 atm
higher pressures, we get CO2 (s), CO2(l), and CO2(g) co-existing.
 Substance - a form of mater that has a fixed composition and distinct
properties
 Homogeneous Substance ->
constant composition throughout
 Heterogeneous Substance
still physically separable
->
components are mixed but are
 Substances have various physical and chemical properties that we use to
characterize and distinguish one substance from another
 Physical property - measured without changing the basic identity of
the substance
 Chemical property - Describes that manner in which one substance
reacts with another substance
e.g. Fe + 3O2 -> Fe2 O3 (rusting)
"Iron + Oxygen gives us rust"
 All measurable properties are either
a) extensive properties - dependent of the amount of matter in the sample;
e.g., mass and volume.
b) Intensive property - the property being measured is independent of the
amount of mater being considered (temp. density melting point etc.) often
used to identify substances
 Physical change - transformation is carved out without changing the basic
identity of the substance e.g. melting ice ->
water and water ->
steam are both examples of physical changes.
 Chemical Change - transform one substance into a chemically different
substance
Mixtures
 homogeneous mixtures - uniform throughout (sugar/water solution)
 heterogeneous mixture - do not have the same composition, properties and
appearance throughout (Iron filings/sand -> easily separated by a magnet)
Elements and Compounds
 Elements - Substance can't be decomposed into simpler substance by
ordinary chemical means
 Compounds - Substances that are composed of two or more elements in
definite chemical proportions (e.g. water H2O -> 2H's 10 not 0.870, 2.29H's
 Elements are the basic building blocks of all matter
3 Classes of Elements
 metals - good conductors of heat/electricity;
(25°C)
all solid at room temp
 nonmetals - poor conductors of heat and electricity; their properties are
usually more varied than those of the metals
 metalloids - properties fall in between the metals and the nonmetals
 Compounds may be broken down into simpler substances (elements)
 the elemental composition of a pure substance is always the same (law
of definite proportions) Joseph Proust ~ 1800.
Units of Measurements
 Chemistry uses SI units (SLIDE OF SI UNITS AND PREFIXES)
 base units such as length (m), temperature (K), mass (g), time (s)
 a # of derived units, e.g., volume (m3), velocity (m/s)
103 kilo
102 deca
10-3 milli etc.
prefixes
Sig Figs and the method of Dimensional Analysis
UNITS OF MEASURE ARE AN INTEGRAL PART OF ANY CHEMICAL
CALCULATION
e.g. the density of methanol is 0.7866 g/mL
What is the volume occupied by 14.37g of CH3OH?
mass
g
g
; 0.7866
= 14.37
volume
mL
v
14.37g
=> V =
= 18.28mL of methanol
0.7866g/mL
Solution
d=
1
 solved via dimensional analysis
 3 of mL of CH3OH = 14.37g CH3OH x 1mL CH3OH / 0.7866gCH3OH
=18.28g methanol
Sig Figs
 Obtain the mass of an NaCl Sample 3.49g
 could be 3.485 g -> 3.4912 g
 there is a fundamental uncertainty in any set of measurements
Two terms
 precision - how well individual measurements agree with one another
 accuracy - how close the measurements are to a "true" or accepted value
e.g. four NaCl measurements
Set 1
3.50g ;
3.51g;
Set 2
3.29g ;
3.29g;
Actual mass of NaCl 3.49g
a) Set 1 accurate/precise
b) Set 2 precise/not accurate
3.51g;
3.28g;
3.52g
3.30g
OPTIONAL
However
add a "0" (2.280g) (3.490g)
 a value between 2.2795 and 2.2804; the "0" is significant (3.489 -> 3.491g)
Rules for significant figures
1. Zeros used to locate decimal points are NOT significant.
2. Exact numbers have an infinite number of zeros (e.g., from definitions).
3. Results from +, - operations should be reported to the same number of
decimal places as that of the least precise term.
4. Results from x,  operations should be rounded off to the same number of
significant figures possessed by the least precise term.
Examples
1. We have 10mg of a sample
0.010g of sample (2 sig figs)
or 1.0 * 101 mg in exponential notation
1.0 * 10-2 g
2. 12.234 + 2.34 = 14.57
2 sig figs after decimal place.
3. 4.83 + 0.1 = 4.9
1 sig fig after decimal place.
4. 4*102 - 30 = 4 * 102
1 sig fig.
5. 4.00*102 - 3 = 3.97 * 102
2 sig figs after decimal place.
# 4 re-done
4 * 102 - 0.3 * 102 = 4 * 102
4.3 x 3.2 = 13.76 = 14.
2 sig. figs.
14.1  13.8 = 1.02 (1.022 rounded off)
6.789 x 0.000032 = 0.000022 = 2.2 x 10-4
(note for this example, your calculator shows 0.000217248)
129 / 13.76 = 9.38
Carrying digits through a long calculation
example:
129 / 13.76 = 9.375 + 10.1 = 19.375 = 19.4
 Usually carry an extra digit through the calculations, and then round off the
first answer.