Chapter 1: Matter and Measurement
Scientific Method: Observation =>
hypothesis
↓
Supports hypothesis ← experimentation → discredits hypothesis → rejected
↓
↓
More supportive experiments ←----------------- modify hypothesis
│
Theory ←---------┴------ (measurable observations stand test of time)-----→ Law
Chemistry- The study of matter, its transformations and the energy changes that
accompany the transformations.
Why study chemistry?
1. Improve health(helpful pharmaceutical drugs, i.e. aspirin, chemicals that fight
diseases).
O
ll
O-C-CH3
CO2H
ASPIRIN
2. Enhance food products(i.e., better fertilizers for more abundance of crops; make
food that are less fattening like sugar substitutes).
3. Protection from and for the environment(synthetic polymers for better clothes,
better houses; using chemicals to decontaminate natural resources, i.e. chelating
chemicals.
4. Improve living conveniences(plastics, better metal composites for making lighter
but stronger cars).
Matter- is the physical material of the universe (and beyond). It is anything that has
mass and occupies space. There are three states (or phases) of matter: solid, liquid and
gas.
State of Matter
Solid
Liquid
Gas
Characteristics of Matter
Volume
Entropy(disorder) Shape
definite
Low(very ordered) definite
definite
moderate
indefinite
indefinite Very high
indefinite
Compressability
Low(dense)
Low(dense)
High(low
density)
Matter can be pure or a mixture.
Pure Matter(Substance)- is matter that has distinct properties and invariant
composition from one sample to another.
Atom- is the building block of matter.
Two Types of Pure Matter (Substance)
Elements
Can not be chemically decomposed
into simpler substances; on the
molecular level, each element is
composed of only one kind of
atom. Examples are members of the
periodic table, i.e., Hydrogen gas,
Helium gas, Oxygen gas, Iron
metal, etc.
Compounds
Can be chemically decomposed
into simpler substances; they are
composed of two or more different
kinds of atoms. Examples are HI
(hydroiodic acid), H2O(water),
NaCl (table salt), H2SO4
Molecule- Substance made up of at least two atoms chemically bonded. Not all
molecules are compounds; but all compounds are molecules.
Mixture- Comprises two or more different kinds of pure substances; yet each
substance retains its chemical identity throughout the mixture.
Two Types of Mixtures
Homogeneous Mixture(Solutions)
Has a high level of uniformity
(same phase) throughout the
mixture. Examples are table salt
dissolved in water, alcohol and
water mix, sugar water or soda pop,
any unreactive mixture of gases like
air (N2 + O2 + CO2), Chemical
solutions, Metal alloys(bronze =
copper + tin, brass = copper + zinc.
Heterogeneous Mixture
Lacks uniformity throughout the
mixture (different phases).
Examples are mixture of salt and
black pepper, ice water (solid and
liquid phases can be seen), oil and
water mix.
Because each substance in a mixture retains its identity and its properties, we can
separate the mixed substances from one another by taking advantage of their retained
properties. For example, a heterogeneous mixture of sand and salt can be separated by
dissolving the mixture in water and decanting the salt water homogeneous mixture
from the sand (solubility property). Then we can separate the salt from the water by
boiling off the water (boiling point property).
Properties of Matter
Physical Properties- are properties one can measure without destroying
the chemical identity or composition of pure matter. Examples are
boiling point, melting point, color, odor, density.
Chemical Properties- are properties that are determined by a chemical
reaction of the matter. Examples are flammability, acidity and basicity,
aromaticity of a hydrocarbon.
Intensive Properties- do not depend on the amount of substance being
studied. Examples are boiling point, melting point, density, color , odor.
Extensive Properties- do depend on the the amount of substance
studied. Examples are mass, volume, size, enthalpy.
Physical Changes of Matter- appearance changes but chemical
composition does not. For example, all changes of state(solid to liquid,
liquid to gas, solid to gas) are physical changes of matter.
Chemical changes of Matter- chemical composition changes; a reaction
takes place and a new product is formed. Such changes are indicated by
burning or combustion, oxidation, evolution of gas, formation of a
precipitate.
Measurements
Type of Measure
Mass
Length
Volume
Time
Temperature
Amt of substance
Electric Current
Energy
Pressure
System Int.(SI)
English Units
Kilogram(kg)
Pound(lb)
Meter(m)
Yard(yd)
m3 / Liter(L)
Second(s)
Kelvin(K)
Mole
Ampere
Joule
Pascal
Quart(qt)
Second(s)
Farenheit(F)
Mole
Ampere
BTU
Atmosphere
Conversion
Factor
1kg = 2.2lbs(or
453.6g = 1lb)
1m = 1.09yds(or
2.54cm = 1in
exactly)
1L = 1.06qt
same
K=5/9(F-32)+273
same
same
9.5 x 10-4 btu/J
1 atm/ 1.013 x 105
Pa
Prefixes
Based on a standard unit like the meter, Kilo means 1000, mili means 1000th, deka
means 10, deci 10th, hecto means 100, centi means 100th. Other prefixes to know:
Prefix
Giga
Mega
nano
pico
femto
Magnitude
109
106
10-9
10-12
10-15
Temperature- the measure of coldness or hotness(or the measure of the average
kinetic energy of molecules). We use a thermometer to measure temperature.
Temperature units are Kelvin, Celcius and Farenheit. Conversions to know :
C = 5/9(F-32)
K = C + 273
K = 5/9(F-32) + 273
Volume- is the measure of the amount of space matter occupies. Units for volume are
liter, mililiter(1000th of a liter), cubic centimeter(cc).
1 liter = 1 decimeter cubed = (0.1m)3 = 1000th meter cubed
1 decimeter = 10 centimeter = 10cm
Therefore, 1 liter = 1 decimter cubed = (10cm)3 = 1000(cm)3 = 1000mL
Density- mass of substance divided by its volume, usually in units of grams per
mililiter (g/mL) or grams per cubic centimeter(g/cc).
The density of water is about 1g/mL
Precision- is a measure of how closely your several quantitative measurements of a
sample agree (reproducibility) and it is usually measured as standard deviation.
( - Xi)/n;  = average; Xi = each measurement; n = number of measurements.
Ex. 10.01mL, 10.02mL,10.04mL, 10.05mL => Xi, χ = 10.03mL, n = 4
S.D. = │10.03-10.01│ + │10.03-10.02│+ │10.03-10.04│+ │10.03-10.05│= 0.02
4
Χ ± S.D. = 10.03mL ± 0.02mL
Accuracy- is a measure of how closely your several measurements agree with the
true or theoretical value and it is usually measured as percent yield (% yield).
% yield = (Yi/Y)  100; Yi = experimental value; Y = Theoretical value.
If a chemical reaction yields 56g of product and theoretical yield is 62g, then
% yield is 56.00g/62.00g x 100 = 90.32%
Significant Figures
To avoid ambiguity in numerical values reported by different researchers who
perform the same experiment, scientists employ the practice of significant figures.
This simply means determining which zeros, if any, are important (significant) in an
experimentally obtained numerical value.
1. All nonzero digits are always significant.
2. Zeros between nonzero digits are always significant.(i.e. 1005, 0.200005,
0.002005)
3. Assume a reported value X with a decimal point. If X < 0.1, the zero(s) at the
beginning of the number are not significant(i.e. 0.002, 0.00003, 0.00000004 ).
4. If X < 1, the zeros that come after the nonzero digits are significant(i.e. 0.900,
0.200000).
5. Zeros sandwiched between a nonzero digit on the left and a decimal point on the
right are always significant (i.e. 6000.)
6. When a number ends in zeros but contains no decimal point, the zeros may or
may not be significant. This ambiguity can be remedied by using scientific
notation and the judgement of the reader. For example the value 103000g can be
written these ways:
1.03  105 g (3 sig. Fig.)
1.030  105 g (4 sig. Fig.)
Rounding Off to Significant Figures
If the left most digit to be dropped is less than 5, the preceeding number is left
unchanged.
i.e. round off 7.248 to 2 significant figures. 7.248  7.2
i.e. round off 3.7883974 to 4 significant figures. 3.7883974  3.788
If the leftmost digit to be dropped is 5 or greater, increase the preceeding number
by one.
i.e. round off 4.735 to 3 significant figure. 4.735  4.74
i.e. round off 2.376 to 2 significant figures. 2.376  2.4
In multiplication and division operations, the result should have the same number of
significant figures as the measurement with the least number of significant figures
i.e. 6.221cm  5.2cm = 32.3492cm2  2 sig fig.  32cm2
In addition and subtraction, the result can have no more decimal places than the
measurement with the fewest decimal places.
i.e. 20.4 + 1.322 + 83 = 104.722
The value 83 has the fewest decimal places (it has none). So the resulting value
should have no decimal places, 105
Exact Numbers
Exact numbers have an infinite number of decimal places. So they are not used to
determine the number of decimal places of a resultant value. For example, there 24
students in your class this semester, not 25.1 or 24.5 or 24.99. For example
1000g/1kg is an exact ratio value. The ratio 2.54cm/in. is an exact ratio value.
Dimensional Analysis
In converting a measurement from metric to English unit(or vice versa), one uses as
many conversion factors as is needed in “bookkeeping” set called dimensional
analysis. The conversion factor are set up so that the unit one immediately wants to
convert to can not be cancelled out.
2.54cm = 1in
conversion factor is (2.54cm/1in.) or (1in./2.54cm)
i.e convert 8.50in. to centimeters: 8.50 in│2.54 cm = 21.6cm
│1 in
Convert 23 femtometers to centimeters.
23 fm│1 m
│102cm = 2.3 x 10-12 cm.
15
│10 fm │1 m
OR
23 fm│ 10-15 m │1 cm = 2.3 x 10-12 cm.
│1 fm
│10-2 m
Density of water is 1g/mL; convert units to ounces per cubic inches(oz/ in3)
1g │1 kg │2.2lbs │16oz │1 mL │ ( 2.54cm)3 = 0.58 oz./in3
1mL│1000g │ 1 kg │1 lb │1(cm)3 │1 in3