Chemistry
of matter.
Each element is made of a
unique kind of atom.
A compound is made of two
or more different kinds of
elements.
Note: Balls of different colors are used to represent
atoms of different elements. Attached balls represent
connections between atoms that are seen in nature.
These groups of atoms are called molecules.
Matter is anything that occupies space and
has mass.
Atoms are the building blocks
Classifications of Matter
State of Matter
Composition of Matter
Methods of Classification
Au, O2, Fe
Cannot be subdivided
Elements united in
fixed ratios
H2O, CO2
Composition of Matter
Uniform
Composition
Air, Salty water, Alloy
Nonuniform
Composition
Oil-water, milk, blood
States of Matter
2Na(s) + Cl2(g)
2NaCl(s)
A chemical change alters the composition or identity of the
substance(s) involved.
A physical change does not alter the composition or identity of a
substance.
sugar dissolving in tea
ice melting
Types of Changes
Boiling point
Density
Color
Hardness
Pressure
Mass
Volume
State (solid, liquid, gas)
Sucrose
heat
Carbon
+
Water
Chemical Properties can only be observed when a substance is
changed into another substance. Eg. Enthalpy of formation, Heat of
Combustion, Chemical Stability, Flammability
Melting (freezing) point
Temperature
Physical Properties can be observed without changing a substance
into another substance.
Identifying Matter: Types of Properties
Color
Temperature
Density
An intensive property of a material does not depend upon
how much matter is being considered.
Volume
Length
Mass
An extensive property of a material depends upon how
much matter is being considered.
Extensive and Intensive Properties
Numbers play a major role in chemistry. Many topics are
quantitative (have a numerical value).
Concepts of numbers in science
Units of measurement
Quantities that are measured and calculated
Uncertainty in a measurement
Significant figures
Dimensional analysis
Numbers and Chemistry
1 kg = 1000 g = 1 x 103 g
weight force that gravity exerts on an object
weight = g x mass
SI unit of mass is the kilogram (kg)
mass measure of the quantity of matter
Matter is anything that occupies space and has mass.
Mass & Weight
International System of Units (SI)
International System of Units (SI)
A milliliter is a cube 1 centimeter (cm)
long on each side, also called 1 cubic
centimeter (cm cm cm = cm3).
1 mL = 1 cm3
1 L = 1000 mL = 1000 cm3 = 1 dm3
Note that volume is not a base unit for
SI; it is derived from length
(m m m = m3).
The most commonly used metric units
for volume are the liter (L) and the
milliliter (mL).
A liter is a cube 1 decimeter (dm) long on
each side.
Volume
1 m3
Temperature Scales
9x 0C + 32
5
32 0F = 0 0C
212 0F = 100 0C
0F =
K = 0C + 273
273 K = 0 0C
373 K = 100 0C
K = 0C + 273
Inexact (or measured) numbers
depend on how they were
determined. Scientific
instruments have limitations.
Some balances measure to 0.1
g; others measure to 0.0001g.
Exact numbers are counted or given by
definition. For example, there are 12 eggs in
1 dozen.
Numbers Encountered in Science
Different measuring devices have different uses and different
degrees of accuracy.
All numbers that originate from measurements have some
degree of inaccuracy.
Uncertainty in Measurements
Precision how close a set of measurements are
to each other
Accuracy how close a measurement is to the
true value
Accuracy versus Precision
4 significant figures
3 significant figures
1 significant figure
2.00 mg
3 significant figures
Zeros at the end of a number are significant
0.08 L
Zeros to the left of the first nonzero digit are not significant
606 m
Zeros between nonzero digits are significant
1.234 kg
Any digit that is not zero is significant
refers to measured digits
3, zeros between nonzero digits are significant.
3, each digit is a nonzero digit.
f) 7000 mL
4, the number is greater than one so all the
zeros written to the right of the decimal point
count as significant figures.
This is an ambiguous case. The number of significant
figures may be four (7.000 103), three (7.00 103),
two (7.0 103), or one (7 103).
e) 1.310 1022 atoms
d) 0.043 kg 2, same reason as in (c).
c) 0.825 m 3, zeros to the left of the first nonzero digit do not count as
significant figures.
b) 6.01 g
a) 478 cm
Determine the number of significant figures in the following
measurements:
Example
7.2
Rounding off 4.735 to 3 SF
4.74
If the leftmost digit removed is 5 or greater the preceding
number is increased by 1.
Rounding off 7.248 to 2 SF
If the leftmost digit removed is less than 5, the preceding
number is left unchanged.
In rounding of numbers, look at the leftmost digit to be removed.
Significant Figures
round off to 90.4
two significant figures after decimal point
round off to 0.79
3.70
-2.9133
0.7867
one significant figure after decimal point
90.432
89.332
+ 1.1
When addition or subtraction is performed, answers are rounded to
the least significant decimal place.
Addition or Subtraction
Significant Figures
round to
3 sig figs
2 sig figs
round to
2 sig figs
6.8 112.04 = 0.0606926
3 sig figs
4.51 x 3.6666 = 16.536366
= 0.061
= 16.5
The number of significant figures in the result is set by the original
number that has the smallest number of significant figures
Multiplication or Division
Significant Figures
= 6.67333 = 6.67
Because 3 is an exact number
6.64 + 6.68 + 6.70
3
=7
The average of three measured lengths; 6.64, 6.68 and 6.70?
Numbers from definitions or numbers of objects are considered to
have an infinite number of significant figures
Exact Numbers
Significant Figures
88.3 mL
(e) 2.64 103 cm + 3.27 102 cm
(d) 0.0154 kg
c) 8.16 m 5.1355
b)
a) 11,254.1 g + 0.1983 g
Carry out the following arithmetic operations to the correct
number of significant figures:
Example
given unit x
given unit
desired unit
= desired unit
given quantity x conversion factor = desired quantity
We use dimensional analysis to convert one quantity to another.
Most commonly, dimensional analysis utilizes conversion factors (e.g., 1
in. = 2.54 cm).
We use the ratio which allows us to change units (puts the units we
have in the denominator to cancel).
Dimensional Analysis
N is a number
between 1 and 10
N x 10n
n is a positive or
negative integer
0.0000000000000000000000199
1.99 x 10-23
The mass of a single carbon atom in grams:
6.022 x 1023
602,200,000,000,000,000,000,000
The number of atoms in 12 g of carbon:
Scientific Notation
1. Write each quantity with
the same exponent n
2. Combine N1 and N2
3. The exponent, n, remains
the same
Addition or Subtraction
n>0
568.762 = 5.68762 x 102
move decimal left
568.762
Scientific Notation
4.70 x 104
4.31 x 104 + 0.39 x 104 =
4.31 x 104 + 3.9 x 103 =
n<0
0.00000772 = 7.72 x 10-6
move decimal right
0.00000772
1. Divide N1 and N2
2. Subtract exponents n1 and n2
Division
1. Multiply N1 and N2
2. Add exponents n1 and n2
Multiplication
Scientific Notation
8.5 x 104 5.0 x 109 =
(8.5 5.0) x 104-9 =
1.7 x 10-5
(4.0 x 10-5) x (7.0 x 103) =
(4.0 x 7.0) x (10-5+3) =
28 x 10-2 =
2.8 x 10-1
Aspirin is composed of 60.0% carbon, 4.5% hydrogen, and
35.5% oxygen by mass, regardless of its source. Use Figure 1.9
to classify aspirin.
Practice Exercise 2
Which of the following is the correct description of the inside of a grapefruit?
(a) It is a pure compound.
(b) It consists of a homogeneous mixture of compounds.
(c) It consists of a heterogeneous mixture of compounds.
(d) It consists of a heterogeneous mixture of elements and compounds.
(e) It consists of a single compound in different states.
Practice Exercise 1
Because the material is uniform throughout, it is homogeneous. Because its composition differs for the two
samples, it cannot be a compound. Instead, it must be a homogeneous mixture.
Solution
relative amounts of gold and palladium they contain. Both samples are uniform in composition throughout. Use
Figure 1.9 to classify white gold.
Sample Exercise 1.1 Distinguishing among Elements,
Compounds, and Mixtures
(a) How many picometers are there in 1 m?
(b) Express 6.0 103 m using a prefix to replace the power
of ten.
(c) Use exponential notation to express 4.22 mg in grams.
(d) Use decimal notation to express 4.22 mg in grams.
Practice Exercise 2
Which of the following weights would you expect to be
suitable for weighing on an ordinary bathroom scale?
(a) 2.0 107 mg (b) 2500 g (c) 5 10 4 kg (d) 4 106 cg
(e) 5.5 108 dg
Practice Exercise 1
We can find the prefix related to each power of ten in Table 1.4:
(a) nanogram, ng;
(b) microsecond, ;
(c) millimeter, mm.
Solution
What is the name of the unit that equals (a) 10 9 gram, (b) 10 6 second, (c) 10 3 meter?
Sample Exercise 1.2 Using SI Prefixes
Using the volume given in the question, 1.05 103 cm3, and
the
definition of density, we have
In this case each quantity has two decimal places. Thus, the
mass
of the gas, 1.38 g, has two decimal places.
In subtracting numbers, we determine the number of
significant figures in our result by counting decimal places in
each quantity.
To calculate the density, we must know both the mass and the
volume of the gas. The mass of the gas is just the difference in
the masses of the full and empty container:
Solution
A vessel containing a gas at 25 C is weighed, emptied, and then reweighed as depicted in Figure 1.26. From the data
provided, calculate the density of the gas at 25 C.
Sample Exercise 1.9 Determining the Number of Significant Figures in a
Calculated Quantity
If the mass of the container in the sample exercise (Figure 1.26) were measured to three decimal places before and
after pumping out the gas, could the density of the gas then be calculated to four significant figures?
Practice Exercise 2
You are asked to determine the mass of a piece of copper using its reported density, 8.96 g/mL, and a 150-mL
graduated cylinder. First, you add 105 mL of water to the graduated cylinder; then you place the piece of copper in
the cylinder and record a volume of 137 mL. What is the mass of the copper reported with the correct number of
significant figures? (a) 287 g (b) 3.5 10 3 g/mL (c) 286.72 g/mL (d) 3.48 10 3 g/mL (e) 2.9 102 g/mL
Practice Exercise 1
In dividing numbers, we determine the number of significant figures our result should contain by counting the number
of significant figures in each quantity. There are three significant figures in our answer, corresponding to the number
of significant figures in the two numbers that form the ratio. Notice that in this example, following the rules for
determining significant figures gives an answer containing only three significant figures, even though the measured
masses contain five significant figures.
Continued
Sample Exercise 1.9 Determining the Number of Significant Figures in a
Calculated Quantity
Thus, converting from km3 to m3 to L, we have
From the back inside cover, we find 1 L = 10 3 m3, but there is no relationship listed involving km3. From our
knowledge of SI prefixes, however, we know 1 km = 103 m and we can use this relationship between lengths to
write the desired conversion factor between volumes:
Solution
109 km3 of water. Calculate the volume in liters.
Sample Exercise 1.12 Converting Volume Units