Chemistry
EIGHTH EDITION
by
Steven S. Zumdahl
Susan Zumdahl
Chapter 1
Chemical Foundations
An Overview


Chemistry: Fundamentally concerned with
how one substance changes to another.
Science:
 A process for understanding nature and
its changes
 Primary objective is to explain
phenomena of the physical world
An Overview

Scientific method – the process that lies at
the center of scientific inquiry. Steps in the
Scientific Method are:
1. Making observations (collecting data)
2. Making a prediction (formulating a
hypothesis)
3. Performing experiments to test the
prediction (testing the hypothesis)
An Overview

Observations:
 Qualitative- does not involve a number
 Quantitative- involves both a number
and a unit

Hypothesis: A possible explanation for an
observation. An educated guess
An Overview

Theory (model): A set of tested hypotheses
that gives an overall explanation of some
natural phenomenon.
 Is an interpretation
 An attempt to explain why it happens

Law is a summary of observed behavior.The
same observation applies to many different
systems (Law of Conservation of Mass)
 A law summarizes what happens
The Various Parts of the Scientific Method
Which of the following is the correct order of
steps to establish a theory?
A.
B.
C.
D.
Conducting experimental work, collecting
observations, making a hypothesis, establishing a
theory.
Establishing a theory, conducting experimental work,
collecting observations, making a hypothesis.
Making a hypothesis, collecting observations,
establishing a theory, conducting experimental work
Collecting observations, making hypothesis,
conducting experimental work, establishing a theory.
Ans. ‘D’
Units of Measurement
Measurement: Compare an unknown quantity
with one that is known
 A Quantitative measurement consists of two
parts
 A number
 A unit
Both parts must be present for the measurement
to be meaningful. Example: 20 grams

Units of Measurement


Measurements – mass, length, time,
temperature etc.
Different systems were adopted in different
parts of the world.
 English system
 Metric system
Units of Measurement
In 1960, an international agreement set up a
system of units called the International
System or the SI system. This system is
based on the metric system and units derived
from the metric system.
Mass/kilogram(kg), length/meter(m)
Volume/not a fundamental SI unit but is
derived from length. Cubic meter (m3)-more
commonly used is liter/(L).
Prefixes
p = pico (0.000000000001)
n = nano (one billionth)
µ = micro (0.000001)
m = milli (0.001)
k = kilo (1000)
M = mega (1000000)
G = giga (1000000000)
10-12
10-9
10-6
10-3
103
106
109
Uncertainty in Measurements





There is a limit to the accuracy of any measurement.
The digit that must be estimated is called uncertain
A measurement always has some degree of
uncertainty
We usually report a measurement by recording all
the certain digits plus the first uncertain digit
(Significant figures)
The accuracy of any result can not be any better
than the data it came from
The overall accuracy is determined by the least
accurate number used in the computation
Measurement of Volume Using a Buret.The volume is read
at the bottom of the liquid curve (called the meniscus).
Precision
Precision: Refers to the degree of agreement
Among several measurements of the same
quantity. Reflects the reproducibility of a
given type of measurement.
1st series of measurements: 34, 35, 37, 37, 38
2nd series of measurements: 30, 35, 40, 42, 47
The precision of the 1st series is better than
the 2nd
Accuracy
Accuracy: Refers to the agreement of a
particular value with the true value
Average = Sum of all the values / number
of values
example: 34, 35, 37, 37, 38
sum of all the values = 181
number of values = 5
average = 36.2
true value = 36.0  good accuracy
The Difference between Precision and Accuracy
Errors

Random Error (indeterminate error): A
measurement has an equal probability of
being high or low. This type of error occurs
in estimating the value of the last digit of
measurement.

Systematic Error (Determinate error): This
type of error occurs in the same direction
each time. It is either always high or always
low, often resulting from poor technique.
Rules for Counting Significant Figures






All non-zero digits are significant figures: 1234  4 significant
figures
Zeros between non-zero digits (captive zeros) are significant
figures: 205  3 significant figures
Zeros beyond decimal point at the end of the number (trailing) are
significant figures: 0.24000  5 significant figures
Zeros preceding the first non-zero digit in a number are not
significant figures: 0.00453  3 significant figures
Zeros at the end of whole numbers are not significant figures unless
you are given information to the contrary: 3400  2 significant
figures, 34.00 X 102  4 significant figures, 3400.  4 significant
figures.
Exact numbers can be assumed to have an infinite number of
significant figures: 2 in 2r (the circumference of a circle), 4 & 3 in
4/3r3 (the volume of a sphere), 1 in = 2.54 cm.
How many significant figures are in each of
the following?
a.
b.
c.
d.
e.
f.
g.
h.
12  2 significant figures (S.F.)
1098  4 S.F.
2001  4 S.F.
2.001 x 103  4 S.F.
0.0000101  3 S.F.
1.01 x 10-5  3 S.F.
1000.  4 S.F. (because of the decimal point)
22.04030  7 S.F.
Rules for Significant Figures in
Mathematical Operations
Multiplication and division: The number
of significant figures in the result is the
same as that in the quantity with the
smallest number of significant figures.
Ex.: 4.56 x 1.4 = 6.38  6.4
(3 S.F.) (2 S.F.) (3 S.F.) (2 S.F.)
The product should have only two
significant figures since 1.4 has two
significant figures.
Addition and subtraction: The result has
the same number of decimal places as the
least precise measurement used in the
calculation.
Ex.: 12.11 + 18.0 + 1.013  31.128  31.1
The correct result is 31.1, since 18.0 has
only one decimal place.
Rules for Rounding
In a series of calculations, carry the extra digits
through to the final result, then round.
If the digit to be removed
a. Is less than 5, the preceding digit stays the
same. Ex. 1.33 rounds to 1.3.
b. Is equal to or greater than 5, the preceding
digit is increased by 1. Ex. 1.36 rounds to
1.4.
Example
Perform the following mathematical operation, and
Express each result to the correct number of significant
figures.
•
•
•
•
97.381 + 4.2502 + 0.99195  102.62315 
102.623
171.5 + 72.915 - 8.23  236.185  236.2
(9.2 x 100.65) / (8.321 + 4.026)  925.98 / 12.347
 74.996  75
8.27 (4.987 - 4.962)  8.27(0.025)  0.20675
0.21
Dimensional Analysis

Dimensional analysis converts a given
result from one system of units to another

The method involves using conversion
factors to cancel units until you have the
proper unit in the proper place.

How do you convert 1.53 minutes to seconds?
a. Find a conversion factor (or factors) :
60 sec = 1 min
b. Set up start-up and ending information with
units 1.53 min. = sec
c. Decide which unit to put on top. We need
an answer in ‘sec’ and we need to get rid of
‘min’. Therefore, 1.53 min X 60 sec/1 min
= 91.8 sec.



To convert from one unit to another, use the
equivalence statement that relates the two
units
Derive the appropriate unit factor by looking
at the direction of the required change (to
cancel the unwanted units).
Multiply the quantity to be converted by the
unit factor to give the quantity with the
desired units.

Convert 10.0 km to miles
Equivalence statements:
1 km = 1000 m, 1 m = 1.09 yd, 1760 yd = 1 mi
1000m 1094
.
yd
1mi
10.0kmX
X
X
 6.22mi
1km
1m
1760 yd

Convert 50 miles/hours to feet/minute
50miles 5280 ft 1hour
ft
X
X
 4400
 4.4 X 103 ft/min
hour
mile
60 min
min
Temperature Measurement
Three systems for measuring temperature
1. The Celsius scale (oC)
2. The Kelvin scale (K)
3. The Fahrenheit scale (oF)
The size of the temperature unit (the degree) is
the same for Celsius and Kelvin scales, the
difference is in their zero points.
Temperature (Kelvin) = temperature (Celsius)
+ 273.15
TK = TC + 273.15
Temperature (Celsius) = temperature (Kelvin)
– 273.15
TC = TK – 273.15
The Three Major Temperature Scales

Both the degree sizes and zero points are different
for Fahrenheit and Celsius scales – need two
adjustments
 One for degree size
 One for the zero point
180 oF = 100 oC
Unit factor is 180 oF /100 oC or 9 oF / 5 oC or
reciprocal
Zero points: 32 oF = 0 oC (subtract 32 from the
Fahrenheit temperature)
(TF – 32 oF ) 5 oC / 9 oF = TC
TF = TC X 9 oF / 5 oC + 32 oF
A person has a temperature of 102.5 oF. What is
this temperature on the Celsius scale? On the
Kelvin scale?
TC = (TF -32 oF) 5 oC / 9 oF
= (102.5 oF - 32 oF) 5 oC / 9 oF = 39.2 oC
TK = TC + 273.2 = (39.2 + 273.2 ) K = 312.4 K
Density
Density: The mass of substance per unit volume
of the substance
Density = mass(g)/volume(cm3)  g/cm3
Density is often used as an “identification tag”
for a substance.
Matter: Anything occupying space and having
mass. Matter exits in three states:
1. Solid is rigid, it has a fixed volume and shape.
2. Liquid has a definite volume but no specific
shape, it assumes the shape of its container.
3. Gas has no fixed volume or shape, it takes on
the shape and volume of its container.
Types of Mixtures

Homogeneous mixture: Having visibly
indistinguishable parts. Physical properties
are the same throughout the material. A
homogeneous mixture a solution (example:
vinegar).

Heterogeneous mixtures: Having visibly
distinguishable parts. Physical properties are
different at different points in a material
(example: bottle of ranch dressing).
Definitions




Pure substance is one with constant composition.
Pure substances can be isolated by separation
techniques – distillation, filtration, chromatography.
Compound is a substance with constant composition
that can be broken down into elements by chemical
processes. Example: electrolysis of water produces
hydrogen and oxygen.
Physical change is a change in the form of a
substance but not in its chemical composition.
Chemical change: a given substance becomes a new
substance or substances with different properties and
different composition
Definitions



Elements are substances that cannot be decomposed
into simpler substances by chemical or physical
means.
Intensive properties: the value is independent on
the amount of material present (color, temperature,
melting point, freezing point, density).
Extensive properties: the value is dependent on the
amount of material present (weight, length, volume).
The Organization of Matter
Summary









Scientific method: Consists of looking at the world, developing
hypothesis about how it operated, and then making prediction based on
these hypotheses.
Measurement: Compare an unknown quantity with a known quantity.
SI unit: International System unit (kg, m, sec).
Uncertainty: Limit to the accuracy of measurement.
Significant figures: Count the number of significant Figures and calculate
after mathematical operation.
Dimensional analysis: Converts from one system of units to another.
Temperature conversion: TC = TK – 273.15
TC = (TF – 32 oF) 5 oC/9 oF
Density = m (g)/V (cm3)
Definitions: matter, homogeneous, heterogeneous, physical change,
chemical change elements, intensive and extensive properties.