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Chapter 1
Chemistry:
The Science of Change
The study of matter,
its properties,
the changes that matter undergoes,
and
the energy associated with these changes.
The Study of Chemistry
Matter is anything that has mass and occupies space.
1.1
The Study of Chemistry
Scientists follow a set of guidelines known as the scientific method:
gather data via observations and experiments
identify patterns or trends in the collected data
summarize their findings with a law
formulate a hypothesis
with time a hypothesis may evolve into a theory
Classification of Matter
Chemists classify matter as either a substance or a mixture of substances.
A substance is a form of matter that has definite composition and distinct properties.
Examples: salt (sodium chloride), iron, water, mercury, carbon dioxide, and oxygen
Substances differ from one another in composition and may be identified by appearance, smell, taste, and other properties.
1.2
A mixture is a physical combination of two or more substances.
A homogeneous mixture is uniform throughout.
Also called a solution.
Examples: seawater, apple juice
A heterogeneous mixture is not uniform throughout.
Examples: trail mix, chicken noodle soup
Chemistry is the Central Science
Theories based on physical principles
Mathematical tools are necessary
Biology, material sciences etc. based on chemistry
Central Science
The Relationship of Chemistry and the Other Disciplines
Why Study Chemistry
Useful
Interesting
A way of thinking
Why Study Chemistry
To be better informed
To be a knowledgeable consumer
To make better decisions for yourself and society
To learn problem-solving skills
To enhance analytical thinking
How to Study Chemistry
Know it
Reason it
Detail of it
Love it
What？
Why？
How？
Chemistry
Inorganic Chemistry 无机化学
Analytical Chemistry 分析化学
Quantitative Chemical Analysis 定量化学分析
Instrumental Analysis 仪器分析
Organic Chemistry 有机化学
Physical Chemistry 物理化学
Structural Chemistry 结构化学
Chemistry Laboratory 化学实验
Classification of Matter
All substances can, in principle, exist as a solid, liquid or gas.
We can convert a substance from one state to another without changing the identity of the substance.
Classification of Matter
Solid particles are held closely together in an ordered fashion.
Liquid particles are close together but are not held rigidly in position.
Gas particles have significant separation from each other and move freely.
Solids do not conform to the shape of their container.
Liquids do conform to the shape of their container.
Gases assume both the shape and volume of their container.
Classification of Matter
A mixture can be separated by physical means into its components without changing the identities of the components.
The Properties of Matter
There are two general types of properties of matter:
1) Quantitative properties are measured and expressed with a number.
2) Qualitative properties do not require measurement and are usually based on observation.
1.3
The Properties of Matter
A physical property is one that can be observed and measured without changing the identity of the substance.
Examples: color, melting point, boiling point
A physical change is one in which the state of matter changes, but the identity of the matter does not change.
Examples: changes of state (melting, freezing, condensation)
The Properties of Matter
A chemical property is one a substance exhibits as it interacts with another substance.
Examples: flammability, corrosiveness
A chemical change is one that results in a change of composition; the original substances no longer exist.
Examples: digestion, combustion, oxidation
The Properties of Matter
An extensive property depends on the amount of matter.
Examples: mass, volume
An intensive property does not depend on the amount of matter.
Examples: temperature, density
Macroscopic, Microscopic & Particulate Matter
Matter can be studied on three levels:
Macroscopic Level: “ Matter that can be seen with the human eye “
Beach Sand, Trees, Cars, Pen, CD, Mountains,
Planets, Galaxies, etc
Length: 101 to 109 meters
Microscopic Level: “ Matter that is too small to be seen by the naked eye, but can be seen under a microscope”
Very small plants, individual bacteria, cellular
structures, DNA Molecule, Semiconductors, etc
Length: 10- 6 meters
Macroscopic, Microscopic & Particulate Matter (cont)
Particulate Level: “ Matter too small to be seen with even the most powerful optical microscope “
Particulate matter consists of the tiny particles
that make up all matter
Molecules, atoms, protons & electron
Length: 10 - 10 meters (1 Angstrom = 10 - 10 meters )
Macroscopic, Microscopic & Particulate Matter (cont)
Scientific Measurement
1.4
Properties that can be measured are called quantitative properties.
A measured quantity must always include a unit.
The English system has units such as the foot, gallon, pound, etc.
The metric system includes units such as the meter, liter, kilogram, etc.
SI Base Units
The revised metric system is called the International System of Units (abbreviated SI Units) and was designed for universal use by scientists.
There are seven SI base units
SI Base Units
The magnitude of a unit may be tailored to a particular application using prefixes.
Mass
Mass is a measure of the amount of matter in an object or sample.
Because gravity varies from location to location, the weight of an object varies depending on where it is measured. But mass doesn’t change.
The SI base unit of mass is the kilogram (kg), but in chemistry the smaller gram (g) is often used.
1 kg = 1000 g = 1×103 g
Atomic mass unit (amu) is used to express the masses of atoms and other similar sized objects.
1 amu = 1.6605378×10-24 g
Temperature
There are two temperature scales used in chemistry:
The Celsius scale (°C)
Freezing point (pure water): 0°C
Boiling point (pure water): 100°C
The Kelvin scale (K)
The “absolute” scale
Lowest possible temperature: 0 K (absolute zero)
K = °C + 273.15
The Fahrenheit scale is common in the United States.
Freezing point (pure water): 32°F
Boiling point (pure water): 212°F
There are 180 degrees between freezing and boiling in Fahrenheit (212°F-32°F) but only 100 degrees in Celsius (100°C-0°C).
The size of a degree on the Fahrenheit scale is only of a degree on the Celsius scale.
Temperature
Temp in °F = ( ×temp in °C ) + 32°F
Derived Units: Volume and Density
There are many units (such as volume) that require units not included in the base SI units.
The derived SI unit for volume is the meter cubed (m3).
A more practical unit for volume is the liter (L).
1 dm3 = 1 L
1 cm3 = 1 mL
d = density
m = mass
V = volume
SI-derived unit: kilogram per cubic meter (kg/m3)
Other common units: g/cm3 (solids)
g/mL (liquids)
g/L (gases)
Derived Units: Volume and Density
The density of a substance is the ratio of mass to volume.
Uncertainty in Measurement
There are two types of numbers used in chemistry:
1) Exact numbers:
are those that have defined values
1 kg = 1000 g
1 dozen = 12 objects
are those determined by counting
28 students in a class
2) Inexact numbers:
measured by any method other than counting
length, mass, volume, time, speed, etc.
1.5
Uncertainty in Measurement
An inexact number must be reported so as to indicate its uncertainty.
Significant figures are the meaningful digits in a reported number.
The last digit in a measured number is referred to as the uncertain digit.
When using the top ruler to measure the memory card, we could estimate 2.5 cm. We are certain about the 2, but we are not certain about the 5.
The uncertainty is generally considered to be + 1 in the last digit.
2.5 + 0.1 cm
Uncertainty in Measurement
When using the bottom ruler to measure the memory card, we might record 2.45 cm.
Again, we estimate one more digit than we are certain of.
2.45 + 0.01 cm
Significant Figures
The number of significant figures can be determined using the following guidelines:
1) Any nonzero digit is significant.
2) Zeros between nonzero digits are significant.
3) Zeros to the left of the first nonzero digit are not significant.
112.1
4 significant figures
305
3 significant figures
0.0023
2 significant figure
50.08
4 significant figures
0.000001
1 significant figure
Significant Figures
The number of significant figures can be determined using the following guidelines:
4) Zeros to the right of the last nonzero digit are significant if a decimal is present.
5) Zeros to the right of the last nonzero digit in a number that does not contain a decimal point may or may not be significant.
1.200
4 significant figures
100
1, 2, or 3 – ambiguous
Determine the number of significant figures in the following measurements: (a) 443 cm, (b) 15.03 g, (c) 0.0356 kg, (d) 3.000×10-7 L, (e) 50 mL, (f) 0.9550 m.
Worked Example 1.1
Solution (a) 443 cm (b) 15.03 g
(c) 0.0356 kg (d) 3.000 x 10-7 L
(e) 50 mL (f) 0.9550 m
Strategy Zeros are significant between nonzero digits or after a nonzero digit with a decimal. Zeros may or may not be significant if they appear to the right of a nonzero digit without a decimal.
3 S.F.
4 S.F.
3 S.F.
4 S.F.
1 or 2, ambiguous
4 S.F.
Think About It Be sure that you have identified zeros correctly as either significant or not significant. They are significant in (b) and (d); they are not significant in (c); it is not possible to tell in (e); and the number in (f) contains one zero that is significant, and one that is not.
In addition and subtraction, the answer cannot have more digits to the right of the decimal point than any of the original numbers.
102.50
+ 0.231
102.731
143.29
- 20.1
123.19
← round to two digits after the decimal point, 102.73
← round to one digit after the decimal point, 123.2
← two digits after the decimal point
← three digits after the decimal point
← two digits after the decimal point
← one digit after the decimal point
Calculations with Measured Numbers
In multiplication and division, the number of significant figures in the final product or quotient is determined by the original number that has the smallest number of significant figures.
1.4×8.011 = 11.2154
11.57/305.88 = 0.0378252
2 S.F.
← fewest significant figures is 2, so round to 11
4 S.F.
← fewest significant figures is 4, so round to 0.03783
4 S.F.
5 S.F.
Calculations with Measured Numbers
Exact numbers can be considered to have an infinite number of significant figures and do not limit the number of significant figures in a result.
Example: Three pennies each have a mass of 2.5 g. What is the total mass
3×2.5 = 7.5 g
Calculations with Measured Numbers
Exact
(counting number)
Inexact
(measurement)
In calculations with multiple steps, round at the end of the calculation to reduce any rounding errors.
Do not round after each step.
Compare the following:
Calculations with Measured Numbers
1) 3.66×8.45 = 30.9
2) 30.9×2.11 = 65.2
1) 3.66×8.45 = 30.93
2) 30.93×2.11 = 65.3
Rounding after each step
Rounding at end
In general, keep at least one extra digit until the end of a multistep calculation.
Perform the following arithmetic operations and report the result to the proper number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L – 2.075 L, (c) 13.5 g ÷ 45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102 g ÷ 4.991x103 g
Worked Example 1.2
Solution (a) 317.5 mL
+ 0.675 mL
318.175 mL
(b) 47.80 L
- 2.075 L
45.725 L
Strategy Apply the rules for significant figures in calculations, and round each answer to the appropriate number of digits.
← round to 318.2 mL
← round to 45.73 L
Perform the following arithmetic operations and report the result to the proper number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L – 2.075 L, (c) 13.5 g ÷ 45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102 g ÷ 4.991x103 g
Worked Example 1.2 (cont.)
Solution
(c) 13.5 g
45.18 L
(d) 6.25 cm×1.175 cm
Strategy Apply the rules for significant figures in calculations, and round each answer to the appropriate number of digits.
← round to 0.299 g/L
= 0.298804781 g/L
3 S.F.
4 S.F.
← round to 7.34 cm2
= 7.34375 cm2
3 S.F.
4 S.F.
Perform the following arithmetic operations and report the result to the proper number of significant figures: (a) 317.5 mL + 0.675 mL, (b) 47.80 L – 2.075 L, (c) 13.5 g ÷ 45.18 L, (d) 6.25 cm x 1.175 cm, (e) 5.46x102 g ÷ 4.991x103 g
Worked Example 1.2 (cont.)
Solution (e) 5.46 x 102 g
+ 49.91 x 102 g
55.37 x 102 g
Strategy Apply the rules for significant figures in calculations, and round each answer to the appropriate number of digits.
= 5.537 x 103 g
Think About It Changing the answer to correct scientific notation doesn’t change the number of significant figures, but in this case it changes the number of places past the decimal place.
An empty container with a volume of 9.850 x 102 cm3 is weighed and found to have a mass of 124.6 g. The container is filled with a gas and reweighed. The mass of the container and the gas is 126.5 g. Determine the density of the gas to the appropriate number of significant figures.
Worked Example 1.3
Solution 126.5 g
– 124.6 g
mass of gas = 1.9 g
density =
Strategy This problem requires two steps: subtraction to determine the mass of the gas, and division to determine its density. Apply the corresponding rule regarding significant figures to each step.
← one place past the decimal point (two sig figs)
1.9 g
9.850 x 102 cm3
← round to 0.0019 g/cm3
= 0.00193 g/cm3
Think About It In this case, although each of the three numbers we started with has four significant figures, the solution only has two significant figures.
Accuracy and Precision
Accuracy tells us how close a measurement is to the true value.
Precision tells us how close a series of replicate measurements are to one another.
Good accuracy and good precision
Poor accuracy but good precision
Poor accuracy and poor precision
Accuracy and Precision
Three students were asked to find the mass of an aspirin tablet. The true mass of the tablet is 0.370 g.
Student A: Results are precise but not accurate
Student B: Results are neither precise nor accurate
Student C: Results are both precise and accurate
Using Units and Solving Problems
A conversion factor is a fraction in which the same quantity is expressed one way in the numerator and another way in the denominator.
For example, 1 in = 2.54 cm, may be written:
1.6
1 in
2.54 cm
2.54 cm
1 in
or
Dimensional Analysis – Tracking Units
The use of conversion factors in problem solving is called dimensional analysis or the factor-label method.
Example: Convert 12.00 inches to meters.
12.00 in ×
Which conversion factor will cancel inches and give us centimeters
1 in
2.54 cm
2.54 cm
1 in
or
= 30.48 cm
The result contains 4 sig figs because the conversion, a definition, is exact.
The Food and Drug Administration (FDA) recommends that dietary sodium intake be no more than 2400 mg per day.
Worked Example 1.4
Solution
2400 mg ×
Strategy The necessary conversion factors are derived from the equalities
1 g = 1000 mg and 1 lb = 453.6 g.
1 lb
453.6 g
453.6 g
1 lb
or
1 g
1000 mg
or
1000 mg
1 g
1 g
1000 mg
1 lb
453.6 g
×
= 0.005291 lb
Think About It Make sure that the magnitude of the result is reasonable and that the units have canceled properly. If we had mistakenly multiplied by 1000 and 453.6 instead of dividing by them, the result
(2400 mg×1000 mg/g×453.6 g/lb = 1.089×109 mg2/lb) would be unreasonably large and the units would not have canceled properly.
An average adult has 5.2 L of blood. What is the volume of blood in cubic meters
Worked Example 1.5
Solution
5.2 L ×
Strategy 1 L = 1000 cm3 and 1 cm = 1x10-2 m. When a unit is raised to a power, the corresponding conversion factor must also be raised to that power in order for the units to cancel appropriately.
1000 cm3
1 L
1 x 10-2 m
1 cm
×
= 5.2 x 10-3 m3
3
Think About It Based on the preceding conversion factors, 1 L = 1×10-3 m3. Therefore, 5 L of blood would be equal to 5×10-3 m3, which is close to the calculated answer.
energy due to the position of the object or energy from a chemical reaction
Potential Energy
Kinetic Energy
energy due to the motion of the object
Energy is the capacity to do work.
Potential and kinetic energy can be interconverted.
Energy is the capacity to do work.
less stable
more stable
change in potential energy EQUALS
kinetic energy
A gravitational system. The potential energy gained when a lifted weight is converted to kinetic energy as the weight falls.
Energy is the capacity to do work.
less stable
more stable
change in potential energy EQUALS
kinetic energy
A system of two balls attached by a spring. The potential energy gained by a stretched spring is converted to kinetic energy when the moving balls are released.
Energy is the capacity to do work.
less stable
more stable
change in potential energy EQUALS
kinetic energy
A system of oppositely charged particles. The potential energy gained when the charges are separated is converted to kinetic energy as the attraction pulls these charges together.
Coulomb Energy
Energy is the capacity to do work.
less stable
more stable
change in potential energy EQUALS
kinetic energy
A system of fuel and exhaust. A fuel is higher in chemical potential energy than the exhaust. As the fuel burns, some of its potential energy is converted to the kinetic energy of the moving car.
Scientific Approach: Developing a Model
Observations :
Natural phenomena and measured events; universally consistent ones can be stated as a natural law.
Hypothesis:
Tentative proposal that explains observations.
Experiment:
Procedure to test hypothesis; measures one variable at a time.
Model (Theory):
Set of conceptual assumptions that explains data from accumulated experiments; predicts related phenomena.
Further Experiment:
Tests predictions based on model.
revised if experiments do not support it
altered if predictions do not support it
A Systematic Approach to Solving Chemistry Problems
Problem statement
Plan
Clarify the known and unknown.
Suggest steps from known to unknown.
Prepare a visual summary of steps.
Solution
Check
Comment and Follow-up Problem
Chapter Summary: Key Points
1
The Scientific Method
States of Matter
Substances
Mixtures
Physical Properties
Chemical Properties
Extensive and Intensive Properties
SI Base Units
Mass
Temperature
Volume and Density
Significant Figures

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