Transcript Classification of Matter

CHEMISTRY
The Central Science
10th Edition
Chapter 1
Introduction: Matter &
Measurement
Why Study Chemistry
• Chemistry is the study of the properties of materials and
the changes that materials undergo.
• Chemistry is central to our understanding of other
sciences.
• Chemistry is also encountered in everyday life.
Chemistry: Catastrophe Prevention?
The space shuttle Columbia disintegrated in 2003 upon reentry into
the Earth’s atmosphere due to a damaged thermal protection system.
The Study of Chemistry
•
•
•
•
•
The Molecular Perspective of Chemistry
Matter is the physical material of the universe.
Matter is made up of relatively few elements.
On the microscopic level, matter consists of atoms and
molecules.
Atoms combine to form molecules.
As we see, molecules may consist of the same type of
atoms or different types of atoms.
Molecular Perspective of Chemistry
Classification of Matter
•
•
•
•
•
•
States of Matter
Matter can be a gas, a liquid, or a solid.
These are the three states of matter.
Gases take the shape and volume of their container.
Gases can be compressed to form liquids.
Liquids take the shape of their container, but they do
have their own volume.
Solids are rigid and have a definite shape and volume.
Classification of Matter
Pure Substances and Mixtures
• Elements consist of a unique type of atom.
• Molecules can consist of more than one type of element.
– Molecules that have only one type of atom (an element).
– Molecules that have more than one type of atom (a compound).
• If more than one atom, element, or compound are found
together, then the substance is a mixture.
• Pure
Substances
and Mixtures
Classification of Matter
•
•
•
•
•
Pure Substances and Mixtures
If matter is not uniform throughout, then it is a
heterogeneous mixture.
If matter is uniform throughout, it is homogeneous.
If homogeneous matter can be separated by physical
means, then the matter is a mixture.
If homogeneous matter cannot be separated by physical
means, then the matter is a pure substance.
If a pure substance can be decomposed into something
else, then the substance is a compound.
Classification of Matter
•
•
•
•
•
•
Elements
If a pure substance cannot be decomposed into
something else, then the substance is an element.
There are 114 elements known.
Each element is given a unique chemical symbol (one or
two letters).
Elements are building blocks of matter.
The earth’s crust consists of 5 main elements.
The human body consists mostly of 3 main elements.
Classification of Matter
Elements
Classification of Matter
Elements
• Chemical symbols with one letter have that letter
capitalized (e.g., H, B, C, N, etc.)
• Chemical symbols with two letters have only the first
letter capitalized (e.g., He, Be).
Classification of Matter
Compounds
• Most elements interact to form compounds.
• Example, H2O
• The proportions of elements in compounds are the same
irrespective of how the compound was formed.
• Law of Constant Composition (or Law of Definite
Proportions):
– The composition of a pure compound is always the
same.
Classification of Matter
Compounds
• If water is decomposed, then there will always be twice
as much hydrogen gas formed as oxygen gas.
• Pure substances that cannot be decomposed are elements.
Classification of Matter
Mixtures
• Heterogeneous mixtures are not uniform throughout.
• Homogeneous mixtures are uniform throughout.
• Homogeneous mixtures are called solutions.
Properties of Matter
Physical vs. Chemical Properties
• Physical properties can be measure without changing the
basic identity of the substance (e.g., color, density, odor,
melting point)
• Chemical properties describe how substances react or change
to form different substances (e.g., hydrogen burns in oxygen)
• Intensive physical properties do not depend on how much of
the substance is present.
– Examples: density, temperature, and melting point.
• Extensive physical properties depend on the amount of
substance present.
– Examples: mass, volume, pressure.
Properties of Matter
Physical and Chemical Changes
• When a substance undergoes a physical change, its
physical appearance changes.
– Ice melts: a solid is converted into a liquid.
• Physical changes do not result in a change of
composition.
• When a substance changes its composition, it undergoes a
chemical change:
– When pure hydrogen and pure oxygen react completely, they
form pure water. In the flask containing water, there is no
oxygen or hydrogen left over.
Properties of Matter
Physical and Chemical Changes
Properties of Matter
Separation of Mixtures
• Mixtures can be separated if their physical properties are
different.
• Solids can be separated from liquids by means of
filtration.
• The solid is collected in filter paper, and the solution,
called the filtrate, passes through the filter paper and is
collected in a flask.
Properties of Matter
•
•
•
•
Separation of Mixtures
Homogeneous liquid mixtures can be separated by
distillation.
Distillation requires the different liquids to have different
boiling points.
In essence, each component of the mixture is boiled and
collected.
The lowest boiling fraction is collected first.
Separation of Mixtures
Units of Measurement
•
•
•
•
Separation of Mixtures
Chromatography can be used to separate mixtures that
have different abilities to adhere to solid surfaces.
The greater the affinity the component has for the
surface (paper) the slower it moves.
The greater affinity the component has for the liquid, the
faster it moves.
Chromatography can be used to separate the different
colors of inks in a pen.
Units of Measurement
SI Units
• There are two types of units:
– fundamental (or base) units;
– derived units.
• There are 7 base units in the SI system.
Units of Measurement
Base SI Units
Units of Measurement
SI Units
Selected Prefixes used in SI System
Class Practice Examples
• What is the name given to the unit that equals (a)
10-9 grams; (b) 10-6 second; (c) 10-3 meter
• What fraction of a meter is a nanometer?
Units of Measurement
SI Units
• Note the SI unit for length is the meter (m) whereas the SI
unit for mass is the kilogram (kg).
– 1 kg weighs 2.2046 lb.
Temperature
There are three temperature scales:
• Kelvin Scale
–
–
–
–
Used in science.
Same temperature increment as Celsius scale.
Lowest temperature possible (absolute zero) is zero Kelvin.
Absolute zero: 0 K = -273.15 oC.
Units of Measurement
Temperature
• Celsius Scale
– Also used in science.
– Water freezes at 0 oC and boils at 100 oC.
– To convert: K = oC + 273.15.
• Fahrenheit Scale
– Not generally used in science.
– Water freezes at 32 oF and boils at 212 oF.
– To convert:
5
C  F - 32 
9
9
F  C   32
5
Class Practice Example
• Make the following temperature conversions:
(a) 68 oF to oC; (b) -36.7 oC to oF
5
C  F - 32 
9
9
F  C   32
5
Units of Measurement
Temperature
Units of Measurement
Derived Units
• Derived units are obtained from the 7 base SI units.
• Example:
units of distance
Units of velocity 
units of time
meters

seconds
 m/s
Units of Measurement
Volume
• The units for volume are
given by (units of
length)3.
– SI unit for volume is 1
m3.
• We usually use 1 mL = 1
cm3.
• Other volume units:
– 1 L = 1 dm3 = 1000 cm3 =
1000 mL.
Units of Measurement
Volume
Units of Measurement
Density
• Used to characterize substances.
• Defined as mass divided by volume:
mass
Density 
volume
• Units: g/cm3.
• Originally based on mass (the density was defined as the
mass of 1.00 g of pure water).
Class Practice Examples
• Answer the following problems:
• (a) Calculate the density of mercury if 1.0 x 102 g
occupies a volume of 7.36 cm3.
• (b) Using the density for mercury, calculate the
mass of 65.0 cm3 of mercury.
Uncertainty in Measurement
• All scientific measures are subject to error.
• These errors are reflected in the number of figures
reported for the measurement.
• These errors are also reflected in the observation that two
successive measures of the same quantity are different.
Precision and Accuracy
• Measurements that are close to the “correct” value are
accurate.
• Measurements that are close to each other are precise.
Precision and Accuracy
Uncertainty in Measurement
Significant Figures
• The number of digits reported in a measurement reflect
the accuracy of the measurement and the precision of the
measuring device.
• All the figures known with certainty plus one extra figure
are called significant figures.
• In any calculation, the results are reported to the fewest
significant figures (for multiplication and division) or
fewest decimal places (addition and subtraction).
Uncertainty in Measurement
•
•
•
•
•
Significant Figures
Non-zero numbers are always significant.
Zeros between non-zero numbers are always significant.
Zeros before the first non-zero digit are not significant.
(Example: 0.0003 has one significant figure.)
Zeros at the end of the number after a decimal place are
significant.
Zeros at the end of a number before a decimal place are
ambiguous (e.g. 10,300 g).
Dimensional Analysis
• Method of calculation utilizing a knowledge of units.
• Given units can be multiplied or divided to give the
desired units.
• Conversion factors are used to manipulate units:
• Desired unit = given unit  (conversion factor)
• The conversion factors are simple ratios:
desired unit
Conversion factor 
given unit
Dimensional Analysis
Using Two or More Conversion Factors
• Example to convert length in meters to length in inches:
Number of in  number of m   conversion m  cm  
conversion cm  in 
100 cm
1 in
Number of in  number of m  

m
2.54 cm
Class Practice Problem
• A person’s height is measured to be 67.50 in. What is
this height in centimeters?
• Perform the following conversions: (a) 2 days to s; (b)
20 Kg to g.
Dimensional Analysis
•
•
•
•
Using Two or More Conversion Factors
In dimensional analysis always ask three questions:
What data are we given?
What quantity do we need?
What conversion factors are available to take us from
what we are given to what we need?