Transcript No Slide Title

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.
Molecules may consist of the same type of atoms or
different types of atoms.
The Molecular Perspective of Chemistry
The Study of Chemistry
Why Study Chemistry
• Chemistry is central to our understanding of other
sciences.
• Chemistry is also encountered in everyday life.
• It is the basis behind many medications
• “Form defines Function”

Classification of Matter

States of Matter
Matter can be a gas, a liquid, or a solid.





Gases have no fixed shape or volume.
Gases can be compressed to form liquids.
Liquids have no shape, but they do have a volume.
Solids are rigid and have a definite shape and volume.
Classification of Matter

Pure Substances
Atoms and Compounds


Molecules can consist of more than one type of element.



Molecules can have only one type of atom (an element).
Molecules can have more than one type of atom (a compound).
Pure substances are



Atoms consist only of one type of element.
The same throughout
Have a set ratio
Mixtures
Two types

Solutions
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.
Pure Substances
and Mixtures
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.
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).
•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).
• 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
Compounds
• Most elements interact to form compounds.
• 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 ratio by mass of the elements in a chemical compound is always the
same, regardless of the source of the compound.
– The law of constant composition can be used to distinguish between
compounds and mixtures of elements:
– Compounds have a constant composition; mixtures do not.
– Water is always 88.8% O and 11.2% H by weight regardless of its
source.
–Brass is an example of a mixture of two elements: copper and zinc. It
can contain as little as 10%, or as much as 45%, zinc.
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 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

Physical and Chemical Changes
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



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
Properties of Matter




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

Powers of ten are used for convenience with smaller or
larger units in the SI system.
Units of Measurement
SI Units
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
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).
Uncertainty in Measurement





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.
Uncertainty in Measurement
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





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
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?