Transcript CH01 MSJ jlm - Department of Chemistry

Chapter 1
The Nature of Chemistry
Robert Boyle 1627-1691.
Geber (Abu Musa Jabir
ibn Hayyan, ‫ )جابر ابن حیان‬721-815. Originally defined the concept
of chemical “element.”
“Father of Arabic chemistry.”
1
The Study of Chemistry
The Atomic and Molecular Perspective of
Chemistry
• Matter is the physical material of the universe.
• Matter is made up of relatively few (ca. 100) elements.
• Elements are the building blocks of matter.
• On the nano (ultramicroscopic) level, matter consists of
atoms. An atom is a “nano-basketball” -- nano = 10 -9.
• Atoms usually are found in the combined state,
commonly molecules.
• Molecules may consist of the same type of atoms
or different types of atoms.
2
CAUTION!!
Not all compounds are made up of molecules.
Many compounds, for example, are composed
of ionic lattices.
For this chapter, however, we will confine
discussions of compounds to the concept of
molecules, which are the combinations of
nonmetallic elements.
3
The Study of Chemistry
The Molecular Perspective of Chemistry
• In these models, red represents oxygen, white
represents hydrogen, and gray represents carbon.
4
Classification of Matter
Pure Substances and Mixtures
A pure substance cannot be separated into simpler
substances by physical means.
A pure substance has definite and constant chemical
and physical properties (i.e., ignition temp., melting
point, magnetic susceptibility, spectral patterns)
A pure element is a pure substance that consists only of
one kind of atom.
Examples of elements: U, P4, Cl2, C60
A pure compound is a pure substance that consists of
more than one kind of atom.
Examples of compounds: HCl, K2SO4 , C2H6O.
5
Molecules are chemical combinations of two or
more atoms which can all be the same,
e.g., Br2 , H2 , S8 , O3 (these are elements)
or different,
e.g., H2O, C2H5Br (these are compounds)
(Elements which have a different number of atoms in
their molecule are called allotropes.
Allotropes have different physical and chemical
properties, e.g., O2 and O3)
6
Classification of Matter
Individual elements or compounds are all pure substances.
For example, individual samples of C, Fe, O2, N2, H2O
NaCl, K2SO4 are all pure substances.
The composition of a pure substance is constant.
(e.g., water is always 88.9% oxygen and 11.1% hydrogen)
If different samples are mixed together, this is called
a mixture. Mixtures do not have constant
compositions.
7
Classification of Matter
Examples of mixtures:
• air (which contains variable amounts of oxygen,
nitrogen, carbon dioxide, water vapor, pollutants,
etc.
• milk (which contains variable amounts of
galactose, minerals, fats, proteins, etc.)
• carbonated drinks (which contain variable
amounts of carbon dioxide, sugars, flavorings, etc.)
• “tap” water (which contains dissolved minerals)
8
Classification of Matter
Pure Substances and Mixtures
pure substances
9
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 (chemically)
into something simpler, then the substance is a
compound.
10
11
Classification of Matter
Elements
• There are 115 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
(O, C, and H).
12
Classification of Matter
Elements
The next five elements are:
Na 2%, K 2%, Mg 2%,
H 1%, Ti 0.5%.
The next six elements are:
N 3%, Ca 1.5%, P 1%,
K,S,Na 0.75%
13
Elements in the Human Body –
including trace elements
14
Classification of Matter
Use your Periodic Table to refer to all elements and
their chemical symbols.
Bring your Periodic Table
to each class!
15
The Periodic Table
Bring your Periodic Table
to each class!
16
Classification of Matter
Compounds
• Most elements interact to form compounds.
• The proportions of elements in a compound is 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 constant (always the
same). For example, water is always 88.9% oxygen and 11.1%
hydrogen (by mass). For reasons we shall later see, the
volumes of hydrogen and oxygen obtained from water are
always in a fixed 2:1 ratio.
17
Properties of Matter
Physical and Chemical Changes
When a substance undergoes a physical change, its
physical appearance changes, but its chemical nature
does not.
Example: the melting of ice (physical change) results in a
solid being converted into a liquid, but it is still water.
Physical changes do not result in a change of composition.
When a substance changes its composition, it
undergoes a chemical change
Example: 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.
18
Units of Measurement
SI Units
There are two types of units:
a) fundamental (or base) units;
b) derived units.
There are 7 base units in the SI system.
Derived units are obtained from the 7 base SI units.
Example:
unitsof distance
unitsof time
me te rs

se conds
 m/s
Unitsof ve locity
19
20
Units of Measurement
Powers of ten are used for convenience with smaller or
larger units in the SI system.
What is a GigaByte?
21
Units of Measurement - 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 = -273.15oC.
Celsius Scale (used in science)
Also used in science.
Water freezes at 0oC and boils at 100oC.
To convert: K = oC + 273.15.
Fahrenheit Scale (used in US engineering and commerce)
Water freezes at 32oF and boils at 212oF.
To convert:
5
9
C 
9
F - 32 
F 
5
C   32
22
Units of Measurement - Temperature
23
Units of Measurement - Temperature
A user-friendly way to view the Celsius Scale:
0° - Cold! (coat)
10° - Cool (sweat shirt)
20° - Pleasant (long sleeves)
25° - Room temperature (short sleeves)
30° - Very warm (T-shirt)
40° - Hot! (swimming pool!)
24
Units of Measurement - Volume
The units for volume are
given by length3
SI unit for volume is 1 m3.
1 mL = 1 cm3
(REMEMBER THIS!)
Other volume units:
1 L = 1 dm3
=1000 cm3
= 1000 mL
25
Units of Measurement - Density
•
•
•
•
Used to characterize substances.
Defined as: density = mass /volume.
Units: g/cm3, also known as specific gravity.
Originally based on mass -- the density was defined as
the mass of 1.00 g of pure water.
26
Uncertainty in Measurement
Precision and Accuracy (actually discussed in Chapter 2)
• Measurements that are close to the “correct” value are
accurate.
• Measurements which are close to each other are
precise.
• Measurements can be:
accurate and precise;
precise but inaccurate;
neither accurate nor precise.
27
Uncertainty in Measurement
Precision and Accuracy
Example: Weigh yourself on your bathroom scales.
Your scales are cheap, and the graduations are every 5
pounds. The precision is bad.
Your scales are expensive, and the graduations are every
0.1 pound. The precision is good.
You know your correct weight from a high-quality scale
at the doctor’s office, to within 0.1 pound. You weigh
yourself on your bathroom scales and find the reading
is off by 10 pounds. The accuracy is bad.
Same situation as above, and find the reading is within
0.1 pound. The accuracy is good.
28
Dimensional Analysis
Method of calculation utilizing a knowledge of units.
• Conversion factors represent “1” (unity).
• Conversion factors are simple ratios.
Example:
1 m or 100 cm
100 cm
1m
Example:
1L
or 1000 mL
1000 mL
1L
are equivalent
are equivalent.
29
Dimensional Analysis
Example: Convert 5.6 km to m:
5.6 km
x 1000 m
km
= 5600 m
Volume conversion. Convert 1 cubic meter to mL.
1 m3 x (100 cm)3 x 1 mL = 1 x 106 mL
(1 m)3
1 cm3
30