When Reading a Piece of Scientific
Equipment:
-
Be sure to note the uncertainty of the measurement
-
Be sure to give all significant figures in the measurement.
Examples: Take a look at the following pieces of equipment:
A Centigram Balance:
So, if you look at the
balance, you can see that the number of certain digits include:
o The
hundreds, the tens, the ones, the tenths, and the hundredths.
o In
other words, you can read for certain 128.19
§
That means 5 certain digits
o Since
all measurements need one unknown digit, the last thousandths digit can be estimated
§
Therefore there would be 6 significant digits in this measurement
§
I would estimate the thousandths place to be a 2
§
Ergo, the total measurement would be:
·
128.192 g
§
Do not forget to list the uncertainty:
·
128.192 +/- .001 g
Try This One:
-
I got: 3.5 +/- .1 C
-
Although there are more lines present, the divisions are not certain enough to allow to you estimate the hundredths
place.
-
You could choose a higher uncertainty value here: +/- .5
2-2 Temperature:
-
Why can not the sense of touch be used to measure temperature?
-
A thermometer is an instrument that gives and accurate and precise reading of temperature.
-
Galileo Galilei (1564-1642) – invented the first temperature instrument
o Modern
thermometers have a bulb filled with mercury or colored alcohol attached to a stem
o Heating
causes liquid to expand and move up the stem
o Cooling
causes liquids to condense and move down the stem
The Fahreheit and Celsius Temperature
Scales
-
Gabriel Fahrenheit- Made thermometers in the late 1600’s and early 1700’s- made up his own temperature
scale.
-
Anders Celsius (1701-1744) developed a scale much more in tune with the metric system
o Freezing
point at sea level = 0 Boiling point at seas level = 100
The Kelvin Temperature Scale
-
The SI scale used to measure temperature is the Kelvin Scale
-
Lord Kelvin (English- 1824-1907) : unit K
o A
degree change of 1 K is the same as a degree change of 1 C
o Zero
point in the Kelvin scale corresponds to absolute zero (-273 C)
§
Absolute zero is where molecular motion stops
-
Some Equations
o C=
K – 273
o K=
C + 273
o (oF-32oF)
x (100oC/180oF) = oC
o (oC
x 180oF/100oC) + 32oF = oF
2-3 Matter
-
Matter is the “stuff” of which things are made
o Has
mass (amount of stuff) and volume (amount of space)
-
Do not know where the “stuff” came from, but it is here and we have learned a lot about the “properties
of matter”
o Has
been a philosophical issue for millennia
States of Matter:
-
Four States of Matter:
o Solid
o Liquid
o Gas
o Plasma
-
Properties of the different states (generalized):
o Solid:
§
High density
§
Density affected little by changes in pressure
§
Shape not affected by the shape of a container
§
Orderly arrangement of particles (ie. Crystals)
o Liquid:
§
High density
§
Density affected little by changes in pressure
§
Adopts the shape of the container
o Gas:
§
Low density
§
Density depends on the pressure
§
Expands to fill the container
o Plasma
§
Low density
§
Density depends on pressure
§
Expands to fill the container
§
Exists only at high temperatures
Changes in State:
-
Can observe changes in states by heating or cooling a substance
Ex: Water at 0 C is changing from liquid to solid
Water at 100 C is changing from liquid to gas
Water from 0 C to 100 C is in the form of a liquid
Ex Mercury at –39 C is changing from a liquid to a solid
Mercury at 357 C is changing from a liquid to
a gas
Mercury from –39 C to 357 C is a slippery
liquid
Properties of Matter:
-
A sample of matter can be identified by observing its characteristics or properties
-
Physical Properties: properties that can be observed without changing the identity of the substance. (density/
color/ melting point, etc)
1)
State: (at standard temperature and pressure): Liquid, Solid, Gas
2)
Quantity: mass, volume, density
3)
Color
4)
Texture
5)
Melting and boiling points
6)
Conductivity
7)
Solubility in Water
-
Chemical Properties: properties that cannot be observed without changing the identity of the substance (flammability,
etc)
1)
Evolution of a gas
2)
Formation of a precipitate
3)
Absorb or Gives off heat
4)
Emission of light
5)
Color Change
Atomic Number: number
of protons in a nucleus (is equal to the number of electrons)
Atomic
Mass: average of all naturally occurring isotopes
Periodic Law: physical
and chemical properties of an atom are periodic functions of the atomic number.
Changes in Matter
-
Physical Changes: changes that do not alter the identity of the substances
o Crushing,
tearing, and changes in state
-
Chemical Changes: changes that do alter the identity of the substance
o Change
in the chemical make up of a substance.
Conservation of Matter:
-
Antoine Lavoisier: “one may take it for granted that in every reaction there is an equal quantity of matter
before and after”
- Antoine
Lavoisier: 1800’s (1743-1794)
-
Mass of substances before a chemical change was always equal to the mass of substances after the change.
-
Conclusion:
o Matter
was neither created nor destroyed during a chemical reaction.
Became known as
the Law of Conservation of Mass
Ex:
By mass,
1 g of H always binds with 8 g of O
So 2 g of H will bind with 16 g of O
3 g of H will bind with 24 g of O
-
Since you know the mass of both reactants, you can figure out the mass of the products:
o 1g
H + 8g O = 9g water
-
Knowing this, you can reverse the reaction:
o Electrolysis-
using electricity to break water
45.0 grams of water:
broken via electrolysis yielded 5.0g H and how many grams O?
45.0g water- 5.0g H= 40g O
** one
of the most important principles in chemistry**
-
Lavoisier got the ax during the Reign of Terror that followed the French Revolution
2.4 Elements and Compounds
Elements:
-
An element is a substance that cannot be separated into simpler substances by a chemical change
o Over
100 known elements
o Named
for famous people, states, planets, countries etc.
-
Element Symbol: a one or two letter abbreviation
o First
letter is always capitalized; second letter is always lower case
o Most
abbreviations come from the English name, others come from the Latin origin.
§
Copper – cuprum Cu
§
Gold – aurum Au
§
Iron – ferrum Fe
§
Lead – plumbum Pb
§
Mercury – Hydragyrum Hg
§
Potassium – Kalium K
§
Silver – Argentum Ag
§
Sodium – Natrium Na
§
Tin – Stannum Sn
§
Tungsten – Wolfram W
-
Periodic table- simple, harmonic, rhythmic way of organizing the elements by innate properties
Compounds
-
A compound is a substance that contains two or more elements combined in a fixed proportion
-
Chemical Symbols are used to represent compounds- merely putting the element symbols in a specific order which
notes the number of each element present.
Distinguishing between elements
and compounds
-
Elements and compounds are pure substances
o Has
a unique set of chemical and physical properties
o Separation
techniques like electrolysis help distinguish between the two
o Careful
measurements of mass help distinguish as well.
2-5 Mixtures
-
A mixture is a blend of two or more pure substances
Types of Mixtures:
-
A mixture that has visibly different parts is called a heterogeneous mixture
-
A mixture that does not have visibly different parts is called a homogeneous mixture
Separating the Components of
a Mixture:
-
Filtration- separation of heterogeneous mixtures of liquids and solids
-
Distillation- separation of homogeneous mixtures of liquids based on different boiling points (one changes to
gas form first)
-
Distillation may also be used to separate impurities from liquids- solids are left behind
-
Crystallization produces solids of very high purity by evaporating the liquid component.
-
Chromatography- separation by flowing along a stationary substance.]
-
All measurements must include both a unit and a number.
o Without
the unit, the number has no meaning.
-
English vs. Metric System:
o English
system – feet, inches, etc are not used in science.
o Metric
system- the international system of measurement is used
§
Common language for all scientists
§
Easy conversions
|
SI Base Unit |
|
|
|
|
|
|
|
Physical Quantity |
|
Unit Name and Symbol |
|
mass |
|
kilogram, kg |
|
length |
|
meter, m |
|
time |
|
second, s |
|
count, quantity |
|
mole, mol |
|
temperature |
|
kelvin, K |
|
electric current |
|
ampere, A |
|
luminous intensity |
|
candela, cd |
|
|
|
|
|
Derived Units Commonly Used in Chemistry |
|
|
|
|
|
Physical Quantity |
|
Unit Name and Symbol |
|
area |
|
square meter |
|
volume |
|
cubic meter |
|
force |
|
newton, N |
|
pressure |
|
pascal, Pa |
|
energy |
|
joule, J |
|
power |
|
watt, W |
|
voltage |
|
volt, V |
|
frequency |
|
hertz, Hz |
|
electric charge |
|
coulomb, C |
The International
System of Units (SI)
- Seven base units (shown above)
Definitions:
o
Length: distance that light travels in a vacuum during a time interval 1/299,792,458 of a second
o
Mass and weight:
§
Mass: amount of material- about 2.2 lbs at sea level
§
Weight: influence of the force of gravity on mass
o
Area and Volume (derived units – combinations of base units)
§
Area = Length x Width
5.0
m x 3.0 m = 15 m2
·
Both units and numbers are multiplied in the answer
§
Volume: amount of space that an object occupies
Non- SI
Units Used Frequently in Chemistry
- Volume: liter, L (there are exactly
1000 L in one cubic meter)
- Pressure: atmosphere, atm; millimeters
of mercury, mm Hg
- Temperature: Celcius degree
- Energy: calorie, cal
Metric
Prefixes
- Prefixes added to the base unit that
make the units larger or smaller
Prefixes
that make the Unit Larger
o
kilo (1 km = 1000 m)
o
mega (1Mm = 1000000 m)
Prefixes
that make the Unit Smaller
o
deci (1 dm = .1 m or
10 dm = 1m)
o
centi (1cm = .01m)
o
milli (1mm = .001 m)
o
micro (1mm = .000001 m)
o
nano (1 nm = .000000001 m)
o
pico (1pm .000000000001 m)
1-5 Uncertainty
in Measurement
***
When making a measurement, you must give all the certain (or exact) digits that the instrument can give + one additional “uncertain”
digit that you estimate***
- Measuring instruments always have some
built in flaws
- Measurement always involves some estimation
On an
electronic balance:
- The measurement is done for you. The last number on the screen is the uncertain digit
o
Be sure to include the unit as well as the measurement
On a scale:
- Imagine a graduated cylinder- measures
volume
o
Liquids curve in the cylinder (known as the meniscus)
o
Measurements are taken from the bottom of the bend
- The certain digits are those given on
the graduated cylinder.
o
The uncertain digit is found by reading between the lines of the instrument
The Uncertainty
of a Measurment
- Generally, the uncertainty of a measurement
reflects the value of the uncertain digit:
Suppose
the measurement was: 32.7 mL
§
The seven in this measurement is uncertain.
§
Since the seven is uncertain, it could possibly as high as an 8 or as low as a 6.
§
Therefore the uncertainty is +/- .1
Reliability
in Measurement
- Precision- several measurements that
are close in value
- Accuracy- how close a measurement is
to the accepted value (a standard)
Reliable
measurements have both high precision and high accuracy.
Can
you have one without the other???
