Download Unit 1: Fundamental Chemistry

Unit 1: Fundamentals of Chemistry
CHEMISTRY:
the science of materials, their composition and
structure, and the changes they undergo.
Six Common Branches of Chemistry
Physical chemistry -Theoretical chem, the
physics contained in chem.
Analytical chemistry - The Best Branch!
Measurement and analysis of materials.
Organic chemistry - Study of covalent carbon
compounds.
Inorganic chemistry - Study of all other
elements and compounds.
Biochemistry - Chemical basis of life.
Nuclear chemistry -Properties and reactions of
atomic nuclei.
Foundation of Chemistry
To study chemical systems and the
CHANGES they undergo.
Initial state →Final State
A + B → AB
All of the principles of chemistry
are based on the study of chemical
reactions
REACTION: a chemical change in which a new
substance is formed.
NaCl + AgOH → AgCl + NaOH
Left side called the REACTANTS
Right side called the PRODUCTS
An important objective of science:
Relate properties of Large samples of matter
(called macroscopic) to the most basic particles of
matter (microscopic)
Ex: What is happening when water boils?
Scientific Method – used all the
time
Step 1: Making Observations
Two types of Observations:
QUALITATIVE: a descriptive term. “Your
shirt is red,” “The solution was bubbling and
was pink”, “The water is a liquid at room
temperature.”
QUANTITATIVE: a quantitative
observation is called a MEASUREMENT.
“The pressure was 1 atm”
Scientific Method
Step 2: Looking for patterns in the observations
Usually results in the formulation of a natural
law
NATURAL LAW: a statement that expresses
generally observed behavior. A natural law is
often expressed as a math formula
Ideal Gas Law: pV = nRT
Scientific Method
Step 3: Formulating Theories
THEORY: May be called a model. It consists
of a set of assumptions put forth to explain the
observations. Can be considered as a
hypothesis.
REMEMBER: an observation is a FACT.
A law is a statement of observations.
A theory is an interpretation or explanation (can
be wrong!!)
Scientific Method
Step 4: Experiment to Test Theories
Experiments may and usually do lead to
modified or changed theories.
Scientific Method
Observations
↓
LAW
↓
Theory ← Modify Theory
↓
↗
Test Theory/Experiment
Measurements:
Think of a whole number greater than 2 but less
than 4.
But, do not think of a graphical or verbal
representation of that number!
Units of Measurement
A MEASUREMENT consists of
two parts: a NUMBER and a
UNIT. Both must be present.
A number without units is
meaningless.
There are two types of units.
FUNDAMENTAL UNITS
These are units upon which all other units are
based.
METER – length
GRAM – measures mass. Mass – quantity of
matter that a body possesses.
WEIGHT – measure of the earth's gravitational field.
Remember: your mass is FIXED but your weight varies
depending on your position from the center of the earth
FUNDAMENTAL UNITS
SECOND – measures time (based on the
vibration of Cesium-133)
MOLE – measures the number of particles and is
equal to 6.02 x 10 to the 23rd.
KELVIN – named after Lord Kelvin. Kelvin
temperature scale is based on absolute zero.
COULOMB – a quantity of electrical charge
DERIVED UNITS
Derived units are units based on
fundamental units. There are lots and
lots of derived units. Two examples:
VOLUME – 1 mL = 1 cubic centimeter = 1 gram
(if water)
1 Liter = 1 cubic decimeter
1 Liter – 1000 cubic centimeters
DERIVED UNITS
Density is another derived unit based
on mass and volume
Density = Mass/Volume
Units for Density are grams/mL
UNCERTAINTY IN MEASUREMENT
All measurements have
some degree of
uncertainty.
How much do you weigh?
5 different honors chemistry
students massed a sample of iron:
student 1 = 16.18 g
student 2 = 16.15 g
student 3 = 16.19 g
student 4 = 16.16 g
student 5 = 16.15 g
Which decimal place do you think was likely to be
rounded? Most exact?? Least exact??
Significant Figures: certain digits
and the first uncertain digit.
(the real reason we need sig figs is to
help us figure out how precisely a
measurement was made)
Sig Figs are important in all fields
of science..
Measuring the sides of a square:
area = (side)(side)
area = (16.4 cm) (22.8 cm)
area = 373.92 cm2
Look at the answer the calculator gives us. It is
IMPOSSIBLE
to have an
answer that is MORE accurate than our
measurements – thus the need for sig figs in the
physical sciences.
(for a plain orange pumpkin to become a golden carriage)
A semi is weighed to be 42,000 lbs (with driver).
The driver picks up a 5 lb cat from the side of the
road. What is the new total weight of the truck?
The rules for SigFigs give a
method for scientists to
indicate how precisely a
value is known.
Rules for Counting Sig Figs
1. Non Zero Integers
Non zero integers always count as significant figures
ex: 3.455 has 4 sig figs
2. Zeros
There are THREE (really 4) rules for Zeros:
1. LEADING ZEROS - are zeros that precede all of the non-zero
digits. They DO NOT count as sig figs. Note that leading zeros
are always in a very small number- Less than 1.
ie. The number 0.000456 has 3 sig figs. The leading zeros are
not significant and are only there to simply indicate the position
of the decimal point.
2. CAPTIVE ZEROS (OR SANDWICHED ZEROS) – are
zeros between two non-zero digits. They are ALWAYS
significant.
ie. the number 1.008 has FOUR sig figs.
3. TRAILING ZEROS – these are zeros at the right end of the
number. There are two rules for trailing zeros:
a. They ARE significant if the number has a decimal point.
b. They are NOT significant if there is NO decimal point.
The number 100 has only 1 sig fig
The number 1.00 x 102 has three sig figs
The number 2306.00 has six sig figs
Now you have fun and practice!!
Determine the # of sig figs in:
236
678.09
1.008
0.000056709
8,900
0.00509080700
Rules for Math and Sig Figs
1. Addition/Subtraction
The result has the same number of decimal places as the
least precise measurement. HINT: Count Decimal Places.
ex. 12.11
18.0 ← here is the limiting term-only 1 dec. place
+ 1.013
31.123
But the CORRECT answer with one decimal place would be 31.1
Essentially, you cannot add or subtract a known digit to or
from an unknown digit!
2. Multiplication/Division
The number of sig figs in the product/quotient is
the same as the number of sig figs in the LEAST
precise measurement. HINT: count the sig figs
ie. (4.56)(1.4) = 6.38 …but you can’t really have
an answer with MORE sig figs than the number
with the least…so the CORRECTED answer
would be 6.4.
(3 sig figs)(2 sig figs) = 2 sig figs
Now YOU get to have fun!!! Give
the answers to the correct # of sig
figs
1.
2.33 + 4.5 + 8.00 + 8 =
2.
9.010 ÷ 3.7 =
3.
9.0 – 3.888 =
4.
(5.66)(1.00)(2.00)(0.0006) =
5.
(7.24+1.5) (9.023) =
Precision vs. Accuracy
These concepts are often confused!!!
ACCURACY – denotes the nearness of a
measurement to its accepted value.
actual beaker mass = 19.0 grams
Your mass = 19.9 g, 24.1 g, and 13.6 g.
How was this student's accuracy?????
PRECISION
An agreement between the numerical values of a
set of measurements that have been made the
same way (think CONSISTENCY!!)
Ex: Beaker mass = 19.0 g
Your mass = 14.1 g, 14.0 g, and 14.1 g
How was your precision?
How was your accuracy?
Dart example
Percentage Error Formula
% error = |experimental - actual|
actual
x 100
PERCENTAGE ERROR
A student was calculating the % of lead (Pb) in
the water at Xenia High School in the drinking
fountains. She came up with the following
values: 16.12%, 16.14%, 16.12% and 16.13%.
The average value was 16.13%. The correct
value according to my scientific calculations was
16.49%.
What can be said about the accuracy?
What can be said about the precision?
Calculate the % error.
SCIENTIFIC NOTATION
Also called exponential notation
Move the decimal to the left – exponent is larger
and POSITIVE!! For example the speed of light
is 30,000,000,000 cm/sec. Put into scientific
notation.
Move the decimal to the right – exponent is
smaller and negative. For example, put 0.000496
m into scientific notation.
To convert
a number to scientific notation:
1. Write only the sig figs.
ex: 103,000,000
write 103
2. Put in a decimal point so there is only 1
digit to the left
ex: 1.03
3. Count the number of spaces that the
decimal point moved from its starting point.
Use this number as the exponent.
ex: 1.03 x 108
4. If the original number was a decimal (less
than 1, greater than 0) then the exponent is
negative. If the original number was greater
than 1 then the exponent is positive.
Fun with Scientific Notation
402,000 =
0.000701 =
100.2 =
(9.24 x 1016 )(6.12 x 1014 ) =
1.96 x 10-8 /2.47 x 10-4 =
DIMENSIONAL ANALYSIS
You will need to become expert at this.
An exciting and fun way of working problems by
using the UNITS to help us along the way.
Defined: a method of changing units by using
conversion factors.
(use the metric/English or English/metric charts)
A conversion factor can be any mathematical
relationship between two units.
Ex: 12 in / 1 ft
45 miles / 1 hr
30 days / 1 month
2.65 g / 1 mL
27 students / 16 desks
1. Set up a grid. What you are given goes to the topleft. Split top and bottom if 2 units!
2. Find conversion factor (or factors) to convert
between given units and needed units.
3. Insert conversion factors into the grid so that
units cancel (top to bottom) so that only needed
units are left.
Doing conversions using
Dimensional Analysis:
Convert 14 Kg to lbs:
Convert 16.9 in to cm:
Convert 8 years to seconds:
Convert 3 gallons to mL
Convert 8 mph to cm/second
Convert 4.66 in2 to cm2
Convert 98.77 yd3 to m3
Convert 4.5 m to Km
Convert 0.455 mL to cL
TEMPERATURE
Three systems: Celsius, Kelvin, Fahrenheit
For Water: BP = 212ºF, 100ºC, 373K
For Water: FP = 32ºF, 0ºC, 273 K
Special Formulas
°C = (°F – 32)5/9
°F = (°C X 9/5) + 32
K = °C + 273
°C = K – 273
OR, my Favorite: (C+40)9/5 = (F+40)
Normal body temperature is
98.6ºF. Convert to ºC and Kelvin.
Liquid nitrogen has a boiling point
of 77K. Convert this to ºF.
DENSITY
Density is defined as the mass of a substance per
unit volume.
Density = M/V
This formula can also be solved for mass and
volume.
M=
V=
OR!!! Use dimensional analysis!!!
The mass of Al is 14.2 g and the
volume is 6.9 mL. Find the density
Calculate the % error (the actual
density is 2.7 g/mL)
The density of Fe is 7.86 g/mL.
You have 29 grams of Fe. How
many mls will it occupy?
Percentage Problems
Percentage is
part/whole x 100
Given: 82 g of a metallic powder. It consists of 31
g of Zn, 3 g Ag, and 48 g of Sn. Find the % of
each.
1.0 lb of salt is dissolved in 1.0 gallon of water.
What is the density of the solution? What percent
of the solution is salt?
1.0 lb of salt is dissolved in 1.0 gallon of
water. What is the density of the
solution? What percent of the solution is
salt?
Flow Chart of Matter
MATTER
Pure Substance
Heterogeneous
Mixture
Mixtures
Homogeneous
Mixture
Pure Substance
Elements
Compound
Atoms
Nucleus
Electrons
Neutrons
Protons
Quarks
Quarks
SEPARATION METHODS
There are Nine (9) ways to separate mixtures in
the lab. Some of these are based on physical
properties and some of these are based on
chemical properties.
1. FILTRATION
Separates based on insoluble/soluble properties
FILTRATE: the soluble substance or liquid that
passes through the filter paper
RESIDUE: the insoluble chunky “stuff” that
remains in the filter paper.
Filtration is a great way to separate a
SUSPENSION: where the particles are larger
than molecular size in the liquid.
SOLUBILITY
Solubility in water is a physical property
SOLUBLE: dissolves
INSOLUBLE: remains undissolved
2. DECANTING
Decanting is used to separate a coarse suspension
of liquid and dense, insoluble solids.
Decanting simply means “to pour off”
Yes...even you can do this separation technique!
3. SIMPLE DISTILLATION
Distillation is used to separate solid solute from
liquid solvent
Distillation is used to make distilled water and
many different alcohol products.
Distillation is based on a phase difference (the
solid remains in the original flask and the liquid
boils, evaporates, then condenses and drips into a
new container in a purified form)
4. Fractional Distillation
Used to separate miscible liquids
MISCIBLE LIQUIDS – liquids that are soluble in
each other (alcohol in water)
IMMISCIBLE LIQUIDS- liquids that are NOT
soluble in each other (oil in water)
Fractional distillation separates based on the
boiling points of the liquids
4. Fractional Distillation
Example: you have a mixture of two liquids –
alcohol and water. Alcohol has a boiling point of
80°C and water has a boiling point of 100°C.
The liquids boil off one by one at their boiling
point temperatures.
Fractional distillation is used in the petroleum
industry (petroleum products)
Many products come from crude oil: drugs (legal
ones of course), cosmetics, kerosene, oil, gas,
plastics, etc.
Petroleum products based on fossil fuels
5. FRACTIONAL
CRYSTALLIZATION
Separates based on soluble solids whose
solubility differs in hot and cold water
Example: solid X is very soluble at all
temperatures
Solid Y is soluble only in hot water
Dissolve both in hot water; cool the water; solid
Y comes out of solution because Y is insoluble in
cold water.
6. EXTRACTION
Uses a device called a SEPARATORY FUNNEL
Extraction is used to separate immiscible liquids
(like oil and water)
IMMISICIBLE LIQUIDS – not soluble in each
other
MISCIBLE LIQUIDS – liquids that are soluble in
each other
7. Chromatography
Separates substances by differences in dissolving
rates
Chromatography is used to separate COLORS
and PROTEINS
Chromatography is used to do various analysis' of
DNA (paternity tests, etc)
8. ELECTROPHORESIS
Separates based on the charge of the particles
Remember: like charges repel; unlike charges
attract
Must have an electrical field with positive and
negative electrodes in order for electrophoresis to
work.
9. CENTRIFUGATION
Uses a device called a centrifuge to settle and
separate sediments
Separates based on the different densities of the
particles in the mixture.
How would YOU separate the
following mixtures?
a. flour and water
b. sugar solution and sand
c. 70% ethanol/water solution
d. oil, water and sand
e. Mercury and water (Hg is a really heavy
metal and is a liquid)
f. chlorophyll pigments
g. sugar and Kool-Aid