Transcript Chem 121 - Pierce College

Chem 121
Introduction to Inorganic Chemistry
What is Matter?
• Matter is anything that has mass and occupies
space.
• Mass is a measurement of the amount of matter
present. The mass is constant, no matter where
it is …on the moon or on earth.
• Weight the a measurement of the gravitational
force pulling the object toward earth. The weight
of an object will be different on the moon that
has a different grav. pull.
Properties and Changes
• Two types of properties:
– Physical : Those that can be observed or measured
without changing or trying to change the
composition of the matter in question * no
original substances are destroyed and no new
substances are created. Like Color and SIZE.
– Chemical: properties that matter demonstrates
when attmepts are made to change it into other
kinds of matter.
Physical Change VS Chemical Change
• Physical Change: Cutting a piece of paper,
Boiling water, freezing water, heat is added or
removed from matter.
• Chemical Change: burn the paper. ,mix the
water with an acid.
A Model of Matter
• Scientific Models are explanations for observed
behavior.
• All matter is made up of particles that are too
small to see. (molecules)
• Molecules: The smallest particle of a pure
substance that has the properties of that
substance and is capable of a stable independent
existence. A molecule is also the limit of physical
subdivision for a pure substance.
Example
• Oxygen: helps a substance burn more
rapidly..like wood.
– A large amt. or a small amt of Oxygen would
behave the same.
– The smallest amt. that would still behave the
same is known as the molecule.
What’s beyond the Molecule?
• John Dalton wanted to know. In 1808 he
proposed the following:
• 1. All matter is made up of tiny particles called
atoms.
• 2. Substances called elemenats are made up of
atoms that are all identical.
• 3. Substances called compounds are
comginations of atoms or two or more elements.
Cont.
• 4. Every molecule of a specific compound
always contains the same number of atoms of
each kind of element found in the compound.
• 5 In chemical reactions, atoms are rearranged,
separated, or combined, but are never created
nor destroyed.
Types of Molecules
• Diatomic Molecule: contains two atoms
• Homoatomic molecule: contain only one kind
of atom.
• Heteratomic molecule: contain more then one
kind of atom
• Triatomic molecule: 3 non-identical atoms
• Polyatomic molecule: more then 3
Classification of Matter
• Pure Substances - not adulterated or mixed with
anything else.
– They have unique and consistent physical and chemical
properties.
• Physical Properties: Melting Point Temperature, Color, Density.
• Chemical Properties: a chemicals ability to react with other pure
substances. In a chemical reaction, substances lose their identity
and form new substances with new chemical and physical
properties.
– It undergoes physical change without losing its identity...
Eg.( melting, freezing, or evaporation)
Mixtures
• Consists of two or more pure substances in varying
proportions.
– Heterogeneous – visibly discontinuous..like salt and
pepper.
– Homogeneous – have a uniform appearance throughout;
like sugar and water. The mixture is called a solution, and
it is described as homogeneous.
• Mixtures can be separated back into their pure
substance components.
• Mixtures have properties that are variable and depend
on the proportions of the components.
Compounds and Elements
• Compounds: Some pure substances are found
to be able to be decomposed into simpler
pure substances.
• Elements: pure substances that cannot be
further decomposed. It cannot be separated
chemically in to simpler substances, nor be
created by combining simpler substances.
Measurement and the Metric System
• Measure: - the size, capacity, extent, volume
or quantity of anything, especially as
determined by comparison with some
standard or UNIT.
• Systeme International d’Unite’s.
Significant Figures
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Communicating Degrees of Uncertainty
4 1 sig fig
4.0 2 sig figs
4.000 4 sig figs
4.0000 5 sig figs
Examples
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Do not over represent the amt. of precision that you have.
Which digits are really giving me info about how precise my measurement is?
0.00700
If you measure the above in km, it could also be 7.00 m (the previous zeros are
determining the units to use, the trailing zeros determine precision)
0.052 …(could be re-written as 52 m)
(do not count leading zero’s before the first non-zero digit)
370. (because they wrote a decimal, it is exactly 370) 3 sig figs.
10.0 (go to nearest 10th) 3 sig figs
705.001 ( zero’s are part of measurement ..between non zero digits)
37,000 (ambiguous) Maybe you measured to the nearest 1000, or nearest 1…you
don’t know. Go with 2 sig figs. (more conservative)
( A trailing Zero as in 4.130 is significant. (This has 4 sig figs)…
Rules of Thumb
• A trailing zero , 4.130 , is significant.
• A zero within a number, 35.06 cm
• A zero before a digit as in 0.082 , is not
significant
• A number ending in zero with no decimal
point , 20 is ambiguous.
Mulitiplication and Division with Sig
Figs
• Let’s say we are calculating the area of a
Rectangle 1.69 m x 2.09m
• Area = 3.5321 m2
Use the least precisice number as the basis for
the amt. of sig figs.. 3 sig figs..
• Area = 3.53 m2
Another Example
• Calculate how many tiles I need for a room 12.07
ft x 10.1 ft.
• Floor Area = 121.907 ft2
• ** Do not round yet! *** go to the end with all
numbers, then establish sig figs and round**
• Tiles in bathroom = 121.907ft2/1.07 ft2
• Tiles = 113.931775701 tiles
• 3 sig figs = 114 tiles
Addition and Subtraction
• Ex: 1.26 (nearest hundredth , 3 sig figs) + 2.3 (nearest
10th, 2 sig figs) = 3.56
• The least precise number went to the 10th, therefore 2
sig figs in result. 3.7
• Or 1.26 + 102.3 = 103.56 (only as precise as the least
precise number)…
• 1.26 has 3 sig figs, 102.3 has 4 sig figs, however the
least precise measurement is 102.3 as it is measured
only to the 10th, not the 100th…therefore the answer
will be 4 sig figs, 103.6
Another Example
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One Block: 2.09 m high
Another block: 1.901 m high
How tall is it to stack them?
1.901 + 2.09 = 3.991
Did I measure the entire stack to the nearst
mm? NO! Only report as precise as the least
precise measurement. 3.99m
Another example
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Building: 350 ft tall (ambiguous)
Radio Tower : 8 ft tall
How tall is building plus the tower: 358 ft
We only measured tower to nearest ft.
You have to round to the nearest 10 ft.
Answer is actually 360 ft. or report it as 3.6 x
10 2
Metric System Units
Using Units in Calculations
aka: Dimensional Analysis
• Step 1: Write Down the known or given quantity.
Include both numerical value and units of the
quantity.
• Step 2: Leave some working space and set the
known quantity equal to the units of the
unknown quantity.
• Step 3: Multiply the known by one or more
factors (conversion factors) to cancel the units of
the known and generate the units of the
unknown.
• Step 4: Do the arithmetic .
Example Problems
• 50 μL (50 – microliter) sample of blood serum
must be expressed as Liters.
• (1μL = 1 x 10-6 L)
• Step 1:
• Step 2:
• Step 3:
• Step 4:
Example
• One of the fastest-moving impulses in the
body travels at a speed of 400 ft/per second.
(ft/s). What is the speed in miles/hr?
Calculating Percentages
• Percent =
number of specific items x 100
total items in the group
% = Part/Total x 100
Non SI units most common
Prefixes for Metric Units
Mass, Volume, & Density
• Mass is the measure of a quantity of matter. It is
measured relative to a standard mass (which is why the
devices to weigh an object are called Balances) Mass is
not Weight.
• Volume is the amount of space a sample occupies. 1
mL = 1cm3
• Denisty – a physical property of a substance =
mass/volume. Since volume increases with
Temperature increase, a density is always reported
with a Temperature.
Example
• A 35.66 g sample of metal as weighed and put
into a graduated cylinder that contained 21.2
mL of water. The water level after the metal
was added was 25.2 mL. What is the density
of the metal in g/cm3
How to measure Density
• Take a substance and weight it. Then add it to a
known volume of water in a volumetric flask and
notice the volume change in the water as the
volume of the substance. Calculate the density.
• To calculate the mass of a liquid, you add the
liquid to a zeroed balance with a volumetric flask
and weight the liquid. Then you have the mass
and the volume, and you can caluclate the
density.
Density Example Problem
• What is the volume of a 32 g sample of
ethanol whose density is 0.789g/cm3 ?
Report volume in cm3
Hydrometer
• If something floats on water, it is less dense
then water.
• If something sinks, it is more dense then
water.
• A hydrometer rises or falls to a density that is
equal to the density of the liquid. It is
calibrated to show the specific gravity of the
liquid.
Specific Gravity
• S.G. = density of test liquid/ density of reference
liquid
• Note that the units cancel.
• The standard reference liquid for measuring the
specific gravity of aqueous solutions is pure water
at 4 deg. C. Density is 1.000g/cm3 .
• S.G. of blood is 1.028.
• This means that blood is 1.028 times the density
of pure water.
Temperature
• Substances can either gain heat or lose it,
depending on whether they are cooler or
hotter then their environments.
• To measure heat, we must have an indication
of how hot or cold something is…that is the
temperature.
• It indicates how hot something is…not the
amount of heat.
Kelvin, C, and F
• K = C + 273
• F = 9/5 C + 32
• C (F-32)(5/9)
Heat and Calorimetry
• Heat is a form of energy.
• Each substance has a different capacity to absorb
heat.
• A unit of heat is defined by its effect (the rise in
temperature) on a fixed mass of a reference
substance.
• SI unit = joule, (J)
• Non-SI commonly used = cal
• 4.184 J = 1 cal
Specific Heat : Cp
• The characteristic response of a given mass of a
given substance to a given amount of heat is
expressed by Cp
• Cp= joules/(grams x Δ°C)
• The specific heat is equal to the heat
absorbed or lost per Celsius degree change
in temperature per gram of substance.
Specific Heat
• The higher a substances specific heat, the more slowly it’s temperature
rises in repsonse to heating.
Calculating Specific Heat
• What is the specific heat of a substance if the
addition of 334 J of heat to 52 g of that
substance causes the temperature to rise from
16 C to 48 C?
Calculating Heat from Cp
• How much heat must be added to 45.0 g of a
substance that has a specific heat of 0.151
J/gC to cause it’s t to rise from 21 C to 47 C ?
Calorimeter
• When the heat produced by some physical or
chemical process is abosrbed into a given
mass of water, the water’s T rise will allow us
to calculate the heat produced by the process.
Example
• What is the Specific Heat (Cp) of an element
that takes 50 Joules to heat up 400 grams
from 34 deg to 76 deg.?
Calorimeter
Basal Metabolic rate
• BMR is the minimum metabolic activity of a
human at rest and with an empty gi tract.
• In nutrition and metabolism, heat is more
commonly given as calories.
• The rate means it is the amount of heat over a
period of time…expressed in kcal/min.