Honors Chemistry IA
Unit 1
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Atoms are the submicroscopic particles that
make up the basic building blocks of matter
“Smallest unit of matter”
These come together to form
molecules (covalent)
and compounds (ionic)
One carbon atom
for each oxygen atom
make up the molecule
carbon monoxide
Two hydrogen
atoms for each
oxygen atom make
up water
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Studying these atoms and how they arrange
is of interest to chemists
“Chemistry” – the science that seeks to
understand the behavior of matter by
studying the behavior of atoms and
molecules
◦ Focusses on matter and the changes they undergo
◦ Energy and matter conservation
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Scientists observe and perform experiments
on the physical world to learn about it
The Scientific Method is a series of steps used
to organize and test hypotheses, collect data,
and formulate conclusions
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Observations often lead scientists to
formulate a hypothesis
◦ Hypothesis is an interpretation or explanation of an
observation
◦ MUST be written in “if/then” form and MUST BE
TESTABLE!!!!
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We then test, or experiment, these
hypotheses to verify if we are correct or if we
need to go back 
Some conclusions may be a Scientific Law or a
Theory.
What is the difference ??
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A Law summarizes past observations and
predicts future ones.
◦ i.e. the Law of Conservation of Mass
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A theory a proposed explanation for
observations based on well-established and
tested hypotheses.
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Collecting observations is a
critical part throughout each
step
You observe to hypothesize
Experiment and then observe
Observe and then analyze
Observe and then form a
conclusion
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You go out in the morning before school in
December and your car wont start. Use the
scientific method to figure out a possible
solution.
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Matter is anything that has mass and takes up
space… in other words: anything with mass
and volume
Matter can exist in three states (or phases)
◦ Solid – atoms are tightly packed together
◦ Liquid – not as tight; able to slide past one another
◦ Gas – very loose; bouncing all over; no definite
shape or volume compressible
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Solid matter may also
exhibit a crystalline
structure.
◦ This is a long-range,
repeating order such as
diamond
◦ Very STRONG and
STABLE
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Liquids are not compressible and are packed
nearly as tightly as solids
They are able to move freely past one another
in a fluid motion
◦ This enables them to be “poured” and explains the
large range of motion of these particles
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Atoms have A LOT of space between molecules
/ atoms
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They are free to move in three dimensions past
and around one another
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They are COMPRESSIBLE!!
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If you are a pure substance, you can either be
a pure elemental or a pure compound
◦ Elemental – consisting of only one type of atom
◦ Compound – composed of two or more elements
(such as water and carbon dioxide)
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Heterogeneous Mixtures:
◦ Composition varies throughout
◦ If you sample from one spot it may not be the same
as a sample from another
◦ Salad, Pizza, ...
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Homogeneous Mixtures:
◦ Same composition throughout; uniform
◦ Kool-Aid, Salt water, ...
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Separation
techniques target
physical properties
to isolate and
separate the
components back
out
Can be very easy or
a little more
elaborate
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Changes that alter only the state or the
appearance but do not change the chemical
composition are physical changes
A Physical Property is one that a substance
displays without changing its composition
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A Chemical Change is a change that alters the
composition or matter
During a chemical change, atoms rearrange and
transform a starting substance into a new
substance
◦ “Bonds are broken, reformed, and gives you something
new”
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A chemical property is one that a substance
displays only by changing its composition via a
chemical change
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Determine whether each of the following
changes is physical or chemical
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The
The
The
The
evaporation of rubbing alcohol
burning of lamp oil
bleaching of hair with hydrogen peroxide
forming of frost on a cold night
◦ A copper wire hammered flat
◦ A nickel dissolves in acid to form a blue-green
solution
◦ Dry ice vaporizes without melting
◦ A match ignites when struck on a flint
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Energy exchange is necessary for a chemical
or physical change to take place
What is energy??
Energy is the “capacity to do work”
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What are two types of energy??
Kinetic and Potential
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Kinetic Energy is the
total energy
associated with its
motion (energy from
motion)
Potential is energy
from rest… “it has
potential – though not
moving yet”
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Thermal Energy is the energy associated with
the temperature of an object
It may got hot or cold… both exhibit a
change in temperatures
Exothermic and Endothermic
(review from bio IB)
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The energy (and mass) put into a system
MUST be recovered back out of the system in
some way shape or form
“Energy (and mass) is neither created or
destroyed”
The Law of Conservation of
Energy (and Mass)
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Systems with high potential energy will
always have the tendency to change in a way
that lowers their potential energy
It “dissipates” out and is absorbed by
surrounding bodies or the atmosphere
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In chemistry UNITS are critical
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Units – the standard quantities used to
specify measurements
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Gives a number meaning, without units they
are nothing
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We also need units that AGREE with one
another regardless of who or where in the
world we are working
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Two main types of measurement:
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English System (The American System) – used
in the U.S.
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The Metric System – used in most other parts
of the world
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Scientists all around the world use the Metric
System a.k.a. the International System of
Units (SI)
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Scientists use Celsius or Kelvin when
measuring temperature
There is nothing “Easy” or “clean” about the
Fahrenheit Scale (not SI units)
When given anything in F, you must first
convert to C or K
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Convert:
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212℃  ?? ℉
47 ℉  ?? ℃
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185 ℃  ?? ℉
275 ℃  ?? ℉
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76 ℉  ?? ℃
123 ℃  ??
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-22 ℉  ?? ℃
-17.1 ℃  ?? K
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4 ℉  ?? K
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The Metric System (SI) is a “base 10” scale
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Meaning, conversions are as simple as
moving the decimal over
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Prefixes are used as multipliers to denote
values
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Ex: kilo- means 103
(1,000)
milli- means 10-3
(0.001)
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Derived units can be made by combining
other units together.
Usually, these units are a measurement “per”
another (such as meters “per” second, or
grams “per” mole)
These units will tell you the mathematical
derivation of the value
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Density is defined as the amount of mass in a
given space (the mass “per” volume)
The unit to represent this is g/mL or g/cm3
As the unit indicates, the mathematical
equation for density is:
𝑑=
𝑚
𝑉
𝐷𝑒𝑛𝑠𝑖𝑡𝑦 =
𝑚𝑎𝑠𝑠
𝑉𝑜𝑙𝑢𝑚𝑒
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Density is an example of an intensive
property
◦ A property that is independent of the amount of the
substance
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Mass, in contrast, is an example of an
extensive property
◦ A property that is dependent (or depends on) the
amount of the substance
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Calculate the density of a sample with a mass
of 4.53 grams and a volume of 0.212 mL
(0.212 cm3)
A metal cube has an edge length of 11.4 mm
and a mass of 6.67 g. Calculate the density
of the metal  use your table on page 20 to
determine the identity of this unknown.
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A man receives a platinum ring from his
fiancé. Before the wedding, he notices that
the ring feels a little light for its size and
decides to measure its density. He places the
ring on a balance and finds that it has a mass
of 3.15 grams. He then find that the ring
displaces 0.233 cm3 of water. Is the ting
made of platinum (Pt)? Or is it a fake???
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Which data set seems to be more certain and
reliable?
Year
Carbon Monoxide
Concentration
(ppm)
Year
Carbon Monoxide
Concentration
(ppm)
1997
15.0
1997
15
1998
11.5
1998
12
1999
11.1
1999
11
2000
9.9
2000
10
2001
7.2
2001
7
2002
6.5
2002
7
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Scientific measurements are reported so that
every digit is certain except the last, which is
always estimated!!
So, that means you measure out as far as you
know for sure!! And thennnn estimate one
more digit.
◦ If it right between two lines you may estimate it to
be 0.5 and so on… the last one is not incorrect but
an estimate
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The non-place-holding digits (those that are
not simply marking the decimal place) are
called significant digits or significant figures
The greater the number of significant figures,
the greater the certainty of the measurement
23.45
23.5
24
certain
less certain
least certain
1.
2.
3.
4.
5.
All nonzero numbers are significant (1, 2, ..)
Sandwiched zeroes are significant (between
two nonzero numbers) (8,008 & 9,000,001)
Leading zeroes (to the left of a nonzero) are
not significant (0.00323 & 0.00006)
Trailing zeroes after a decimal point are
always significant (12.00 & 1.000x104)
Trailing zeroes with no decimal are not
significant (1200 & 145,000)
careful tho… 1200. makes them
significant
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Exact numbers are always significant, regardless of
zeroes
Counted values, conversion factors, constants are exact
◦ “I have 600 skittles in my pocket… not 597 rounded up… this is
an exact counted number
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Calculators DO NOT present values in the proper
number of sigfigs!
Exact Values have unlimited sigfigs
How many sigfigs do the following values
have?
46.3 lbs
40.7 in.
580 mi
87,009 km
0.009587 m
580. cm
0.0009 kg
85.00 L
580.0 cm
9.070000 cm
400. L
580.000 cm
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Multiplying / Dividing
The answer cannot have more sigfigs
than the value with the smallest number
of original sigfigs
ex:
12.548 x 1.28 = 16.06144
This value only has 3 sigfis,
therefore the final answer must
ONLY have 3 sigfigs!
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Multiplying / Dividing
The answer cannot have more sigfigs
than the value with the smallest number
of original sigfigs
ex:
12.548 x 1.28 = 16.06144
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This value only has 3 sigfis,
therefore the final answer must
ONLY have 3 sigfigs!
=16.1
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How many sigfigs with the following
FINAL answers have? Do not calculate.
12.85 * 0.00125
4,005 * 4000
48.12 / 11.2
4000. / 4000.0
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Adding / Subtracting
The result can be NO MORE certain than
the least certain number in the
calculation (total number)
ex:
12.4
Line up the decimal
points FIRST, then round
18.387
and chop off
+ 254.0248
284.8118
The least certain number is only certain to the “tenths”
place. Therefore, the final answer can only go out one
past the decimal.
ex:
12.4
18.387
+ 254.0248
284.8118
Least certain number (total
number)
=284.7
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Both addition / subtraction and
multiplication / division
Round using the rules after each
operation.
Ex: (12.8 + 10.148) * 2.2 =
22.9 * 2.2 = 50.38 = 50.
Scientific Notation – a number written a
the product of two values:
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A number out front & A x10 to a power
This notation allows us to easily work
with very, very large numbers or very,
very small numbers.
The number out front MUST be written
with ONLY one value prior to the decimal
point
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Examples:
a. 3.24x104g= 32,400 grams
b. 2.5x107mL = 250,000,000 mL
The exponent (x104) value can have a power
that is positive or negative, depending on if
you are dealing with a SMALL number or a
LARGE number
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Examples:
a. 8.55x104g
= 85,500 grams
b. 4.67x10-4 L
= 0.000467 Liters
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Addition / Subtraction
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6.2 x 104 + 7.2 x 103
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Addition / Subtraction
6.2 x 104 + 7.2 x 103
First, make exponents the same
62 x 103 + 7.2 x 103
Do the math and put back in Scientific
Notation
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Multiplication / Division
3.1 x 103 * 5.01 x 104
The “mantissas” are multiplied and the
exponents are added.
(3.1 * 5.01) x 103+4
16 x 107 = 1.6 x 108
Do the math and put back in Scientific
Notation (with correct number of sigfigs)
 Accuracy
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Vs. Precision
Measuring and obtaining data
experimentally always comes with some
degree of error.
Human or method errors & limits of the
instruments
We want BOTH accuracy AND precision
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Selecting the right piece of equipment is
key
Beaker, Graduated Cylinder, Buret?
Measuring 1.5 grams with a balance that
only reads to the nearest whole gram
would introduce a very large error.
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So what is Accuracy?
Accuracy of a measurement is how close
the measurement is to the TRUE value
“bull’s-eye”
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An experiment calls for 36.4 mL to be
added
Trial 1: delivers 36.1 mL
Trial 2: delivers 36.6 mL
Which is more accurate???
Trial 2 is closer to the actual value
(bull’s-eye), therefore it is more accurate
that the first delivery
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Now, what about Precision??
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Precision is the exactness of a measurement.
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It refers to how closely several measurements
of the same quantity made in the same way
agree with one another.
“grouping”
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Maximizing Accuracy and Precision will help
to Minimize ERROR
Error is a measure of all possible “mistakes”
or imperfections in our lab data
As we discussed, they can be caused from us
(human error), faulty instruments
(instrumental error), or from simply selecting
the wrong piece of equipment (methodical
error)
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Error can be calculated using an “Accepted
Value” and comparing it to the “Experimental
Value”
The Accepted Value is the correct value
based on reliable resources (research,
textbooks, peers, internet)
The Experimental Value is the value YOU
measure in lab. It is not always going to
match the Accepted value… Why not??
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Error is measured as a percent, just as your
grades on a test.
Percent Error = accepted – experimental
accepted
x100%
This can be remembered as the “BLT” equation:
•
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bigger minus littler over the true value 
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See “Dimensional Analysis” interactive slide
show