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Matter and Measurements
Honors Chemistry IAUnit 1
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)
Atoms and Molecules
One carbon atomfor each oxygen atommake up the moleculecarbon monoxide
Two hydrogen atoms for each
oxygen atom make up water
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
Chemistry
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
The Scientific Method
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!!!!
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 ??
A Law summarizes past observations and predicts future ones. ◦ i.e. the Law of Conservation of Mass
A theory a proposed explanation for observations based on well-established and tested hypotheses.
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
The Scientific Method
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.
Practice
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
Classification of Matter
Solid matter may also exhibit a crystalline structure.
◦ This is a long-range, repeating order such as diamond
◦ Very STRONG and STABLE
Solids
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
Liquids
Atoms have A LOT of space between molecules / atoms
They are free to move in three dimensions past and around one another
They are COMPRESSIBLE!!
Gases
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)
Pure Substances
Heterogeneous Mixtures:
◦ Composition varies throughout◦ If you sample from one spot it may not be the
same as a sample from another◦ Salad, Pizza, ...
Homogeneous Mixtures:
◦ Same composition throughout; uniform◦ Kool-Aid, Salt water, ...
Mixtures
Separation techniques target physical properties to isolate and separate the components back out
Can be very easy or a little more elaborate
Separating Mixtures
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
Physical Changes and Properties
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”
A chemical property is one that a substance displays only by changing its composition via a chemical change
Chemical Properties and Changes
Chemical Properties
Determine whether each of the following changes is physical or chemical◦ The evaporation of rubbing alcohol◦ The burning of lamp oil◦ The bleaching of hair with hydrogen peroxide◦ The 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
Practice
Energy exchange is necessary for a chemical or physical change to take place
What is energy??
Energy is the “capacity to do work”
What are two types of energy??
Kinetic and Potential
Energy
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”
Kinetic vs. Potential
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)
Thermal Energy
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)
Principle or Energy #1
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
Principle or Energy #2
In chemistry UNITS are critical
Units – the standard quantities used to specify measurements
Gives a number meaning, without units they are nothing
We also need units that AGREE with one another regardless of who or where in the world we are working
Units of Measurement
Two main types of measurement:
English System (The American System) – used in the U.S.
The Metric System – used in most other parts of the world
Scientists all around the world use the Metric System a.k.a. the International System of Units (SI)
Units of Measurement
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
What Units do we want??
Convert:
212℃ ?? ℉ 47 ℉ ?? ℃
185 ℃ ?? ℉ 275 ℃ ?? ℉
76 ℉ ?? ℃ 123 ℃ ??
-22 ℉ ?? ℃ -17.1 ℃ ?? K
4 ℉ ?? K
The Metric System (SI) is a “base 10” scale
Meaning, conversions are as simple as moving the decimal over
Prefixes are used as multipliers to denote values
Ex: kilo- means 103 milli- means 10-3 (1,000) (0.001)
Metrics Made Easy
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
Derived Units
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:
Density: A derived unit
Density is an example of an intensive property ◦ A property that is independent of the amount of
the substance
Mass, in contrast, is an example of an extensive property◦ A property that is dependent (or depends on) the
amount of the substance
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.
Practice with Calculations
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???
Which data set seems to be more certain and reliable?
Reliability and SigFigs
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
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
Read each to the correct number of SigFigs
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 certain 23.5 less certain 24 least certain
Counting SigFigs
1. All nonzero numbers are significant (1, 2, ..)2. Sandwiched zeroes are significant (between
two nonzero numbers) (8,008 & 9,000,001)3. Leading zeroes (to the left of a nonzero) are
not significant (0.00323 & 0.00006)4. Trailing zeroes after a decimal point are
always significant (12.00 & 1.000x104)5. Trailing zeroes with no decimal are not
significant (1200 & 145,000)careful tho… 1200. makes them significant
Rules
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
Calculators DO NOT present values in the proper number of sigfigs!
Exact Values have unlimited sigfigs
Exceptions
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
Practice
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
Calculating with SigFigs
This value only has 3 sigfis, therefore the final answer must ONLY have 3 sigfigs!
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
=16.1
Calculating with SigFigs
This value only has 3 sigfis, therefore the final answer must ONLY have 3 sigfigs!
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
Practice
Adding / Subtracting The result can be NO MORE certain
than the least certain number in the calculation (total number)
ex: 12.4 18.387
+ 254.0248 284.8118
Calculating with SigFigs
The least certain number is only certain to the “tenths” place. Therefore, the final answer can only go out one past the decimal.
Line up the decimal points FIRST, then round and chop off
ex: 12.4 18.387
+ 254.0248 284.8118 =284.7
Calculating with SigFigs
Least certain number (total number)
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.
Calculating with SigFigs
• Scientific Notation – a number written a the product of two values:• 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.
Scientific Notation
• The number out front MUST be written with ONLY one value prior to the decimal point
Examples: a. 3.24x104g= 32,400 gramsb. 2.5x107mL = 250,000,000 mL
Scientific Notation
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
Examples: a. 8.55x104g b. 4.67x10-4 L = 85,500 grams = 0.000467 Liters
Scientific Notation
Addition / Subtraction
6.2 x 104 + 7.2 x 103
Scientific Notation
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
Scientific Notation
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)
Scientific Notation
Accuracy 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
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.
Experimental Error
So what is Accuracy?
Accuracy of a measurement is how close the measurement is to the TRUE value
“bull’s-eye”
Accuracy
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
Accuracy
Now, what about Precision??
Precision is the exactness of a measurement.
It refers to how closely several measurements of the same quantity made in the same way agree with one another.
“grouping”
Precision
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)
Error
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??
Error
Error is measured as a percent, just as your grades on a test.
Percent Error = accepted – experimentalx100% accepted
• This can be remembered as the “BLT” equation:
bigger minus littler over the true value
Error
See “Dimensional Analysis” interactive slide show
Conversions
