Chapter 1. Introduction: Matter and Measurement Watch Bozeman Videos & other videos on my website for additional help: Big Idea 1: Molecules/Elements Chemical Analysis Mass Spectrometry Conservation of Atoms 1.1 The Study of Chemistry • Chemistry: • is the study matter and changes that they undergo. Why Study Chemistry? • Considerable impact on society (health care, food, clothing, conservation of natural resources, environmental issues etc.). • Chemistry serves biology, engineering, agriculture, geology, physics, etc.. Chemistry is the central science. 1.2 Classifications of Matter Matter: has mass and occupies space. 3 States of Matter Solids Have a definite shape and definite volume. The particles in a solid are packed tightly together and only vibrate gently around fixed positions. Incompressible. Liquids Has no shape of their own. Takes shape of container. A liquid has a definite volume. The particles in a liquid are free to move. No collusions – Slide past each other. Incompressible. Gases Have no shape or volume. The particles in a gas spread apart filling all the space of the container available to them. Move at very high speeds. Collusions Compressible.
Transcript
Chapter 1. Introduction: Matter and Measurement Watch Bozeman Videos & other videos on my website for additional help: Big Idea 1:
Molecules/Elements
Chemical Analysis
Mass Spectrometry
Conservation of Atoms
1.1 The Study of Chemistry • Chemistry: • is the study matter and changes that they undergo.
Why Study Chemistry?
• Considerable impact on society (health care, food, clothing, conservation of natural resources, environmental issues etc.).
• Chemistry serves biology, engineering, agriculture, geology, physics, etc.. Chemistry is the central science.
1.2 Classifications of Matter
Matter: has mass and occupies space.
3 States of Matter
Solids Have a definite shape and
definite volume.
The particles in a solid are
packed tightly together and
only vibrate gently around
fixed positions.
Incompressible.
Liquids
Has no shape of their own.
Takes shape of container.
A liquid has a definite
volume.
The particles in a liquid are
free to move.
No collusions – Slide past
each other.
Incompressible.
Gases Have no shape or volume.
The particles in a gas spread
apart filling all the space of
the container available to
them.
Move at very high speeds.
Collusions
Compressible.
Matter Chart
0
10
20
30
40
50
60
70
80
90
1st Qtr 2nd Qtr 3rd Qtr 4th Qtr
East
West
North
2 Types of Matter: Pure Substances and Mixtures
Pure Substances: (2 types)
1. Elements: Building Blocks of Matter • Cannot be decomposed into simpler substances - only one kind of atom • There are 116 known elements.
• Names derived from a wide variety of sources (e.g., Latin or Greek, mythological characters, names of people or places).
• Each is given a one- or two-letter symbol derived from its name.
2. Compounds • Consist of two or more different elements - Metals and Nonmetals (NaCl) • Compounds have different properties than their component elements (e.g., water is liquid, but
hydrogen and oxygen are both gases at STP).
• Law of Constant (Definite) Proportions (Proust): A compound always consists of the same combination of elements (e.g., water is always 11% H and 89% O).
Molecules: • are combinations of atoms held together in specific shapes. • ONLY NONMETALS!
Mixtures: (2 types) A mixture has varying composition….can be physically separated.
1. Homogeneous. Uniform in composition throughout a given sample but the composition
and properties may vary from one sample to another; e.g. a solution of salt water.
ALL HOMOGENEOUS MIXTURES ARE CALLED SOLUTIONS!!!
2. Heterogeneous. Have separate, distinct regions within the sample. As a result the
composition and properties vary from one part of the mixture to another; e.g. a chocolate
chip cookie.
1.3 Properties of Matter
Property: allows us to recognize a particular type of matter and to distinguish it from other types of matter.
Physical and Chemical Properties and Changes All matter exhibits physical and chemical properties by which it can be classified.
Introduction: Matter and Measurement 3
Physical properties: color, odor, density, hardness, solubility, melting point, and boiling point.
Chemical properties: when a substance reacts with other substances.
Examples of chemical properties are reactions with acids and bases, oxidation and reduction and
a huge number of other chemical reactions.
Each substance has a unique set of physical and chemical properties.
• Intensive properties do not depend on the amount of substance present (e.g., temperature, melting point etc.).
• Extensive properties depend on the quantity of substance present (e.g., mass, volume etc.). Physical Change: If some aspect of the physical state of matter is altered, but the chemical composition
remains the same, the change is a physical change. The most common physical changes are changes of
state. (example: liquid water to vapor to ice – It’s still water)
Chemical Change: In a chemical change, which is often called a chemical reaction, the atoms of a
substance are rearranged to form new substances. A chemical change requires that the new substance or
substances formed have a different chemical composition to the original substance or
substances. Chemical changes are often accompanied by color changes and/or heat changes.
Separation of Mixtures
• Key: separation techniques exploit differences in properties of the components. • Filtration: remove solid from liquid. • Distillation: boil off one or more components of the mixture.
• Chromatography: exploit solubility of components.
The Scientific Method
• The scientific method provides guidelines for the practice of science. • Eventually after several experiments the hypothesis may become a theory. A theory
gives a universally accepted explanation of the problem.
Chapter 1 4
• Theories are different to laws. Laws state what general behavior is observed to occur naturally.
E.g. The Law of Conservation of Mass exists since it has been consistently observed that during
all chemical changes mass remains unchanged (i.e. it is neither created nor destroyed).
1.4 Units of Measurement
• Many properties of matter are quantitative (numbers). • A measured quantity must have BOTH a number and a unit. • The units most often used for scientific measurement are those of the metric system.
SI Units
• 1960: All scientific units use Système International d’Unités (SI Units).
• THESE ARE CALLED BASE UNITS!!!
Base Quantity Name of Unit Symbol
Mass Kilogram kg
Length Meter m
Time Seconds s
Amount of Substance Mole mol
Temperature Kelvin K
Volume Liter L
Prefix Symbol Meaning
Metric Conversion
Tera = 1012
T
Giga = 109 G
Mega = 106 M
kilo = 103 k
hecto = 102 h
deka = 101 da
centi = 10-2
c
milli = 10-3
m
micro = 10-6
nano = 10-9
n
pico = 10-12
p
femto = 10-15
f
Introduction: Matter and Measurement 5
Prefixes Base units Prefixes
Remember: G __ __ M __ __ k h da m d c m __ __ µ __ __ n
L
g
s
***Giant Mighty king henry drank milk during christmas morning until noon***
Note: EVERY UNIT HAS A BASEUNIT…BUT NOT EVERY UNIT HAS A PREFIX!!
Mass vs. Weight – chemists are quite guilty of using these terms interchangeably.
mass (g or kg) – a measure of the resistance of an object to a change in its state of motion; the
quantity of matter present.
weight (a force ∴Newtons) – the response of mass to gravity; since all of our measurements
will be made here on Earth. We “weigh” chemical quantities on a balance NOT a scale!!
Converting Units - One unit can be converted to another by using a conversion factor.
9.00 inch x 2.54 cm = 22.9cm
1 inch
Temperature
Temperature is the measure of the hotness or coldness of an object.
3 scales: Celsius (oC), Fahrenheit (
oF) and Kelvin (K).
Scientific studies use Celsius and Kelvin scales. • Celsius scale: water freezes at 0
oC and boils at 100
oC .
• Kelvin scale: (SI Unit) water freezes at 273.15 K and boils at 373.15 K .
• 0 K = –273.15oC. ABSOLUTE ZERO(lowest possible temp)
• Fahrenheit scale: water freezes at 32oF and boils at 212
oF .
Chapter 1 6
Temperature Conversion factors
Convert the following temperatures from one unit to the other.
(i) 1390 oC to K
(ii) 38 K to oF
(iii) 13 oF to
oC
Derived SI Units:
• These are formed from the seven base units. (example is m/s)
Volume
• Units of volume = (units of length)3 = m
3.
• cm3 [also known as cc (cubic centimeters)]
1cm3 (solid) = 1mL (liquid)
1dm3 = 1L (gases)
Density
• Defined as mass divided by volume.
• Units: g/cm3 (solids); g/mL (liquids); g/L (often used for gases).
Density = mass
volume
1.5 Uncertainty in Measurement When reading the scale on a piece of laboratory equipment such as a graduated cylinder, there is always a
degree of uncertainty in the recorded measurement. The reading will often fall between two divisions on
273.15CK
273.15KC
32)F(9
5C
32C5
9F
Introduction: Matter and Measurement 7
the scale and an estimate must be made in order to record the final digit. This estimated final digit is said
to be uncertain and is reflected in the recording of the numbers by using +/-.
The certain and the uncertain numbers taken together are called significant figures.
Rules of significant figures
1. Any non-zero integers are always counted as significant figures.
2. Leading zeros are those that precede all of the non-zero digits and are never counted as
significant figures. (Example: 0.0003 has one significant figure.) 3. Captive zeros are those that fall between non-zero digits and are always significant figures.
(Example: 100405 has six significant figures.)
4. Trailing zeros are those at the end of a number and are only significant if the number is
written with a decimal point. (Example: 1000.0 has five significant figures.) 5. In scientific notation the 10x part of the number is never counted as significant.
Determine the number of significant figures in the following numbers.
(i) 250.7
(ii) 0.00077
(iii) 1024
(iv) 4.7 x 10-5
(v) 34000000
Significant Figures in Calculations
• Multiplication and division: • Report to the least number of significant figures
4.56 × 1.4 = 6.38 6.4 corrected • Addition and subtraction: • Report to the least number of decimal places 12.11
18.0 ← limiting term
1.013 31.123 31.1 corrected
Using a calculator carry out the following calculations and record the answer to the
correct number of significant figures.
(i) 34.5 x 23.46
(ii) 123/3
(iii) 2.61 x 10-1 x 356
(iv) 21.78 + 45.86
(v) 23.888897 - 11.2
(vi) 6 - 3.0
Rounding off
Calculators will often present answers to calculations with many more figures than the significant
ones. As a result many of the figures shown are meaningless, and the answer, before it is
presented, needs to be rounded off.
In a multi-step calculation, leave the rounding until the end!
Chapter 1 8
Leave all numbers on the calculator in the intermediate steps, then round to the correct number of
significant figures at the end of the calculation.
Look at the significant figure one place beyond your desired number of significant figures….
If greater then 5 = round up
If less then 5 = leave alone
Precision and Accuracy
•Precision: refers to how close two or more measurements of the same quantity are to one another.
•Accuracy: relates to how close the measured value is to the actual “true” value.
Consider three sets of data that have been recorded after measuring a piece of wood
that is exactly 6.000 m long.
SET X SET Y SET Z
5.864m 6.002m 5.872m
5.878m 6.004m 5.868m
Average 5.871m 6.003m 5.870m
(i) Which set of data is the most accurate?
(ii) Which set of data is the most precise?
Percentage Error
The data that is derived from experiments will often differ from the actual value.
You want a very low percentage!!!! Less than 15% is pretty good lab results.
Percent error = l accepted value – experimental value l
