Transcript Ch. 1 Introduction: Matter and Measurement

Chemistry: A Molecular Approach
 The properties of matter are determined by the
properties of atoms and molecules.
 The understanding of matter at the molecular level
gives us the ability to control that matter.
 First semester General Chemistry – the primary goal is
to learn the fundamental principles of Chemistry.
Why are you here?
 Why study chemistry?
 Reasons
 1)
 2)
 3)
Ch. 1: Matter, Measurement…
 A comparison of two molecules:
 CO consists of one carbon and
one oxygen atom.
C
O
 CO2 consists of one carbon and
two oxygen atoms.
 One is a poison while the other is
harmless.
O
C
O
Ch. 1: Matter, Measurement…
 Matter can exist in three different states:
 Solid – Is rigid, has both a definite shape and volume. Solids
can by crystalline or amorphous.
 Liquid – Is a fluid – molecules or atoms can flow.
However, it has a fixed volume but no fixed shape.
 Gas – Also a fluid. However, it has no fixed volume or
shape. Gases are unique in that they can be compressed.
 Liquids and Solids cannot be compressed.
 Molecular View
Learning Check
 What phase(s) can be compressed?
 What phase(s) have a fixed volume?
 What phase(s) have a fixed shape?
Classification of Matter
 Pure substance – has a fixed composition and
distinct properties. A pure substance can be
either an element or a compound.
 Ex) Table sugar (sucrose) and Carbon are pure
substances.
Classification of Matter
 Elements are substances that cannot be decomposed
into simpler substances. Each element is composed of
only one kind of atom.
 Compounds are composed of two or more elements
that have been chemically combined.
 Ex) Sulfur and Sodium chloride
Classification of Matter
 Most matter we encounter is a mixture of
substances.
 The elements and / or compounds in a mixture
always retain their identity.
 Ex) Gatorade and Milk
Classification of Matter
 Homogenous – the mixture is
uniform throughout. Most
homogenous mixtures are
solutions.
 Crystal Light is a solution.
 Heterogeneous – the mixture is
not uniform throughout.
 A chocolate chip cookie is a
heterogeneous mixture.
Separation of Mixtures
Matter
Variable composition?
No.
Yes.
Pure Substance
Mixture
Separation by
________________?
____________?
No.
Yes.
No.
Yes.
Element
Compound
Heterogeneous
Homogeneous
Learning Check
Decide whether the following is a mixture or a
pure substance.
 Tomato juice
 Iodine crystals
 Sand
 Baking soda
 A substance does not have a variable composition
and can be separated by chemical means. This is a
_________________________.
Separating Mixtures
 Separating mixtures into their
components is something that
scientists often do.
 Simple: filtering sand from
water.
 Simple: using a magnetic to
collect iron fillings from sand.
 Hard: Distillation of two liquids.
 Hard: Chromatography.
Properties of Matter
 Physical property - a characteristic that can be
observed for a material without changing its chemical
identity.
 extensive property - dependent on the amount

ex) mass, volume, heat content
 intensive property - independent of the amount

ex) density, temperature, melting point
Physical Properties
 Copper
 Reddish, shiny
 Melting point = 1085oC
 Density = 8.94 g/mL
 Specific heat = 0.385 J/g C
 Crystal structure = face
centered cubic
Properties of Matter
 Chemical property - describes
how a substance may change
or react towards other
substances.

Ex) Propane burns in air to form
carbon dioxide and water
Physical and Chemical Changes
 Physical change - a change in the form of matter but
not it’s identity - can often easily return to the former
state



ex) melting of ice
ex) dissolving of salt in water
ex) ripping a piece of paper
Physical and Chemical Changes
 Chemical change - a change in which one or more
kinds of matter are transformed into new kinds of
matter - difficult to return to the former state (by any
physical means)


ex) rusting of iron
ex) burning of wood
Learning Check
 Would classify these common daily activities as a
physical or chemical change?
 Getting a haircut
 Applying bleach to turn your hair blond
 Baking a cake
 Adding a Crystal Light© packet to water
 Leaves on a tree turning red or yellow
 Rubbing alcohol evaporating from your skin
 Pouring spaghetti sauce over pasta
Energy
 Most physical and chemical
changes involve changes in
energy.
 Example – water evaporating
from your skin.
 Example – Burning propane in
an outdoor grill.
 Energy = capacity to do work.
 Work = Force times distance.
Types of Energy
 Potential Energy = energy of
position or composition.
 Kinetic Energy = energy of
motion.
 Law of conservation of
energy.
 Tendency of systems with
high PE.
Units of Measurement
 Units used in scientific measurements are those of the
metric system.
 Although we still use the English system, the metric
system is becoming more common.
SI units
 1960 agreement on a set of internationally accepted group
of seven base units from which all others are derived.
 Mass = kilogram
 Length = meter
 Time = second
 Temperature = Kelvin
 Amount of Substance = mole
 Electric current = Ampere
 Luminous intensity = Candela
Metric Prefixes
 Common ones include:
 Kilo (k) = 103
 Centi (c) = 10-2
 Milli (m) = 10-3
 Micro (m) = 10-6
 Nano (n) = 10-9
 Non standard unit, Angstrom’s (Å)
 1 Å = 1 x 10-8 cm
Temperature Scales
 Temperature is a
measurement of the hotness
or coldness of an object.
 English scale = Fahrenheit
 Celsius scale – assigns
temperatures based on
melting and boiling points of
water.
 Kelvin scale – based on
absolute zero being the
coldest possible temperature
 K = oC + 273.15
Derived Units
 Examples are units like volume, density, and velocity.
 Volume of a cube = length cubed
 Density of a substance = mass divided by volume.
Uncertainty in Measurement
 Exact numbers – a number that is known to be
precisely that value. These have no effect on sig.
figs.
 Ex) 12 inches = 1 foot
 Ex) 15 apples in a bag
 Inexact (measured) numbers – have some
amount of uncertainty.
 Ex) A coin has a mass of 2.52g
 Ex) A bottle of soda has a volume of 591mL
Uncertainty in Measurement
 Any measurement contains some uncertainty.
 precision - the closeness of a group of figures to each
other - standard deviation
 accuracy - the closeness of a single value or an average
to the accepted value - percent error
Density of Al
Student A
Student B
Student C
Trial 1
2.5 g/mL
2.58 g/mL
2.68 g/mL
Trial 2
3.3 g/mL
2.51 g/mL
2.72 g/mL
Trial 3
3.1 g/mL
2.55 g/mL
2.69 g/mL
Average
3.0 g/mL
2.55 g/mL
2.70 g/mL
%error / %rad
Study Check
 Do the following represent Exact or Measured
numbers?
 A store has 35 bicycles on display.
 The density of an object is found to be 1.8g/mL.
 1 meter is equivalent to 100 centimeters.
 Planck’s constant is listed in the book as 6.626 x 10-34J s.
 There are about 454 grams in one pound.
Significant Figures
 Significant figures are the digits measured in a number
such that all certain digits plus one uncertain digit is
included.
 Certain digits – all performing measurement would
agree on these.
 Uncertain digit – a “best guess” – it is each individual’s
best interpretation of the measurement.
“Best Guess” examples
Rules:
 Non-zero numbers are always significant.
 Zeros between non-zero numbers are always significant.
 Zeros at the beginning of a number are NEVER
significant; they are merely placeholders.
 Zeros that fall at the end of the number and after the
decimal point are always significant.
 When a number ends in zeros without a decimal point,
the zeros may or may not be significant. We will err on
the least number.
Scientific Notation
 Puts all numbers in the form of: A x 10n, where A is a
number between 1 and 10 and n is the exponent
equally to the number of places the decimal point
must be moved.
 Scientific Notation removes ALL ambiguity from
determining significant figures.
 See Appendix A for a review of scientific notation.
Examples
 How many significant figures do these measured
numbers have?
0.0092
0.00920
9.20
92,000
92,000.0
Sig. Figs. In Calculations
 Multiplication and Division - your answer will keep only
the same number of sig. figs. as the measurement that had
the fewest number of sig. figs.
 Addition and Subtraction - your answer will have the same
number of sig. figs. as the one with the fewest decimal
places.
 Note – in series of calculations, do not round until the very
end. In mixed calculations, follow rules for each individual
calculation.
Rounding of Numbers
 Rounding is the process of dropping nonsignificant digits in a calculation and adjusting the
last digit.
 Rules:
 if the leftmost digit (to be dropped) is 5 followed with
any non-zero digits, then round the final digit up one.
 if the leftmost digit is less than 5, round down.
 if the digit is 5 only, and the digit to be rounded is even,
round down. if it is odd, round up.
Examples
 Round each calculation to the correct number of
significant figures.
1.305 x 0.056 = 0.07308
105.2 x 0.00057 = 0.059964
495.0 ÷ 0.23 = 2152.173913
Examples
 Round each calculation to the correct number of
significant figures.
25.0 + 2.86 = 27.86
69.72 – 67.92 = 1.8
121 – 3.89 = 117.11
Examples
(10.397 – 10.147) x 10.00 = 2.5
(85.2 + 79.9) ÷ 180.5 = 0.91468144
(48.32 – 4.5) ÷ 85.72 = 0.51119925