Unit 1 - Introduction to Chemistry (2017/2018)

Introduction to
Chemistry
Unit 1

1.1 What is
Chemistry?
I will be able to…
• Define chemistry and
chemical.
• Explain how chemistry is an
interdisciplinary science.

1.1 What is Chemistry?
•Chemistry – the study of the
composition, structure, properties, and
change of matter.

Where is the CHEMISTRY?

Where is the CHEMISTREE?

•Chemical – any substance that has a
definite composition.
• It’s always made of the same stuff, no
matter where the chemical comes from.

•Interdisciplinary Study – a science that
involves many fields of study.

1.1 Assignments
•1.1 What is Chemistry? WS

1.2 Scientific
Method
• List and describe the steps of
the scientific method.
• Apply the steps of the scientific
method to a real world
example.
• Identify dependent and
independent variables.
• Identify experiment and control
groups.
• Identify qualitative and
quantitative observations.
• Define scientific law and
scientific theory.

1.2 Scientific Method
•Scientific Method – a series of steps
followed by scientists to solve problems.

•Step 1 – Ask Questions or Identify a Problem
• Who?
• What?
• Where?
• Why?
• How?

•Step 2 – Construct a Hypothesis
• Hypothesis – a testable explanation for
an observation.
• A hypothesis is an educated guess based
on prior knowledge, research and
observations.
• A hypothesis must be
• observable.
• measurable.

•Step 3 – Perform an Experiment
•Experiment – a procedure designed to
test a hypothesis under controlled
conditions.
• One variable is tested.
• One control is tested.

•Variable – a factor that could
affect the results of an
experiment.
•Independent Variable –
variable being changed.
•Dependent Variable –
variable being measured.

•Experiment Group – the group in the
experiment that receives treatment.
•Control Group – the group in the
experiment that does not receive
treatment.

• Independent Variable =
• Dependent Variable =
• Control Group (s) =
• Experimental Group(s) =

• Step 3 – Perform an Experiment
• Observation – a piece of information we
gather using our senses or by taking
measurements.
• Qualitative Observation – observations of
qualities.
• Word descriptions.
• Quantitative Observation – observations of
quantities.
• Number descriptions.
• Quantitative Observation = Measurement

EXAMPLES
• Blue ________________________
• 10 meters ________________________
• 12 ˚C ________________________
• Old ________________________
• 1.2 light years ________________________
• Salty ________________________

•Step 4 – Analyze Data
•Data – the measurable observations
gathered during an experiment.
• Graphs
• Charts
• Tables

• Step 5 – Create a Conclusion
• If the data supports your hypothesis, you
accept the hypothesis.
• Your data can then be published.
• The experiment can then be repeated by
other scientists following the same
procedure (peer review).
• If the experiment is repeated by many
other scientists your hypothesis may
become a scientific theory or scientific law.

•Step 5 – Create a Conclusion
•Scientific Theory – an explanation for
some phenomenon that is based on
observation, experimentation, and
reasoning.
•Scientific Law – a summary of many
experimental results and observations;
a law tells how things work.

•If the data does not support your
hypothesis, you reject the hypothesis.
•The hypothesis can then be revised and
changed based on the new information
gathered.
•The revised hypothesis can then be
retested.

1.2 Assignments
•1.2 Scientific Method WS
•Scientific Method WS
•Penny Drop LAB

1.3 Laboratory
Basics
• Identify and apply
appropriate safety
procedures.
• Identify, properly name, and
use required lab equipment.

1.3 Laboratory Basics
•Graduated Cylinder –
used to measure
liquid volume.
• Very Accurate

•Beaker – used to
stir, heat, and
measure liquid
volume.
• Rough Estimates
•Beaker Tongs –
used to handle
hot beakers.

•Erlenmeyer Flask –
used to heat and
store substances.
• Reduces splatter
when used to heat
a substance.
•Rubber Stoppers –
used to plug a
flask.

•Test Tube – used
mix, heat, or store
substances.
•Test Tube Holder –
used to hold hot
test tubes.

•Funnel – aids in
pouring liquids into
small openings
without spilling.

•Meter Stick –
used to measure
length in the
Metris System.

•Eye Dropper –
used to transfer
small amounts of
liquids.

•Triple Beam Balance –
used to measure mass
in grams.
• Manually balanced.
•Digital Scale – used to
measure mass in grams.
• Electronically balanced.

•Thermometer – used to measure
temperature.

• Ring Stand – a stand
used to support a ring
clamp or test tube
clamp.
• Ring Clamp – used to
clamp onto a ring stand
to sit a beaker or flask.
• Test Tube Clamp –
used to clamp onto a
ring stand to hold test
tube.

•Hot Plate – used to
heat materials
electronically.
•Bunsen Burner –
used to heat
materials over an
open flame.

1.3 Assignments
•1.3 Laboratory Basics WS
•Student Safety Contract
•Science Safety Test
•Safety in the Science Room WS
•Laboratory Basic Skills LAB

1.4 Scientific
Notation
• Express numbers in both
standard notation and
scientific notation.
• Solve addition, subtraction,
multiplication, and division
problems involving numbers
written in scientific notation.

1.4 Scientific Notation
•Very large and very small number are
often written in scientific notation.

•Coefficient = a number greater than or
equal to 1 and less than 10.
•Base = must be 10
•Exponent = shows the number of decimal
places that the decimal needs to moved
to change the number to standard
notation.

EXAMPLE
•The Earth is
approximately
93,000,000 miles
away from the Sun.
Write 93,000,000
miles in scientific
notation.

EXAMPLE
•A human hair is 0.00001 meters wide.
Write 0.00001 meters in scientific notation.

EXAMPLES
(2.4 ∗ 104
) + (6.8 ∗ 105
) =
(5.7 ∗ 103
) − (2.4 ∗ 102
) =

EXAMPLES
(3.2 ∗ 102
) ∗ (6.8 ∗ 107
) =
(4.5 ∗ 10−6
) ∗ (4.5 ∗ 107
) =

EXAMPLES
(9.3 ∗ 108
)
(2.3 ∗ 106)
=
(1.3 ∗ 10−8
)
(7.3 ∗ 106)
=

1.4 Assignments
• 1.4 Scientific Notation WS
• Using Scientific Notation in
Measurements LAB
• Operations with Scientific Notation WS
• Scientific Notation WS

1.5 Significant Figures
• Define uncertainty.
• Identify the number of significant
figures in a measurement.
• Round numbers to the correct
numbers of significant figures or
decimals.
• Calculate answers and determine
the proper number of significant
figures or decimals.

•Uncertainty – the possibility of error in
a measurement.
• When measurements are taken most tools
are not precise and accurate enough to
get exact measurements. To compensate
for this scientists, use significant figures.

• Significant Figure – digits that
carry meaning in a
measurement.
• Significant Figures = Sig Figs
• Sig figs are certain (known)
numbers.
• Sig figs determine how answers
are rounded during calculations.
• Sig figs are necessary in science
because they represent
measurements as accurately as
possible.

•When measurements are taken
• All certain digits are recorded.
• The last digit is uncertain and you must
estimate the digit.

•Rounding Review

EXAMPLES
•Round the following number to the
nearest thousands place.
• 8,832
• 45,832
• 625.36

EXAMPLES
nearest hundreds place.
• 11,694
• 149.49
• 38

EXAMPLES
nearest ones place.
• 893.67
• 21,232.1
• 9.3354

EXAMPLES
nearest hundredths place.
• 0.26598
• 1,612.3658
• 10

•How do you determine sig figs?

•The Atlantic Ocean
is on our right when
we look at a map.
•The Pacific Ocean is
on our left when we
look at a map.
•You are a swimmer.

• If a decimal is ABSENT
you start swimming
on the ATLANTIC side
of the number.
• You can only “swim”
through zeros.
• Once you hit a
number between 1
and 9 you stop
“swimming”.
• All the numbers left
(including zeros) are
significant.

EXAMPLES
•Determine the number of significant
figures in each number.
• 345,000
• 43,001
• 235

• If a decimal is PRESENT you
start swimming on the
PACIFIC side of the number.
• You can only “swim”
through zeros.
• Once you hit a number
between 1 and 9 you stop
“swimming”.
• All the numbers left
(including zeros) are
significant.

EXAMPLES
•Determine the number of significant
figures in each number.
• 1,230.00
• 2.3600
• 214.0001

•Multiplying and Dividing
• The answer should have the same number
of sig figs, as the number with the fewest
sig figs in your problem.

EXAMPLES
•How many significant figures should each
answer have? Calculate the answer.
• 834 * 1.002 =
• 7.3 / 2342 =
• 43 * 3.453 =

•Adding and Subtracting
• The answer should have the same number
of decimal places, as the number with the
fewest decimal places in the problem.

EXAMPLES
•How many decimal places should each
answer have? Calculate the answer.
• 834.7 + 1.002 =
• 7.3 - 2342 =
• 43.4345 + 3.453 =

EXAMPLES
•Write the following numbers in scientific
notation, using the given number of sig figs.
• 1,000,000 with two significant figures.
• 1,000,000 with three significant figures.
• 2,232,450 with two significant figures.

1.5 Assignments
•1.5 Significant Figures WS
•Significant Figures Practice WS
•Significant Figures in the Lab LAB

1.6 Units
• Identify metric and English
units of measurement.
• Explain why scientists use SI
units.
• Measure quantities using
appropriate units for
measurement.

1.6 Units
•Unit – a quantity adopted as a standard
of measurement.
• Measurements must include a quantity
and a unit.
•The two most commonly used units of
measurement are
• English (Imperial) System
• Metric System

1.6 Units
•The English System
• Distance = inch, foot, yard, mile
• Mass = ounce, pound, ton, slug (1 slug = 12 blobs)
• Volume = ounce, cup, pint, quart, gallon
• Time = second, minute, hour, day, year
• Temperature = Fahrenheit
Many English units were based off body
parts of influential people and varied
from region to region.

1.6 Units
•The Metric System
• Distance = cm, m, km
• Mass = gram, kg
• Volume = milliliter, liter
• Time = second, minute, hour, day, year
• Temperature = Celsius
The international prototype kilogram is
made of 90% platinum and 10% iridium.
This mixture of metals is extremely
resistant to environmental factors that
may affect its mass. It is held under very
tight security in St. Cloud, France.

1.6 Units
•Since 1960, scientists worldwide have
used a set of units called the International
System (Le Systeme Internationale in
French) or SI.

1.6 Units
1.6 Assignments
• 1.6 Units WS
• Creating Units and Unit Conversions LAB

1.7 Unit Conversions
• Define conversion factor.
• Convert from one unit to
another using conversion
factors and dimensional
analysis.

•To change between units of the same
measurement scientists use conversion
factors.

•Conversion Factor – a ratio of the equality
of two different units of the same
measurement.
• Can be used to convert from one unit to
another.

1 hr = 60 min 1 min = 60 sec 1 km = 1000 m 7 days = 1 week
24 hrs = 1 day 1 kg = 2.2 lbs 1 gal = 3.79 L 264.2 gal = 1 m3
1 mi = 5,280 ft 1 kg = 1000 g 1 lb = 16 oz 20 drops = 1 mL
365 days = 1 yr 52 weeks = 1 yr 2.54 cm = 1 in 1 L = 1000 mL
0.621 mi = 1.00 km 1 yd = 36 inches 1 cc is 1 cm3 1 mL = 1 cm3

EXAMPLE
• A ruler is 12 inches long. How many
centimeters long is the ruler?

EXAMPLE
• A meter stick is 100 cm long. How
many inches is the meter stick?

EXAMPLE
• A high school cross country race is 5
kilometers. How many miles is a cross
country race?

EXAMPLE
country race? How many feet?

EXAMPLE
country race? How many feet? How
many inches?

•Temperature Conversions
• T°C = Temperature in Degrees Celsius
• T°F = Temperature in Degrees Fahrenheit
• TK = Temperature in Degrees Kelvin

EXAMPLES
•Convert the following temperatures from
Kelvin to Celsius.
𝑻 𝑪 = 𝑻 𝑲 − 𝟐𝟕𝟑. 𝟏𝟓
•34 K
•300 K
•214.6 K

EXAMPLES
Celsius to Kelvin.
𝑻 𝑲 = 𝑻 𝑪 + 𝟐𝟕𝟑. 𝟏𝟓
•49 °C
•24.3 °C
•-36 °C

EXAMPLES
Celsius to Fahrenheit.
𝑻 𝑭 = 𝟏. 𝟖𝟎 𝑻 𝑪 + 𝟑𝟐
•123 °C
•-12 °C
•47.3 °C

EXAMPLES
Fahrenheit to Celsius.
𝑻 𝑪 =
(𝑻 𝑭−𝟑𝟐)
𝟏. 𝟖𝟎
•101 °F
•36 °F
•-45 °F

1.7 Assignments
• 1.7 Unit Conversions WS
• Creating Units and Unit Conversions LAB
• Conversions WS
• Temperature Conversions WS

1.8 Precision and
Accuracy
• Define and differentiate
between precision and
accuracy.
• Calculate percent error.

1.8 Precision and Accuracy
•Accuracy – how close a measurement is
the actual value.
• How “close” your attempts are to the target.

•Precision – how close the measurements
are to on another.
• How “close” your attempts are to one
another.
• The exactness of a measurement.

•Precision and accuracy depend on…
• What you are measuring with.
• Who is doing the measuring.

•Percent Error – calculation of the accuracy
of measurements in an experiment.
• Theoretical Value = actual, known value
• Experimental Value = value determined
during an experiment

EXAMPLE
• The theoretical value for the mass of piece of lead is
59.8 grams. If the measured value is 56.1 grams, what
is the percent error?

1.8 Assignments
•1.8 Precision and Accuracy WS
•Accuracy LAB
•Precision and Accuracy Practice WS

Unit 1: Introduction to Chemistry
Unit 1 Test Review ASSIGNMENT
•Unit 1 Test Review WS

Unit 1 - Introduction to Chemistry (2017/2018)

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