CHEMISTRY Chapter 1 & 2 Matter, Measurements, and
Calculations
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Chapter 1 Section 1 Objectives: 1.Define chemistry 2.List
examples of branches of chemistry 3.Compare and contrast basic
research, applied research, and technological development
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What objects in this room are related to chemistry? Plastics
Fabrics Clothes Cooking oil Motor oil Make-up Radio Batteries
Computers
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Chemistry in our daily lives. Antibiotics Food Transportation
Sports Farming Military Industry
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Chemistry Study of the composition and properties of matter and
the changes that matter undergoes -What something is made of -What
is the internal arrangement
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Chemical Any substance that has a definite composition
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6 Main Branches of Chemistry 1.Organic substances containing C
2.Inorganic substances other than organic 3.Biochemistry living
things 4.Physical chemistry changes of matter 5.Analytical
chemistry id components of materials 6.Theoretical chemistry use
math and computers to understand chemical behavior
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All branches involve some type of research. Basic research to
increase knowledge -how and why Applied research to solve problems
Technological development production and use of products - lags
behind discoveries - application of knowledge
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Review and Assignment 1. Define chemistry 2. List examples of
branches of chemistry 3. Compare and contrast basic research,
applied research, and technological development Assignment: WS
1-1
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Quiz 1.Name two branches of chemistry. 2.List two ways that
chemistry affects our daily lives. 3.Definition of chemistry.
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Chapter 1 - Matter
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Chapter 1 Section 2 Objectives: 1.Distinguish between a mixture
and a pure substance. 2.Define what matter is.
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Matter -anything that has mass and occupies space -includes
almost everything -exceptions are light, heat, and sound
-properties are used to measure matter ex. mass Mass measure of
quantity of matter - not affected by temp, location, or any other
factor
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Demo. Mass vs. matter What caused the change in mass? Is air
matter?
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Matter (cont.) Classified into 2 groups: 1. pure substances 2.
mixtures Pure substance matter that has the same properties
throughout ex. element or compound
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Pure Substances Element substance that cannot be broken down by
ordinary chemical change - only 1 type of atom - symbols
abbreviated w/1 or 2 letters - can be an allotrope allotrope one of
a number of different molecular forms of an element in the same
state Compound substance made up of 2 or more elements chemically
combined - can be broken down by chemical change - more than 1 type
of atom
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Compounds 1.Elements that make up a compound are combined in
definite proportion by mass ex. 100 g water has 11.2 g H and 88.8 g
of O 2. Chemical and physical properties of compound differ from
those of its parts ex. water is liquid, H and O are gases 3.
Compounds can be formed from simpler substances by chem change and
can be broken down into simpler substances
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example 100 of water has 11.2 g H and 88.8 g O How many g of H
is in a 120g sample of water? 120 g water| 11.2 g H = 13.4 g H |
100 g water
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Mixtures - contain 2 or more substances that have different
properties - vary in composition and properties from sample to
sample ex. rock, wood, salt water -Not chemically combined -Can be
separated by simple physical means -ie. filtration, evaporation,
distillation
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Formation of Mixtures A mixture can be formed 3 ways: 1.Element
mixed w/1 or more other elements ex. carbon w/sulfur 2. Compound
mixed w/ 1 or more other compounds ex. salt w/sugar 3. 1 or more
elements mixed w/1 or more compounds ex. sulfur w/sugar
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Characteristics of Mixtures - retain properties of each of its
parts ex. iron and sulfur - iron remains magnetic - composition can
vary widely - can be homogeneous or heterogeneous
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Types of mixtures Homogeneous uniform composition throughout -
called solutions ex. alloys, pop, air, coffee Heterogeneous not
uniform throughout ex. concrete, soil, dry soup, spaghetti and meat
balls
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Matter Pure substance ElementCompound Mixture
HomogeneousHeterogeneous
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Review and Assignment 1. Distinguish between a mixture and a
pure substance. 2. Define what matter is. Assignment: WS
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Chapter 1 Section 2 Objectives: 1. Distinguish between the
physical properties and chemical properties of matter. 2. Classify
changes of matter as physical or chemical. 3. Explain the gas,
liquid, and solid states in terms of particles.
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Properties of Matter -allow us to distinguish btwn substances
-characteristics of a substance -what can be observed -way that a
substance behaves ex. color, taste, odor, gas, liquid, solid
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Properties (cont.) - can be extensive or intensive Extensive
d/o amount of matter ex. volume, weight, mass, and E Intensive does
not d/o amount of matter ex. melting point, boiling point, density,
and conductivity
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Demonstration Properties - water and glycerin How do they
compare? - look, feel, weight, flow - water and salt water How do
they compare? - conductivity
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Physical Properties Can be observed or measured w/out changing
the substance Can describe the substance Odor, taste, hardness,
density, melting point, and boiling point Metals ductile (pulled
into wire), malleable (hammered into sheets), luster (shine), good
conductors
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Chemical Properties A transformation of a substance into a
different one rusting, flammability, tarnishing, new substance
formed
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Physical Change No new substance is formed CHANGE IN PHASE,
pounding, grinding, cutting Changes of phase When a substance
changes phase there is no change in composition Physically
different, chemically the same Solid, liquid, or gas are the three
states of matter
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States of Matter Solid definite volume and shape Particles are
in fixed positions Held w/strong attractive forces Liquid definite
volume and no definite shape Takes shape of container Particles can
move past each other
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States of Matter (cont.) Gas neither definite volume nor
definite shape Particles move easily and are very far apart Plasma
high temperature state in which atoms lose their electrons
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Chemical Change One or more substance is changed to something
new Rusting, burning, gas formed, digestion, heat or light added,
explosion, color change, odor change, water formed
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Review and Assignment 1. Distinguish between the physical
properties and chemical properties of matter. 2. Classify changes
of matter as physical or chemical. 3. Explain the gas, liquid, and
solid states in terms of particles. Assignment: p. 18 and WS
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CHEMISTRY Chapter 1 Section 3 Objectives: 1.Perform density
calculations. 2.Describe conservation of mass.
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Properties of Matter -E is always involved in both physical and
chemical changes -Physical are not at noticable -Chemical are more
noticable -Heat and light are given off
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Density is a physical property is always the same for a solid
substance in gases and some liquids a change in temperature will
change the density increase in temperature will decrease density D
= m/V
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Density problem Use the 5 steps in problem solving to solve the
following problem. Lead has a mass of 22.7 g and its volume is 2.00
cm 3. What is its density? m = 22.7 g V = 2.00 cm 3 D = m/V = 22.7
g/2.00 cm 3 = 11.4 g/ cm 3
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Examples
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Conservation of Mass In reactions matter cannot be created or
destroyed by a chemical change - mass stays the same, it may just
change form
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Density Lab Results Group 1 Group 2 Group 3 Group 4 Group 5
-
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Review and Assignment 1. Perform density calculations. 2.
Describe conservation of mass. Assignment: WS and Density lab
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Chapter 2 - Sec.1 Objectives: 1.Describe the purpose of the
scientific method. 2.Distinguish between qualitative and
quantitative observations. 3.Describe the steps to making a graph.
4.Distinguish between inversely and directly proportional
relationships.
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Scientific Method - a logical approach to solving problems 1.
Make observations -observe your surroundings 2. State the problem -
stated as a question 3. Collect data 4. Form hypothesis - testable
statement 5. Test hypothesis 6. Conclusion 7. Modify hypothesis and
retest
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Observing Involves making measurements and collecting data Data
can be qualitative or quantitative Qualitative non-numerical
information - descriptive (the sky is blue) Quantitative numerical
information - the mass is 25.7 grams
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Conclusion Can be explained by using models Model explanation
of how phenomena occur or how things are related - visual - verbal
- mathmatical
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Theory -models may become part a theory Theory broad
generalization that explains facts or phenomena - must be able to
predict results ex. kinetic-molecular theory collision theory
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Controlled Experiments Use manipulated variable (independent)
Use responding variable (dependent) One variable manipulated at a
time Measurements are called data
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Making a Graph Shows results of an experiment in a meaningful
pattern Dependent variable is on the vertical axis 1. Always
include a title 2. Determine variables 3. Set up scale 4. Plot
points 5. Draw best-fit line
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Oxygen obtained from electrolysis of water
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Relationships in graphs Directly proportional if dividing one
by the other gives you a constant value If one increases so does
the other If started at point (0,0) Inversely proportional if their
product is constant If one increases the other decreases Produce a
curve
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Review and Assignment 1.Describe the purpose of the scientific
method. 2.Distinguish between qualitative and quantitative
observations. 3.Describe the steps to making a graph. 4.Distinguish
between inversely and directly proportional relationships.
Assignment: graphing WS
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Quiz 1.List three steps of the scientific method. 2.List two
steps in making a graph.
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Chapter 2 Sec.2 Objectives: 1. Distinguish between a quantity,
a unit, and a measurement standard. 2. Name SI units for length,
mass, time, volume, and density. 3. Distinguish between mass and
weight.
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Measurements Basic part of science Make observations more
meaningful Needs to be more than just a number or quantity Need a
common system of units For consistency Measure your desk w/anything
you have available
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SI System -The International System of Units -Used in all
science -A standard -Based on 10 -Makes it easier to convert from
one unit to another
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SI System (continued) -7 base units 1. Length meter (m) 2. Mass
kilogram (kg) 3. Time second (s) 4. Amount mole (mol) 5.
Temperature Kelvin (K) 6. Electric Current ampere (amp) 7. Luminous
intensity candela (cd)
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Weight vs. mass Mass quantity of matter - how much space it
takes up - measured w/a balance - unit kg Weight F gravity pulls on
matter with - measured w/spring scale - unit Newton On the moon
will our weight or mass stay the same?
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SI Prefixes You must know these. Kilo- 1000 Deca 10 Base unit
(m, s, L) Centi 1/100 or 0.01 Milli 1/1000 0r 0.001
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Derived Units -combination of base units Examples - Area = m 2
- Volume = m 3 - Density = kg/m 3 - Newton = m kg/s 2
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Derived Units (cont.) Area determined by multiplying 2 lengths
Volume determined by multiplying 3 lengths for a solid - for
liquids unit is cm 3 or mL ** 1 mL = 1 cm 3
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Review and Assignment 1. Distinguish between a quantity, a
unit, and a measurement standard. 2. Name SI units for length,
mass, time, volume, and density. 3. Distinguish between mass and
weight. Assignment: WS 2-2 and p. 42 ~1-3
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Quiz 1.What is the base SI unit for mass? 2.Kilo = ______
3.Centi = _____ 4.What is a derived unit? 5.1 cm 3 = _____ mL
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Chapter 2 - Sec.3 Objectives: 1.Distinguish between accuracy
and precision. 2.Determine the number of significant figures in
measurements. 3.Perform mathematical operations involving
significant figures.
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Accuracy and Precision Accuracy closeness of a measurement to
correct value Precision closeness of a set of measurements to each
other Consistency Do not have to be correct d/o measuring
instrument
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Bullseyes
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Significant Figures -digits in a measurement that are know with
certainty and one digit that is estimated -CALCULATORS DO NOT KEEP
TRACK OF SIGNIFICANT FIGURES
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Significant Figure Rules 1. Digits other than zero are ALWAYS
significant ex. 61.4 3 sig. fig. 2. All zeros at the end of a
number and to the right of the decimal with a # preceding the
decimal are ALWAYS sig ex. 4.7200 km 5 sig. fig. 3. Zeros used only
for spacing are NOT significant ex. 7000 1 sig. fig. 201 sig. fig.
100.04 sig. fig. 4. Zeros between sig. fig are significant 5. Zeros
in front of a non-zero are NOT sig. - dont count until you get to 1
st non-zero from lf to rt 0.004 1 sig. fig. 0.00091 sig. fig.
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Significant Figures 1,000 = _____ sig figs 100.0 = _____ sig
figs 0.00012340 = _____ sig fig 10.0340 = _____ sig fig
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Calculating w/Significant Figures Addition and Subtraction -
use same # of decimal places as the measurement w/the least decimal
places ex. 2.098 3 DECIMAL places +6.21 DECIMAL place 8.298round to
1 Decimal 8.3 is the final answer
Calculating w/sig. figs (cont.) Multiplication and Division -
use same # sig. fig. as the measurement w/the least sig. fig.
ex.2.38 3 sig. fig x 9.02 sig. fig 21.42 round to 2 sig. Fig 21 is
the final answer
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Multiplying and Dividing 100.0 x 10 = _____ 34.56 x 23.45 =
_____ 12.045 x 34.008 = _____ 50.04 x 23 = _____
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Review and Assignment 1.Distinguish between accuracy and
precision. 2.Determine the number of significant figures in
measurements. 3.Perform mathematical operations involving
significant figures. Assignment: WS 2-6 and sig fig WS
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Quiz How many significant figures are in the following numbers?
1. 8,000 _____ 2. 100.01 _____ 3. 0.00056_____ 4. 4500.10 _____ 5.
What is precision?
Percent Error Observed value based on lab measurements True
value based on generally accepted references Error exists in any
measurement d/o measurer, instrument, conditions
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Percent Error % error = true value obs. value x 100 true value
Example atomic mass of Al = 28.9 g measured mass = 27.0 g What is
the % error? 28.9 g 27.0 g x 100 = 7.00 % 28.9 g
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Review and Assignment 1.Perform mathematical operations
involving percent error. Assignment: WS 2-5 and % error WS
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Quiz How many significant figures are in the following numbers?
1. 8,104 _____ 2. 100.01 _____ 3. What does % error tell us? 4.
What is accuracy? 5. What is precision?
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CHEMISTRY Chapter 2 Sec.3 Day 3 Objectives: 1.Use dimensional
analysis to convert measurements. 2.Convert measurements into
scientific notation. 3.Perform mathematical operations using
exponents.
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Problem Solving Rules Write down what is known. - mass = 346
gvolume = 34.6 cm 3 2. Write down unknown. - density = ? 3. Write
the equation to use. D = m/V 4. Fill in knowns. D = 346 g/34.6 cm 3
5. Solve for unknown and label. D = 200 g/cm 3 6. Check your
work.
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Dimensional Analysis - use with conversion factors to change
from one unit to another Steps: convert 2550 m to km 1. Determine
conversion factor - 1000 m to 1 km 2. Set up T-bars 3. Write given
# in first box
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Dimensional Analysis (cont.) 4. Write conversion factor in 2 nd
box - unit on bottom matches unit of given # 5. Matching labels
cancel - if 1 from conversion factor is on top divide - if 1 from
conversion factor is on bottom multiple
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Scientific Notation -Used to represent very large or very small
numbers -There are two parts -Basic form is M x 10 n -M is a number
- n is a number representing how many places to move the
decimal
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Scientific Notation (cont.) If n is negative, your number is a
decimal If n is positive, your number is a large number Examples:
60,000,000 = 6 x 10 7 0.000005 = 5 x 10 -6 125,000 = 1.25 x 10
5
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Scientific Notation (cont.) Write the following in scientific
notation. 1,000,000,000 23,456 0.0005678 0.034 14,239.1
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Scientific Notation (cont.) Write the following in long hand.
1.1 x 10 -9 2.3.5 x 10 5 3.7.123 x 10 -3 4.5 x 10 2 5.4.56 x 10
-2
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Multiplication w/exponents Step 1 Multiply coefficients Step 2
Add exponents ex. (2 x 10 2 ) (2.5 x 10 5 ) = 5 x 10 7
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Division w/exponents Step 1 Divide coefficients Step 2 Subtract
exponents ex. (5 x 10 -2 ) (1.0 x 10 7 ) = 5 x 10 -9
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Addition & Subtraction w/exponents All numbers must be
written in the same power of 10 ex. 5.8 x 10 3 + 2.16 x 10 4 -
change to 0.58 x 10 4 + 2.16 x 10 4 = 2.74 x 10 4
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Scientific Notation & sig figs All numbers in front of the
x 10 are significant ex. 2.00 x 10 2 = 3 sig fig 2 x 10 2 = 1 sig
fig
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Scientific Notation & calculators 5.44 x 10 7 /8.1 x 10 4
5.44 (EE or exp) 7 / 8.1 (EE or exp) 4 = 6.7 x 10 2
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Review and Assignment 1.Use dimensional analysis to convert
measurements. 2.Convert measurements into scientific notation.
3.Perform mathematical operations using exponents. Assignment: p.
57 ~ 1-7 and WS