Unit 1 Measurement and Matter - Mr Palermo
Transcript of Unit 1 Measurement and Matter - Mr Palermo
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Unit 1 Math and Measurement
Mr. Palermo
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LESSON 1:
Metric Conversions
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What is Chemistry?
The study of Matter and the changes it
undergoes…..
What is Matter?
Matter is anything that has mass and takes up
space.
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Matter Can be Described as:
Qualitative measurements: descriptive, non-numerical observations
Quantitative Measurements: are in the form of NUMBERS and UNITS.
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Quantitative Measurements:
The METRIC SYSTEM (SI): System of
measurement used in science and in most
countries
The BASE UNITS of measurement:
(Found on Reference Table D)
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Table D (base units)
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PREFIXES: Used to modify base units of measurement.
(Found on Reference Table C)
Example:
gram (g)
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Converting Units Using table C
1. Find the difference between the exponents
of the two prefixes on Table C.
2. Move the decimal that many places.
To the left
or right?
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Move the decimal to the:
LEFT when going from a smaller prefix to
a larger prefix
RIGHT when going from a larger prefix to
a smaller prefix 52.010
52.010
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Where are the base units?
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Example 1: convert 5.2 cm = ____ mm
The difference between the two factors (-2 and -3) is 1.
Since you are moving from a larger prefix to a smaller prefix you move the decimal one place to the right.
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Example 2: convert 45.5 mm = ____ m
The difference between the two factors (-3 and 0) is 3.
Since you are moving from a smaller prefix to a larger prefix you move the decimal three places to the left.
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Example 3
Convert the following:
20 cm = _________________ m
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LESSON 2:
Density
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Quantitative Calculations:
Mass: the amount of matter an object
contains. (This is different than weight,
which is mass plus gravity)
Volume: The amount of space a
substance occupies
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How do we measure mass in
the lab?
Electronic Balance
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How can we measure volume?
l x w x h (regular solid)
ex. V = 1cm3
Graduated cylinder (liquids)
Read bottom of MENISCUS
ex. V = 27.5 mL
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Reading a Meniscus
10
8
6
proper line of sight reading correct
graduated
cylinder
10 mL
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Measuring Volume: Irregular Solid
Water displacement
method
1. Measure initial volume
2. Measure final volume
with object
3. The Difference is the
volume of the object
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Example: What is the volume of
the solid?
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Density
Ratio of mass of an object to its volume
Use density formula
Located on Table T
V
MD
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Example 1
What is the density of an object with a
mass of 60 g and a volume of 2 cm3?
V
MD
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Example 2
An object has a volume of 825 cm3 and a density of 13.6 g/cm3. Find its mass.
V
MD
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How to solve for mass or volume if
density is not given: Use TABLE S
Example:
The volume of an aluminum sample is 251
cm3. What is the mass of the sample?
The density of aluminum on table S is
2.70g/cm3
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LESSON 3:
Temperature Conversions
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Temperature:
Measure of average kinetic Energy
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Temperature Scales
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Celsius Scale
Freezing point of water at 0°C.
Boiling for water at 100°C.
Below 0 is NEGATIVE.
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Kelvin Scale Water freezes at 273K and boils at 373K
Theoretical point of ABSOLUTE ZERO is when all
molecular motion stops
NO NEGATIVE NUMBERS
Divisions (degrees) are the same as in Celsius
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Converting Between Temperature
Scales
Formula: K = °C + 273
Located on Table T
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Example 1:
What is the temperature in Kelvin of an
object that is 55°C ?
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Example 2:
What is the temperature in Celsius of an
object that is 150 K?
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LESSON 4:
Percent Error
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Accuracy vs. Precision
Accuracy - how close a measurement is to
the accepted or true value
Precision - how close a series of
measurements are to each other
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Example
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EXAMPLE:
Student A
(g/cm3)
Student B
(g/cm3)
Student C
(g/cm3)
Trial 1 1.54 1.40 1.70
Trial 2 1.60 1.68 1.69
Trial 3 1.57 1.45 1.71
Avg. 1.57 1.51 1.70
Range 0.06 0.28 0.02
These students were asked to determine the density of sucrose. Sucrose has a density of 1.59 g/cm3. Which student is more accurate?
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STOP AND THINK
Which student is more precise in the
previous example?
Student A
(g/cm3)
Student B
(g/cm3)
Student C
(g/cm3)
Trial 1 1.54 1.40 1.70
Trial 2 1.60 1.68 1.69
Trial 3 1.57 1.45 1.71
Avg. 1.57 1.51 1.70
Range 0.06 0.28 0.02
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Percent Error Measurement of ACCURACY
the % that the measured value is “off” from accepted value
Measured value = value you “get”
Accepted value = value you “should get”
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Formula is found on Table in your Reference Table:
If answer is negative, your measured value is LESS THAN the accepted value
If answer is positive, your measured value is GREATER THAN the accepted value
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Example A student determines the density of a substance to
be 1.40 g/mL. Find the % error if the accepted value of the density is 1.36 g/mL.
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Stop and Think
What is the percent error if a student measures an object to be 56.2 cm and the actual length is 56.9 cm?
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Stop and Think
In a lab experiment, you are told by your teacher that the actual amount of sugar in a can of Coke is 39 g. You experimentally determine it to be 37 g based on your own data and calculations. What is your percent error?
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LESSON 5:
Precision & Sig Figs
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Significant Figures
Indicate PRECISION of a measurement.
Recording Sig Figs
Sig figs in a measurement include the known digits plus a final estimated digit (precision of instrument)
2.38 cm
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Example measuring length:
We know for sure that the object is more than ____, but less than ____
We know for sure that the object is more than ____, but less than ____
This ruler allows us to estimate the length to ______
2 3
2.8 2.9
2.85 cm
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2 Runners finish the race in 8
seconds. Who won?
1
2
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Runner 1
Runner 2
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Stop and Think:
What is the length of the red line?
cm 0 1 2 3 4 5
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Stop and Think:
What is the length of the red line?
cm 0 1 2 3 4 5
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Stop and Think:
What is the volume of the liquid?
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LESSON 6:
Counting Sig Figs
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Rules for counting sig figs
1. All non-zero digits are significant. 2. Leading zeros are never significant. ex. 0.421 (3 sig figs)
3. All captive zeros are significant. (Captive is a zero between 2 other non-zero digits.) ex. 4012 (4 sig figs)
4. For Trailing zeros: (zeros after last non-zero digit) - Decimal point significant -No decimal point not significant ex. 114.20 (5 sig figs)
ex. 11,420 (4 sig figs)
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How to Count Sig Figs 1. Start counting from LEFT to RIGHT at first
NONZERO number.
2. If decimal point is present then count any trailing zeros
3. If decimal is not present don’t count trailing zeros
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EXAMPLE
1) 2545.300 g (7 sig figs)
2) 4530 km (3 sig figs)
3) 0.00453 m (3 sig figs)
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Stop and Think:
How many sig figs in 23.500 m?
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Stop and Think:
How many sig figs in 53,000 km?
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Stop and Think:
How many sig figs in 0.0800 g?
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Lesson 7: Rounding Sig Figs
in Calculations
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LESSON 7:
Rounding Sig Figs in Calculations
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What do I round my answer to?
Every measurement has some error in it.
When performing calculations AN
ANSWER CAN NEVER BE MORE
PRECISE THAN YOUR LEAST PRECISE
MEASUREMENT
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Rounding: Sig Fig in Calculations
Multiply/Divide - Round answer to the least
number of significant figures.
Example:
(13.91g/cm3)(23.3cm3) = 324.103g
324 g
4 SF 3 SF 3 SF
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Stop and Think
Round answer to proper # of sig figs
15.30 g ÷ 6.4 mL =
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Stop and Think
Round answer to proper # of sig figs
18.9 g x 0.84 g =
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Calculating Sig Figs (con’t) Add/Subtract – Round to the least place value
Example:
3.75 mL
+ 4.1 mL
7.85 mL
224 g
+ 130 g
354 g 7.9 mL 350 g
3.75 mL
+ 4.1 mL
7.85 mL
224 g
+ 130 g
354 g
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Stop and Think
Round answer to proper # of sig figs
28.9 g - 0.85 g =
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Stop and Think
Round answer to proper # of sig figs
80.4 mm – 16.532 mm =
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LESSON 8:
Scientific Notation
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Scientific Notation
A way to represent large or small numbers
For example:
The mass of a hydrogen atom is
0.00000000000000000000000167g.
2 g of H2 contains
602,000,000,000,000,000,000,000
molecules.
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Scientific Notation is written as:
The product of two numbers: a coefficient and a 10 raised to a power.
The coefficient (number written first) is always a number from 1 to 9
Example:
1.67 x 10-24 g
2 g of H2 is composed of 6.02 x 1023
molecules.
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Converting from Expanded form into Scientific Notation
1. For #’s greater than 1 move decimal to the
LEFT until there’s 1 digit to its left. The
number of places moved = exponent
number
Example: 45,450 g =
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2. For #’s less than 1 move decimal to
RIGHT stopping after the first non zero
number. The number of places moved =
negative exponent number
Example: 0.00453 ml = 4.53 x 10-3 ml
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Stop and Think
Convert to Scientific Notation
45,700 m =
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Stop and Think
Convert to Scientific Notation
0.00009 cm =
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Converting from Scientific Notation to Standard Notation
1. Move the decimal place the number of
times indicated by the exponent.
2. To the right if it is positive.
3. To the left if it is negative.
Example:
4.5 x 10-2 = 0.045
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Stop and Think
Convert to Standard Notation
9.6 x 103 m =
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Stop and Think
Convert to Standard Notation
1.2 x 10-4 g =
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Calculating with Sci Notation using a Calculator
Ex. (5.44 × 107 g) ÷ (8.10 × 104 mol) =
5.44 EXP
EE ÷
EXP
EE ENTER
EXE 7 8.10 4
= 671.60493 = 672 g/mol = 6.72 × 102 g/mol
Type on your calculator:
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![Page 82: Unit 1 Measurement and Matter - Mr Palermo](https://reader033.fdocuments.us/reader033/viewer/2022060422/6295748d65648a0e2b298db3/html5/thumbnails/82.jpg)
Stop and Think
3.95 x 102 ÷ 1.5 x 106 =
![Page 83: Unit 1 Measurement and Matter - Mr Palermo](https://reader033.fdocuments.us/reader033/viewer/2022060422/6295748d65648a0e2b298db3/html5/thumbnails/83.jpg)
Stop and Think
3.5 x 102 x 6.45 x 1010 =