Transcript Matter and Measurement

Matter and Measurement
Chapter 1
What Is Matter?
• Matter is anything that takes up space and has
mass.
• Mass is the amount of matter in an object.
• Mass is resistance to change in motion along a
smooth and level surface.
Types of Matter
• Substance- a particular kind of matter – pure
• Fixed composition
• Distinct properties
• Ex. Water, Salt
• Mixture- more than one kind of matter
• Compositions vary
• Each substance retains its own chemical identity and
properties.
Substances
• Elements- simplest kind of matter
• Cannot be broken down into simpler
• All one kind of atom.
• Compounds- are substances that can be
broken down by chemical methods
• When they are broken down, the pieces have
completely different properties than the
compound.
• Made of molecules- two or more atoms
Mixtures
• Heterogeneous- mixture is not the same
from place to place.
• Chocolate chip cookie, gravel, soil.
• Variable Composition
• Homogeneous- same composition
throughout.
• Kool-aid, air.
• Every part keeps its properties.
Solutions
Homogeneous mixture
•
•
•
•
•
•
•
Mixed molecule by molecule
Can occur between any state of matter.
Solid in liquid- Kool-aid
Liquid in liquid- antifreeze
Gas in gas- air
Solid in solid - brass
Liquid in gas- water vapor
Solutions
• Like all mixtures, they keep the properties of
the components.
• Can be separated by physical means
• Not easily separated- can be separated.
Gold
Gold
Copper
Silver
24 karat gold
24/
24
atoms Au
18 karat gold
18/
24
atoms Au
14 karat gold
14/
24
atoms Au
An alloy is a mixture of metals.
• Brass = Copper + Zinc
• Solid brass
• homogeneous mixture
• a substitutional alloy
Copper
Zinc
Solid Brass
• Brass = Copper + Zinc
• Brass plated
• heterogeneous mixture
• Only brass on outside
Copper
Zinc
Brass Plated
Steel
•
•
•
•
Alloys
Stainless steel
Tungsten hardened steel
Vanadium steel
We can engineer properties
– Add carbon to increase strength
– Too much carbon  too brittle and snaps
– Too little carbon  too ductile and iron bends
Nitinol Wire
• Alloy of nickel and titanium
• Remembers shape when heated
Applications:
surgery, shirts that do not need to be ironed.
Properties
• Physical Properties- A change that changes
appearances, without changing the
composition, thus identity is preserved.
• ex. color, odor, density….
• Chemical Properties- a property that can only
be observed by changing the type of
substance.
• ex. flammability
Physical and Chemical
Changes
Chemical
Changes
Physical
Changes
Rusting Nail
Melting Ice
Bleaching a Stain
Boiling Water
Burning a Log
Sawing a log in half
Tarnishing Silver
Tearing Paper
Fermenting of
Grapes
Souring of Milk
Breaking a Glass
Pouring of Milk
Intensive Vs. Extensive Properties
• Intensive Properties
• Do not depend on on
the amount of sample
being examined
• Aid in the identification
of a substance.
• ex. temperature,
density, melting point….
• Extensive Properties
• Depend on the quantity
of sample
• ex. mass, volume,
area….
The SI System
Physical
Quantities
Name of Unit
Abbreviation
Mass
Kilogram
Kg
Length
Meter
M
Time
Second
S
Electric Current
Ampere
A
Temperature
Kelvin
K
Luminous Intensity
Candela
Cd
Amount of
Substance
Mole
mol
• The SI system has seven base
units from which all others
are derived
SI Units (Con’t)
Prefix
Abbreviation
Meaning
Mega-
M
106
Kilo-
k
103
Deci-
d
10-1
Centi-
c
10-2
Milli-
m
10-3
Micro-

10-6
Nano-
n
10-9
Pico-
p
10-12
Femto-
f
10-15
•
These prefixes indicate
decimal fractions or
multiples of various units
Temperature Conversions
• °C = 5/9 ( °F-32)
• °F = 9/5 (°C ) +32
•
K = °C + 273.15
At home you like to keep
the thermostat at 72 F. While
traveling in Canada, you find
the room thermostat calibrated
in degrees Celsius. To what
Celsius temperature would you
need to set the thermostat to
get the same temperature you
enjoy at home ?
Derived Units
• SI units are used to derive the units of
other quantities.
• These units express speed, velocity, area
and volume.
• They are either base units squared or
cubed, or they define different base units
Volume
• Calculated by multiplying L x W x H
• Basic SI unit of volume is the cubic meter (m3 ).
• Smaller units are sometimes employed ex. cm3, dm3 ….
• Volume is more commonly defined by liter (L).
Density
Density is an intensive property of all substances; this
means that it is a universal characteristic of the
substance and does not change because of extensive or
accidental properties such as amount, size or location.
Density is the relationship between the mass and volume of
an object. It is defined as:
• Density = Mass of substance
Volume of substance
• Density is usually expressed in g/ml
Specific Gravity
Specific Gravity is a ratio between the density of a
substance and the density of water. It is defined by:
Specific Gravity = Density of Sample
Density of Water
In calculations involving s.g., the units of density must
match in order for the units to cancel
Problems Involving Density
1. The mass of 325 mL of the liquid methanol is found to be 257 g.
What is the density of methanol?
2. A lead weight used in the belt of a scuba diver has a mass of 226g. When
the weight is carefully placed in a graduated cylinder containing 200.0ml of
water, the water level rises to 220.0ml. What is the density of the lead weight
(g/ml)
3. A sample of a certain material has a mass of 2.03x 10-3 g. Calculate the
Volume of the sample given that the density is 9.133 x 10-1 g/cm3
Uncertainty and Measurement
Exact numbers
• Numbers whose values are
known exactly
• e.x. 12 eggs in a dozen, 1000g in
a kg
Inexact numbers
• Numbers obtained by measuring
a quantity
• e.x. height, weight or
temperature
• There is always a degree of
uncertainty in measured values.
Why?
Precision Vs Accuracy
• When comparing sets of data points, scientist
want to know two things: which set of data points
are precise and which set are accurate.
Precision
•
How many times a given
measurement can be repeated
with results close in value to
each other.
• Which of the following data is
more precise:
119g, 120g, 128g or
101g, 100g, 99g
Good precision poor
accuracy
Accuracy
• How close an experimental
measurement is to the true,
actual, or book value
• Usually the more accurate a
measurement the more
precise it will be.
Good Accuracy and Precision
Precision and Accuracy
•
Accuracy refers to the agreement of a particular value with
the true value.
Precision refers to the degree of agreement among several
measurements made in the same manner.
Neither accurate
nor precise
4/11/2016
Precise but not
accurate
Precise AND
accurate
28
Significant Figures
• When using a measuring device to measure
anything the last number of the
measurement is always estimated
• Measured quantities are usually reported in
such a way that only the last digit is
uncertain. All digits including the uncertain
digit are called significant figures
Significant Figures
A number is not significant if
it is:
• A zero at the beginning of a
decimal number
ex. 0.0004lb, 0.075m
• A zero used as a placeholder in a
number without a decimal point
ex. 992,000,or 450,000,000
A number is a S.F. if it is:
• Any real number ( 1 thru 9)
• A zero between nonzero digits
ex. 2002g or 1.809g
• A zero at the end of a number or
decimal point
ex. 602.00ml or 0.0400g
• Any digit in the coefficient of a
number written in scientific
notation
ex. 4.0 x 105 m, 5.70 x 10-3
Sig Fig Practice #1
How many significant figures in each of the following?
4/11/2016
1.0070 m 
5 sig figs
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
0.0054 cm 
2 sig figs
3,200,000 
2 sig figs
31
Addition and Subtraction
• When adding or subtracting the
answer is rounded so it has the
same number of decimal places
as the number with the least
number of decimal places
6.23 cm
39.24 cm +
677.1 cm
722.6 cm
Multiplication and Division
• When multiplying or dividing
the answer has the same
number of significant digits as
the number with the least
number of S.F.
2.85ml x 67.4ml = 192ml
ml
49.618g =
43.8ml
Sig Fig Practice #2
Calculation
Calculator says:
Answer
3.24 m + 7.0 m
10.24 m
10.2 m
100.0 g - 23.73 g
76.27 g
76.3 g
0.02 cm + 2.371 cm
2.391 cm
2.39 cm
713.1 L - 3.872 L
709.228 L
709.2 L
1818.2 lb + 3.37 lb
1821.57 lb
1821.6 lb
2.030 mL - 1.870 mL
0.16 mL
0.160 mL
4/11/2016
34
Odd Numbers
• Odd numbers are rounded up, even numbers are left alone
when the remainder is 5
ex. 22.15 g if rounded to 3 SF = 22.2g
22.25 g if rounded to 3 SF = 22.2
Sig Fig Practice #3
Calculation
Calculator says:
Answer
3.24 m x 7.0 m
22.68 m2
100.0 g ÷ 23.7 cm3
4.219409283 g/cm3 4.22 g/cm3
23 m2
0.02 cm x 2.371 cm 0.04742 cm2
0.05 cm2
710 m ÷ 3.0 s
236.6666667 m/s
240 m/s
1818.2 lb x 3.23 ft
5872.786 lb·ft
5870 lb·ft
1.030 g ÷ 2.87 mL
2.9561 g/mL
2.96 g/mL
4/11/2016
36
Scientific Notation
In science, we deal with some very
LARGE numbers:
1 mole = 602000000000000000000000
In science, we deal with some very
SMALL numbers:
Mass of an electron =
0.000000000000000000000000000000091 kg
4/11/2016
37
.
2 500 000 000
9 8 7 6 5 4 3 2 1
Step #1: Insert an understood decimal point
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n
4/11/2016
38
2.5 x
9
10
The exponent is the
number of places we
moved the decimal.
4/11/2016
39
0.0000579
1 2 3 4 5
Step #2: Decide where the decimal must end
up so that one number is to its left
Step #3: Count how many places you bounce
the decimal point
Step #4: Re-write in the form M x 10n
4/11/2016
40
5.79 x
-5
10
The exponent is negative
because the number we
started with was less
than 1.
4/11/2016
41
Review:
Scientific notation expresses a
number in the form:
M x
1  M  10
4/11/2016
n
10
n is an
integer
42
Conversion Factors
Fractions in which the numerator and denominator are
EQUAL quantities expressed in different units
Example:
1 in. = 2.54 cm
Factors: 1 in.
and
2.54 cm
2.54 cm
1 in.
Learning Check
Write conversion factors that relate each of the
following pairs of units:
1. Liters and mL
2. Hours and minutes
3. Meters and kilometers
How many minutes are in 2.5 hours?
Conversion factor
2.5 hr x 60 min
1 hr
= 150 min
cancel
By using dimensional analysis / factor-label method, the UNITS ensure that you
have the conversion right side up, and the UNITS are calculated as well as the
numbers!
Dimensional Analysis
• Many problems in chemistry and health sciences require a
change in units, thus dimensional analysis is used to make
these conversions
• The key is the correct use of a conversion factor to change 1
unit to another
ex. 1hr = 60 min
this allows us to write the relationship:
60 min
1hr
and
1hr
60min
Problems Involving Dimensional
Analysis
1. On a bicycle trip, Maria averaged 35 miles per day.
How many days did it take her to cover 175 miles?
2. To prevent bacterial infection, a doctor orders 4
tablets of amoxicillin per day for 10 days. If each
tablet contains 250 mg of amoxicillin, how many
ounces of the medication are given in 10 days?
Solution to Problems
Step 1: Given
175 miles
35 miles = 1 day
Step 2:
Unit plan
MilesDays
175 miles x 1 day = 5.0 days
35 miles
Solution to Problem
Step 1: Given
10 days
4 tablets/day 250 mg each
Step 2: Unit Plan
mgglboz
1000mg x 1g
x 1lb
x 16 oz = 0.35 oz
1000mg 453.59g 1lb
Dealing with Two Units
If your pace on a treadmill is 65 meters per
minute, how many seconds will it take for you
to walk a distance of 8450 feet?
Learning Check
• A Nalgene water
bottle holds 1000
cm3 of dihydrogen
monoxide (DHMO).
How many cubic
decimeters is that?
Solution
1000 cm3
1 dm
10 cm
(
)
3
= 1 dm3
So, a dm3 is the same as a Liter !
A cm3 is the same as a milliliter.