Download Chapter 1. Introduction: Matter and Measurement

Chapter 1. Introduction: Matter and Measurement
Chemistry: Study of properties, composition and structure of matter and the
chemical and physical changes that matter undergoes.
•
Chemistry applies to all aspects of life. E.g. production of food, clothes,
gasoline, development of drugs, conservation of natural resources, etc.
•
Chemistry is the central science - helps you understand related fields such
as biology, physics, geology & engineering.
MATTER: Anything that has mass & occupies space.
• Approximately 100 Elements constitute all matter
3 Physical States of Matter
1. Gas: indefinite shape; indefinite volume; molecules are far apart & move
very rapidly
2. Liquid: indefinite shape; definite volume; molecules are close together, but
still move pretty fast (slower than for gas)
3. Solids: definite shape; definite volume; molecules are packed closely & are
relatively rigid
Composition of Matter
1. Pure substance: Consist of 1 Substance with fixed composition and distinct
properties; cannot be separated by physical means.
A. Element: Made up of unique atoms & cannot be chemically separated
into simpler substances. E.g. He, O2 , C
B. Compounds: Consist of 2 or more different elements & can be
chemically broken down into simpler substances. E.g. H2O, NaCl, C6H6
Atom: smallest unit particle of an element that retains the chemical identity
of that element.
Molecule: 2 or more atoms chemically combined.
2. Mixture: Physical combination of 2 or more substances; variable
composition; components can be separated by physical means. E.g. filtration
or distillation can be used to separate different substances
A. Heterogeneous: 2 or more different phases present; composition
and properties are not uniform. E.g. sand, rocks, egg
B. Homogeneous: 1 phase with same composition, appearance, and
properties. E.g. salt water, air, gin & tonic
Elements
•
>112 known elements
•
Main Elements found in earth's crust: O, Fe, Si, Al, Ca; human body: H, O, C
Compounds
•
Compounds have different properties than elements contained in compound
(E.g. H2O vs. H2 and O2)
Law of Constant Composition: A compound always consists of the same
combination of elements (% composition is fixed)
PROPERTIES OF MATTER
1. Physical properties: properties observed without changing a substance.
e.g. color, odor, taste, melting pt, physical state, density
2. Chemical properties: describe the chemical reactivity of a substance. E.g.
H2 is flammable, dynamite is explosive.
3. Intensive properties: Do not depend on amount of substance present. E.g.
melting point, density, color, temperature
4. Extensive properties: Depend on quantity of substance present.
E.g. mass, volume, pressure
CHANGES:
Physical Change: Substance changes physical form w/o changing its identity.
Often involves change in physical state. E.g. melting of ice, breaking of glass
Chemical Change: Chemical composition of the substance changes. Old
substance is destroyed; new substance is formed. E.g. burning of gas,
decomposition of H2O
Scientific Method:
1. Design experiment & collect data (observe, experiment)
2. Analyze data & develop hypothesis
hypothesis: tentative explanation of observations
3. Test hypothesis to prove or disprove it
theory: hypothesis that has been tested and validated
law: concise statement that summarizes facts about a certain
phenomena - not an explanation. A law involves a measurable quantity &
usually is expressed as a mathematical equation.
1.4 SI Units (System International)- preferred metric units
•
7 Base Units (length, mass, time, temperature, mole - Table 1.4)
•
Metric prefixes are used to change the size of the unit to larger or
smaller units. (Know prefixes from Table 1.5: G, M, k, d, c, m, µ, n)
1) Length: SI unit is meter (1 m = 1.0936 yd)
•
Know 2.54 cm = 1 inch
2) Mass: measurement of amount of matter; not affected by gravity
SI unit = kilogram
(1 kg = 2.2 lb)
3) Temperature: SI unit is Kelvin
•
Temperature Scales based on freezing & boiling point of water
°C
100
0
BP
FP
•
•
K
373
273
°F
212
32
Kelvin & Celsius have same size degree, scale is just shifted by 273
O K is lowest temperature - absolute Zero; no molecular motion
Temperature conversions:
K = °C + 273
°F =
9
°C + 32
5
Derived SI units
1. Volume: SI unit = m3;
•
cm3 and L more commonly used
1 cm3 = 1 cc = 1 ml
2. density: mass per unit volume
m
v
•
d=
•
units usually expressed as
d (H2O) = 1.0
•
g
cm 3
g
g
g
or
or
3
cm
ml
L
d (air) = 0.001
g
cm 3
d (Au) = 19.3
g
cm 3
density of gases is low; density of liquids & solids is high (solids usually
have highest densities)
1.5 Uncertainty in Measurement
Precision: how well measured quantities agree with each other.
Accuracy: how well measured quantities agree with the true value.
E.g. dart board analogy
Two kinds of numbers
1. Exact: counted or defined numbers; infinite number of significant figures
E.g. 12 eggs in a dozen; 2.54 cm = 1 inch
2. Inexact: measured numbers have finite number of significant figures
(measured by ruler, scale, speedometer, etc.)
Significant Figures: All digits known plus one uncertain digit
Rules for counting
1) Non-zero numbers are always significant. E.g. 185.27 has 5 sig figs
2) Zeros between numbers are always significant. E.g. 305.6 has 4 sig figs
3) Zeros before the first non-zero digit are not significant - they just locate the
position of the decimal pt. E.g. 0.0049 has 2 sig figs
4) Zeros at the end of the number and after a decimal point are significant.
E.g. 6.7000 has 5 sig figs
5) Zeros at the end of a number without a decimal point are ambiguous.
E.g. 28500 has 3,4 or 5 sig figs; use exponential notation to indicate exact
number of sig figs. E.g. 2.85 x 103 has 3 sig figs, 2.8500 x 103 has 5 sig figs
Calculations
Multiplication and division
Answer has the same number of significant figures as the value with the least
number of significant figures. E.g. 6.221 cm x 5.2 cm = 32 cm2 (not 32.3492)
Addition and subtraction
Answer is reported to the least number of decimal places.
E.g. 20.4 g - 1.3222 g = 19.1 g
1.6 Dimensional Analysis
• Method of calculation in which the given quantity is multiplied by 1 or more
conversion factors to obtain desired quantity.
• A conversion factor is a ratio where the numerator & denominator are
100 cm
equivalent, but possess different units. E.g.
1m
• For correct set-up, all units cancel except for desired units.
* Note the back cover of text has some of the important conversion factors.
Example Problems.
1. Perform the following conversions:
a. 336 Mg to g:
 10 6 g 
 = 3.36 x 108 g
336 Mg 
1
Mg


b. 2.75 kg to cg:
-2
c. 4.6 x 10
 103 g  1cg 
 −2  = 2.75 x 105 cg
2.75 kg 
 1 kg  10 g 
-2
µm to mm:
4.6 x 10
 10 −6 m 

µm 
1
µ
m


 1 mm 
 −3  = 4.6x10-5 mm
 10 m 
d. Convert 1.35 x 109 km3 to L.
 1000 m 

1.35x10 km 
 1km 
9
3
3
 10 dm 


 1m 
3
 1L 

 = 1.35x1021 L
3 
1
dm


2. A cube 1.5 cm on a side has a mass of 1.9 g. What is the density in g/cm3 ?
l = 1.5 cm
m = 1.9 g d = ? g/cm3
Volume = l3 = (1.5 cm)3 = 3.375 cm3
d=
m
1.9 g
=
= 0.56 g/cm3
3
v
3.375 cm
(2 sig figs)
3. The density of a piece of ebony wood is 1.20 g/cm3 . What is the volume of
5.74 kg of wood?
 1000 g  cm 3 

 = 4.78 x 103 cm3
5.74 kg 
 1 kg  1.20 g 
4. a) Convert 82 °F to °C
5
5
5
°C = (°F − 32) = (82-32) = (50.) = 28 °C
9
9
9
b) Convert 25 °C to K & °F
K = °C + 273 = 25 + 273 = 298 K
9
9
°F = °C + 32 = (25) + 32 = 77 °F
5
5
5. Convert 45.7 in/hr to mm/s
45.7
in
hr
 2.54 cm  1 m  1000 mm  1 hr




 1in  100 cm  1 m  60 min
 1 min

 60 s

mm
 = 0.322
s

6. Density of air is 1.19 g/L. What is the mass in kg of air in a room that
measures 12.5 x 15.5 x 8.0 ft? 1 m = 1.0936 yd
g
V = 1550 ft3
d = 1.19
l
Paths: ft3 → yd3 → m3 → L & g → kg
 1 yd 

1550 ft 
3
ft


3
3
 1m



1
.
0936
yd


3
 1dm 
 −1 
 10 m 
3
 1 L  1.19 g  1 kg 



 = 52 kg
3 
1
L
1000
g
1
dm




7. A car travels 28 miles per gallon. How many kilometers per Liter will it go?
1 mile = 1.6093 km; 1 gallon = 3.7854 L
28
miles
gallon
 1 gallon  1.6093 km 


 = 12 km/L
3
.
7854
L
1
mile


