Transcript Chapter 1
Chapter 1 Introduction: Some Basic Concepts
Welcome to the World of Chemistry
Chemistry
The study of matter – its nature, its structure (how it is related to its atoms and molecules), properties, transformations, and its interactions with energy Gold Mercury
Matter
Anything that has mass and occupies space
Mass vs Weight
weight = force = mg g: gravitational acceleration
Mass is a measurement of the quantity of matter in a body or sample
Weight is the magnitude of Earth’s attraction to such a body or sample
Physical States (Phases)
Example of bromine, Br 2 , a halogen
Solid
definite shape and volume
made of particles (atoms, molecules, or ions) held close together and rigidly in place
reasonably well understood.
Example: Graphite — layer structure of carbon atoms reflects physical properties.
Liquid
definite volume but indefinite shape
made of particles (atoms, molecules, or ions) held close together but allowed to move relative to each other
fluid and may not fill a container completely
not well understood
Gas
indefinite volume and indefinite shape
the same shape and volume as their container
made of particles (atoms or molecules) separated from each other by large distances and that move very fast
fluid
good theoretical understanding
Physical Property
characteristic of matter that can be observed without changing the basic identity of the matter
characteristics that are directly observable
eg. state, size, mass, V, color, odor, melting point (T m ), boiling point (T b ), density, solubility...
Chemical Property
characteristic of matter that requires change in identity of the matter for observation (a chemical reaction)
characteristic that describes the behavior of matter
eg. flammability, corrosiveness, bleaching power, explosiveness, ...
Scientific Method
Observation Hypothesis * Law
*
*
Theory * experiment and then modify
Scientific Method
Procedure designed to test an idea Tentative explanation of a single or small number of observations Careful noting and recording of natural phenomena General explanation of natural phenomena Generally observed occurrence in nature 12
Relationships Between Pieces of the Scientific Method
13
Hypothesis explanation for an observation
–
falsifiable – a tentative interpretation or – confirmed or refuted by other observations
–
tested by experiments invalidated – validated or when similar observations are consistently made, it can lead to a Scientific Law
–
a statement of a behavior that is always observed
–
summarizes past observations and predicts future ones
–
Law of Conservation of Mass A theory is a unifying principle that explains a body of facts and the laws based on them.
It is capable of suggesting new hypotheses.
Classification of Matter
Mixture
A combination of pure substances in which the components retain their identities (no reaction)
Can be separated into simpler mixtures and/or pure substances by
Physical Separation Methods
mechanical: eg. sand and iron filings
filtration: eg. sand and water
extraction: eg. washing clothes, decaffeinating coffee
distillation
chromatography
Distillation
Simple - for separation of volatile component from non-volatile component(s)
Distillation
Fractional - for separation of multiple volatile components from each other. Employed in many chemistry labs, labs, and in crude oil refining.
Chromatography
Mixture placed in mobile phase (gas or liquid). Mobile phase flows over and through stationary phase (solid or liquid). Mixture components separate based on relative affinity for mobile and stationary phases.
Heterogeneous Mixture
inconsistent composition
atoms or molecules mixed not uniformly
contains regions within the sample with different characteristics
eg. pizza, carpet, beach sand, ...
Homogeneous Mixture
solution
consistent composition throughout
atoms or molecules mixed uniformly
eg. air in a room, glass of tap water
Compound
can be broken down to 2 or more elements by chemical means
constant composition
eg. water, H 2 O, by mass H:O = 1:8 hydrogen peroxide, H 2 O 2 , H:O = 1:16
elements combined lose individual identities
more than 20 million compounds are now known
Elements
basic substances of which all matter is composed
pure substances that cannot be decomposed by ordinary means to other substances.
made up of atoms
~ 117 known at this time
given name and chemical symbol Aluminum Bromine
Element Symbols
1, 2 or 3 letters:
first letter always capitalized
usually first letter(s) of name H hydrogen O oxygen C carbon N nitrogen Na sodium Cl chlorine Mg magnesium Al aluminum P phosphorus K potassium Po polonium
learn Latin names where appropriate, antimony - Sb - stibium gold - Au - aurum tungsten - W – wolfram sodium – Na – natrium potassium – K - kalium
elements from 104 to 111 are named after scientists; 112-118 have 3 letter symbols based on Latin name for number
112 Uub ununbium 113 Uut ununtrium 114 Uuq ununquadium 115 Uup ununpentium 116 Uuh ununhexium Homework: learn the names of first 36 elements in the periodic table
Periodic Table
a listing of the elements arranged according to their atomic numbers, chemical and physical properties
VERY useful and important
Physical Change
transformation of matter from one state to another that does not involve change in the identity of the matter
examples: boiling, subliming, melting, dissolving (forming a solution), ...
Chemical Change
transformation of matter from one state to another that involves changing the identity of the matter
examples: rusting (of iron), burning (combustion), digesting, formation of a precipitate, gas forming, acid-base neutralization, displacing reactions...
Intensive Property
independent of amount of matter
eg. density, temperature, concentration of a solution, specific heat capacity...
Extensive Property
depends on amount of matter
eg. mass, volume, pressure, internal energy, enthalpy, ...
Density
mass (g) mass (g) Density =
=
volume (cm 3 ) volume (mL)
density of H 2 O is 1.00 g/cm 3 (pure water at ~ 4 °C) 1cm 3 = 1mL Mercury Platinum Aluminum liquid 13.6 g/cm 3 21.5 g/cm 3 They sink in water 2.7 g/cm 3
Know and Own and Practice Well
Metric System
SI Units
Unit Conversions
Learn a Conversion Factor Between English and Metric for
–
length, mass, volume, pressure
SI Units
Système International d’Unités
A different base unit is used for each quantity.
Prefixes
A
prefix
in front of a unit increases or decreases the size of that unit.
makes units larger or smaller than the initial unit by one or more factors of 10.
indicates a numerical value.
prefix
1 kilo meter 1 kilo gram
= value
= 1000 meters = 1000 grams
Metric and SI Prefixes
Learning Check
Indicate the unit that matches the description.
1. A mass that is 1000 times greater than 1 gram.
1) kilogram 2) milligram 3) megagram 2. A length that is 1/100 of 1 meter.
1) decimeter 2) centimeter 3) millimeter 3. A unit of time that is 1/1000 of a second.
1) nanosecond 2) microsecond 3) millisecond
Learning Check
Select the unit you would use to measure A. your height.
1) millimeters 2) meters 3) kilometers B. your mass. 1) milligrams 2) grams 3) kilograms C. the distance between two cities.
1) millimeters 2) meters 3) kilometers D. the width of an artery.
1) millimeters 2) meters 3) kilometers
Volume
1 m = 10 dm (1m) 3 = (10 dm) 3 1m 3 = 1000 dm 3 = 1000 L 1 dm = 10 cm (1dm) 3 = (10 cm) 3 1dm 3 = 1000 cm 3 = 1000mL
Equalities
• • Equalities
use two different units to describe the same measured amount. are written for relationships between units of the metric system, U.S. units, or between metric and U.S. units. For example,
1 in = 2.54 cm 1 m = 1000 mm 1 ft = 12 in 1 lb = 16 oz 1 mile = 5280 ft 2.205 lb = 1 kg 1 L = 1.057 qt 1 hour = 60 min 1 lb = 454 g 1 gal = 4 qt 1 cm 3 = 1 cc = 1 mL
Conversion Factors
•
A
conversion factor
is a fraction obtained from an equality.
Equality :
1 in. = 2.54 cm • •
is written as a ratio with a numerator and denominator.
can be inverted to give two conversion factors for every equality.
1 in.
2.54 cm
and
2.54 cm 1 in.
Learning Check
Write conversion factors for each pair of units.
A. liters and mL Equality: 1 L = 1000 mL B. hours and minutes Equality: 1 hr = 60 min C. meters and kilometers Equality: 1 km = 1000 m D. micrograms and grams Equality: 1 µg = 10 -6 g
Conversion Factors in a Problem
• •
A
conversion factor
may be obtained from information in a word problem.
is written for that problem only.
Example 1:
The price of one pound (1 lb) $2.39.
1 lb red peppers $2.39
and of red peppers is $2.39
1 lb red peppers
Example 2:
The cost of one gallon (1 gal) 1 gallon of gas and $3.95
of gas is $3.95.
$3.95
1 gallon of gas
Percent as a Conversion Factor
• • • •
A
percent factor
gives the ratio of the parts to the whole.
% = Parts x 100 Whole uses the same unit to express the percent.
uses the value 100 and a unit for the whole.
can be written as two factors. Example: A food contains 30% (by mass) fat. 30 g fat 100 g food and 100 g food 30 g fat
Density as a conversion factor
Density of a mineral oil = 0.875 g/mL 0.875 g oil 1 mL and 1 mL 0.875 g oil
Learning Check
Write the equality and conversion factors for each of the following.
A. square meters and square centimeters B. jewelry that contains 18% (by mass) gold C. One gallon of gas is $4.00
Solving: Given and Needed Units
• •
To solve a problem Identify the
given
unit Identify the
needed
unit.
Example: A person has a height of 2.0 meters. What is that height in inches?
The
given unit
is the initial unit of height.
given unit = meters (m)
The
needed unit
answer.
is the unit for the
needed unit = inches (in.)
Problem Setup: Dimensional Analysis
• • • •
Write the given and needed units.
Write a unit plan to convert the given unit to the needed unit.
Write equalities and conversion factors that connect the units.
Use conversion factors to cancel the given unit and provide the needed unit.
Unit 1 x Unit 2 Given x
unit
= Unit 2 Unit 1
Conversion = Needed factor unit
Setting up a Problem
How many minutes are 2.5 hours?
Given unit Needed unit Unit Plan
= = 2.5 hr
=
min hr → min Setup problem to cancel hours (hr). Given Conversion Needed unit factor unit 2.5 hr x 60 min = 150 min (
2 SF )
1 hr
Learning Check
A rattlesnake is 2.44 m long. How many centimeters long is the snake?
1) 2440 cm 2) 0.0244 cm 3) 24.4 cm 4) 244 cm
Using Two or More Factors
•
Often, two or more conversion factors are required to obtain the unit needed for the answer.
Unit 1 → Unit 2 → Unit 3
•
Additional conversion factors are placed in the setup to cancel each preceding unit Given unit x factor 1 x factor 2 = needed unit Unit 1 x Unit 2 Unit 1 x Unit 3 Unit 2 = Unit 3
Example: Problem Solving
How many minutes are in 1.4 days?
Given unit:
1.4 days
Plan:
Factor 1 Factor 2 days → hr → min
Set up problem:
1.4 days x 24 hr x 60 min = 2.0 x 10 3 min 1day 1 hr
2 SF Exact Exact = 2 SF
Learning Check
A bucket contains 4.65 L of water. How many gallons of water is that?
Unit plan:
L → qt → gallon
Equalities:
1.06 qt = 1 L 1 gal = 4 qt
Set up Problem:
4.65 L x 1.06 qt x 1 gal = 1.23 gal 1 L 4 qt 3 SF Exact Exact 3 SF
Learning Check
If a ski pole is 3.0 feet in length, how long is the ski pole in mm?
Solution:
3.0 ft x 12 in x 2.54 cm x 10 mm = 1 ft 1 in. 1 cm Calculator answer: Needed answer: 914.4 mm 910 mm (2 SF rounded)
Check factor setup: Check needed unit:
Units cancel properly mm
Learning Check
If your pace on a treadmill is 65 meters per minute, how many minutes will it take for you to walk a distance of 7500 feet?
Solution: Given: Plan:
7500 ft 65 m/min Needed: min ft → in. → cm → m → min
Equalities:
1 ft = 12 in. 1 in. = 2.54 cm 1 m = 100 cm 1 min = 65 m (walking pace)
Set Up Problem:
7500 ft x 12 in. x 2.54 cm x 1m 1 ft 1 in.
x 1 min 100 cm 65 m = 35 min final answer (2 SF)
(# 11): Ethylene glycol, C 2 H 6 O 2 , is an ingredient of automobile antifreeze. Its density is 1.11 g/cm 3 at 20 °C. If you need exactly 500. mL of this liquid, what mass of the compound, in grams, is required?
Needed: m(g) Given: d(g/cm 3 ) and V(mL) 1 cm 3 = 1 mL 500. mL
1.11 g ───── = 555 g 1 mL 3 SF
(# 13): A chemist needs 2.00 g of a liquid compound with a density of 0.718 g/cm 3 . What volume of the compound is required?
Needed: V(cm 3 ) Given: d(g/cm 3 ) and m(g) 2.00 g
1 cm 3 ────── = 2.78 cm 3 0.718 g 3 SF
(# 15): A sample of 37.5 g of unknown metal is placed in a graduated cylinder containing water. The levels of the water before and after adding the sample are 7.0 and 20.5 mL respectively. Which metal in the following list is most likely the sample?
Metal d(g/mL) Metal d(g/mL) Mg Fe Ag 1.74
7.87
10.5
Al Cu Pb 2.70
8.96
11.3
The volume of sample = volume of water displaced in cylinder = 20.5 – 7.0 = 13.5 mL one dec. place
After placing the piece of metal Before
(# 15):
Needed: d(g/cm 3 ) Given: V(mL) and m(g) d = m 37.5 g ── = ────── = 2.78 g/mL V 13.5 mL 3 SF From the list, the metal is Al.
Accuracy: nearness of the measurement to accepted value of the quantity.
Precision: reproducibility; how well several determinations of the same quantity agree.
Consider a sample that was analyzed for lead content and was known to contain 49.3 ppm lead. Two analyses Analysis A Analysis B diff. from Trial ppm Pb 1 38.9
2 3 23.2
55.9
4 80.1
5 46.9
average = 49.0 ppm Pb Trial ppm Pb average 1 48.9 4.6
2 59.8 6.3
3 54.5 1.0
4 49.0 4.5
5 55.3 1.8
53.5 ppm average diff. 15.5 ppm 3.6 ppm More accurate More precise
Numbers
magnitude, value
direction: sign (+ or −)
type of measurement: units
precision of original measurement: significant figures
Measured Numbers
A measuring tool
•
is used to determine a quantity such as height or the mass of an object.
provides numbers for a measurement called
measured numbers
.
~4.56 mL
Reading a Meter Stick
. l 2 . . . . l . . . . l 3 cm . . . . l . . . . l 4 . . •
The markings on the meter stick at the end of the orange line are read as the first digit 2
• •
plus the second digit 2.7 The last digit is obtained by
estimating
.
The end of the line might be estimated between 2.7
of 2.7
5
–2.8 as half-way (0.5) or a little more (0.6), which gives a reported length cm or 2.7
6
cm.
Known & Estimated Digits
In the length reported as 2.76 cm,
• • •
The digits 2 and 7 are
certain (
known
).
The final digit 6 was
estimated (
uncertain
).
All three digits (2.76) are
significant including the
estimated digit.
Significant Figures in Measured Numbers
Significant figures •
obtained from a measurement include all of the known digits
plus the estimated digit.
•
reported in a measurement depend on the measuring tool.
Significant Figures
Examples of Counting SF
143.22
143.0
300592
0.0020930
100.0
100.
100
Exact numbers have an unlimited number of significant figures
A number whose value is known with complete certainty is exact
–
from counting individual objects
–
from definitions 1 cm is exactly equal to 0.01 m
–
from integer values in equations in the equation for the radius of a circle, the 2 is exact
SF in Calculations
Addition and/or Subtraction
perform operation
round answer to same number of digits after decimal as number in calculation with the fewest
Example
132.09 + 35.94376 – 0.0173 =
132.
09 + 35.94376 – 0.0173 = 168.01646
must have 2 digits after decimal 168.
02
Multiplication and/or Division
perform operation(s)
round answer to same number of significant figures as number in the calculation with the fewest
Example
(26.894)(0.0837)/13 = (26.894)(0.0837) = 0.1731560
13
must have 2 SF 0.
17
Log and/or Antilog
number of digits in mantissa of log = number of significant figures in antilog
Example
log (14.8003) =
log ( 14.8003
) = 1.17027051857
antilog
1.
170271
mantissa
Rounding
if first digit to be eliminated is ≥ 5, round preceding digit up one
if first digit to be eliminated is <5, truncate
Examples
Round each of the following to 4 SF
10.02700
10.03
10.02495
10.02
10.02502
10.03
10.02500
10.03
10.01500
10.02
10.02500000000000000000000001
10.03
Exponentials
Scientific notation: very large or very small numbers are expressed in the following general form: exponent term, n = ± integer N
x 10
n digit term, between ± 1 and 9.9999… (coefficient) eg −12,760,000 = −1.276 x 10 7 0.000012760 = 1.2760 x 10 -5
Write the following in scientific notation: 22,400 = 2.24 x 10 4 22,400. = 2.2400 x 10 4 892 x 10 5 = 8.92 x 10 7
-0.00198 x 10 -10 =
-1.98 x 10 -13
127.60 x 10 -5 =
1.2760 x 10 -3
Write in fixed notation: 5.720 x 10 -2 = 0.05720
-1.982 x 10 4 = -19,820
Exponentials in calculations
5.750 x 10 3 + 7.25 x 10 2 = 1.75 x 10 -3 x 6.45 x 10 2 57.50 x 10 2 + 7.25 x 10 2 1.75 x 10 -3 x 6.45 x 10 2 = 64.75 x 10 2 1.75 x 10 -3 x 6.45 x 10 2 = 64.75 x 10 2 1.75 x 6.45 10 -3 x 10 2 =
5.74 x 10 3 3 SF
(0.000345 – 0.0001273) x 6.730x10
3
= 154.00 6 dec places (we keep the 7 7, 0.000217
7 x 6.730x10
3
= 154.00
though) 2.17
7 has 3 SF only 2.17
7 x10
4 x 6.730x10
3
= 0.00951
3 1.5400 x 10 2 = 9.51x10
3 3 SF
Temperature
• • • Temperature
is a measure of how hot or cold an object is compared to another object.
indicates that heat flows from the object with a higher temperature to the object with a lower temperature. is measured using a thermometer.
Temperature Scales
Temperature Scales are Fahrenheit, Celsius, and Kelvin.
have reference points for the boiling and freezing points of water.
Temperature Scales
Fahrenheit ( °F )
Celcius or Centigrade ( °C ) Kelvin (K)
9 °F = ── 5 °C +32 = 1.8 °C + 32
5 ( °F - 32) °C =
( °F -32) =
9 1.8
K = °C + 273 273.15 (exact) ΔT(K) = ΔT(°C) variation of temperature
Learning check
a. The normal temperature of a chickadee is 105.8
°F. What is that temperature on the Celsius scale? 1) 73.8
°C 2) 58.8°C 3) 41.0
°C b. A pepperoni pizza is baked at 235 °C. What temperature is needed on the Fahrenheit scale? 1) 267 °F 2) 508 °F 3) 455°F c. Convert 204.3 K into °C.
Other elements to remember
Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn
Ru, Rh, Pd, Ag, Cd
Pt, Au, Hg