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Unformatted text preview: CHEMISTRY 1307
General Chemistry I
Today’s Agenda 1. Welcome/Introduction
2. Discussion of the Syllabus http://www.uhd.edu/ 3. Class Objectives/Goals http://www.uhd.edu/ 4. Chapter 1 Discussion 1 CHEMISTRY 1307
General Chemistry I
Instructor: Dr. Byron K. Christmas
Office: N-809 Phone: (713) 221-8169
Phone: E-Mail: [email protected]
Office Hours: See Syllabus
Text: Chemistry: A Molecular Approach – Nivaldo J. Tro
Other Requirements: i>Clickers and MGC
Other 2 CHEMISTRY 1307
Prerequisites: CHEM 1305 or High School Chemistry
Co-requisite: Credit or Enrollment in MATH 1301 and
On-Line Mastering General Chemistry Homework/Quizzes:
Examinations: 3 Examinations & 1 Comprehensive Final
Examination 3 CHEMISTRY 1307
EVALUATION OF PERFORMANCE: 3 Examinations (Final Exam Replaces Lowest) 51% 4 MGC Quizzes (Lowest Quiz is Dropped!) 15%
15% 11 MGC Homework Assignments (Lowest 2 Scores Drop!) 9%
1 i>Clicker Participation 4%
4% 1 Final Examination (Standardized, Comprehensive!) 21% 4 CHEMISTRY 1307
We will now do “Classroom Enrollment” for your i-Clicker!
When you see your name scrolling down, note the letter next to
Your name and press that key on your remote. A second letter
Should then appear by your name. Click on that letter to
Confirm you registration! 5 CHEMISTRY 1307
i>Clicker Question # 1
The measurement, 0.0100 g has how many significant figures?
E.1 6 Why Study Chemistry?
1. To better understand the world: what it is made of
and how it works.
2. Because it is the most practical and relevant of the
sciences - chemistry is the study of EVERYTHING!
3. It is the “Central Science” - All other sciences
intersect at and depend on chemistry.
4. It is essential to the national and local economies.
(Houston is at the center of the world’s largest petrochemical complex) 7 Why Study Chemistry?
5. It is required for virtually every major involving
science, mathematics, or engineering.
6. An awareness of the principles of chemistry is essential
to being an informed and responsible citizen in a highly
7. It is incredibly fascinating and a lot of fun!
8. And MANY MORE!
8. And 8 The Language of Chemistry
3 Alphabet - Chemical Symbols of the Elements (Memorize the
first 103 names and symbols)
H Sc Zr Ta Xe Nd Np
3 Words - Chemical Formulas of Compounds
3 3 Mg Re H2O
HNO3 C12H22O11 NaCl NH4ClO4
Sentences - Chemical Equations
2 Na (s) + Cl2 (g) -----> 2 NaCl (s) + energy
Paragraphs - Reaction Mechanisms
(Dealt with in Chapter 13 in CHEM 1308) 3 Using the Language to Express Ideas Definitions, concepts, mathematical skills, etc. Master the Language and
You Master the Subject!
You 9 The Periodic Table of the Elements 10 The Periodic Table of the Elements
Page 59 1 Million pounds of Molybdenite gave 1 gram in 1928! 11 Chapter 1
Matter, Measurement, &
Problem Solving 12 CHEMISTRY: (Memorize this Definition!)
Chemistry is the study of the properties, composition,
and structure of matter, the physical and chemical
matter the physical
changes it undergoes, and the energy liberated or
absorbed during those changes.
Got your notebooks? Write this DOWN! MATTER:
Matter is anything that occupies space and
the earth paint
rocks 13 Let’s Think About Matter!
What Are the States of Matter? B. Gases
C. Liquids Temperature A. Plasma D. Solids
Page 6 in Textbook 14 Atomic and Molecular Concepts
Atoms or Molecules Liquid Temperature Nuclei Atoms or Molecules Crystalline Solid 15 Classification of Matter
Substances Elements Compounds Mixtures
Heterogeneous Memorize 16 Atomic and Molecular Concepts
Classification of Matter
• Substance - A distinct type of matter. All
samples of a substance have the same properties. Elements and compounds are subElements and compounds are substances. • Mixture - A sample of matter consisting of
two or more substances which are NOT
chemically combined. 17 Page 7 in Textbook 18 Classification of Matter (Substances)
3 Element - A substance that cannot be broken down
(decomposed) into simpler substances by chemical
reactions. 3 Compound - A substance composed of two or more
elements chemically combined in fixed ratios by mass.
Water - H2O
Carbon dioxide - CO2
Sodium Chloride - NaCl
Na Iron(II) sulfide – FeS
Iron(III) sulfide – Fe2S3
Fe 19 Classification of Matter (Mixtures)
Mixtures 3 Homogeneous - A mixture having only one phase; it is uniform (the same) throughout and has the
same properties throughout. These are called
3 Heterogeneous - A mixture with more than one
phase. It is non-uniform and does NOT have the
same properties throughout. 20 Classification of Matter
Substances Elements Compounds Mixtures
Heterogeneous Memorize 21 i>Clicker Question # 2: A homogeneous mixture is also known as
A.A molecule 22 i>Clicker Question # 3: An example of a pure substance is
A.Both C. and D. are correct! 23 Classification of Matter
Substances Elements Compounds Mixtures
Heterogeneous Memorize 24 i>Clicker Question # 4: Which of the following is correct for the material
pictured? a gaseous pure substance
3 a liquid pure substance
3 a gaseous mixture
3 a solid mixture
3 none of the above
3 25 Matter and Change
Physical Change - A change in which
each substance involved in the change
retains its original identity and no new
elements or compounds are formed.
Evaporating 26 Matter and Change
Page 8 in Textbook Physical Changes 27 Matter and Change
Page 9 in Textbook Physical Changes 28 Matter and Change
Page 11 in Textbook i-Clicker Question # 5:
Which one below represents
H2O after vaporization? H2O (l) molecules in pan
pan 29 Matter and Change
3 Chemical Change - A change in which one or
more elements or compounds (substances) are
formed. 2 H2 (g) + O2 (g) 2 H2O (l)
“Reacting” AgNO3 (aq) + HCl (aq)
HNO AgCl (s) +
AgCl 30 Matter and Change
Page 9 in Textbook Chemical
Change! 31 Matter and Change
Page 10 in Textbook 32 Properties of Matter
3 Physical Properties - Properties that do NOT involve
substances changing into other substances.
Volume 3 Chemical Properties - Properties that involve
substances changing into other substances.
Oxidation Potential 33 Measurement
3 Chemistry is an Observational science.
science. 3 Chemistry is a Quantitative science.
Quantitative science. 3 Measurement - A quantitative
observation. Page 21 in Textbook o 103.3 F
103.3 34 Measurement
All measurements have three parts:
1. A value
26.9762 g 2. Units
3. An Uncertainty
Examples: 33.2 mL
33.2 72.36 mm 426 kg 31 people 35 i>Clicker Question # 6: When reading a graduated cylinder, read the volume at
the bottom of the meniscus.
What volume of liquid is in the graduated cylinder?
E. 4 mL
4.57 36 Measurement
Systems of Units - Standards of Measurement
Standards 1. The Need for Standards
2. The English System (What a pain!!!)
16 oz/lb 3 ft/yd
2000 lb/ton 3. The Metric System - A decimal system
gram (g) - Length
Mass 37 Measurement
1 m = 1000 mm
1 mL = 0.001 L
1 kg = 1000 g = 1 000 000 mg
10 cm = 0.01 m = 0.000 01 km
23 kL = 23 000 000 000 µ L 4. The SI System - Système International d’Unitès
A. A complete system of units adequate for
the entire realm of physical science.
the 38 SI System of Measurement Prefixes for the Metric and SI Systems
Page 14 in Textbook! 39 SI System of Measurement Prefixes for the Metric and SI Systems
Page 17 in Textbook! 40 SI System of Measurement Common Conversions for the Metric and SI Systems
Page 18 in Textbook! 41 SI System of Measurement
5. Rules for Using the SI Systems
1. Use only singular form of units and do NOT
use a period after the symbol for the unit.
2. Use a dot on the base line for the decimal
3. Group digits in threes around the decimal
point and do NOT use commas. 200 000.000 003 km 42 SI System of Measurement
4. Do NOT use spaces for four-digit
measurements. 1645 mL
1645 or 0.2367 µ g 5. Do NOT use the degree sign (o) for
temperature recorded for the Kelvin
temperature scale. 78.6 K not 78.6 o K 43 Measurement
Conversion Factors - A fraction whose
numerator and denominator contain the same
quantity expressed in different units.
1 mile = 5280 ft 1 mile = 5280 ft = 1
1 mile 1 cm = 0.01 m 1 cm
= 0.01 m = 1
1 cm 1 in = 2.54 cm 2.54 cm =
1 in = 1
2.54 cm 44 Measurement
6. Uncertainty in Measurements Exact Measurements: Measured values determined
by counting or when a value is defined.
by counting Examples: 31 people
31 27 rocks
6 2.54 cm = 1 in 10 µ L = 1L The uncertainty in these measurements = 0
Non-exact Measurements: All other measurements.
The last digit recorded is uncertain; it is estimated!!
uncertain it estimated!! Examples: 27.5 g
27.5 32.7 mm 12 467 km 2 1.156 x 10 mL 45 i>Clicker Question # 7: What is the temperature recorded
on the thermometer? o A. 103.3 C
o B. 103.3 F
o 46 Accuracy, Precision, & Sensitivity
Accuracy - The degree to which a measured
value agrees with the true or “accepted” value.
Precision - The reproducibility of a measured
Sensitivity - The “fineness” of a measured
value; the number of significant figures it has.
23.5673 g is a more sensitive measurement than 23.57 g.
23.57 47 Accuracy vs. Precision Page 26 in Textbook! 48 i>Clicker Question # 8: A student measures the mass of a penny 4 times
and records the following data. What can be said
about the data if the actual mass of the penny is
2.4987 g? The data is both accurate
3 The data is neither
accurate nor precise.
3 The data is accurate, but
3 The data is not accurate,
but it is precise.
1 Mass, g
2.5104 2 2.5106 3 2.5102 4 2.5109 49 Measurement
Significant Figures: Each digit obtained as a result
of a measurement. This includes all of the certain
digits and the first uncertain digit. The number of
significant figures in a measurement is an indicator
of the SENSITIVITY of the measurement.
1.5 How many significant figures are in the following:
65 mL 2 173.4 g 4 12.2 m 3 1 x 109 ns
ns 1 50 Measurement
The Problem with Zero:
207.1 mm 0.002 36 mm 260.1 mm 0.123 00 mm
.123 00 2040.0 mm 3600 mm Rules for Significant Figures:
All non-zero digits are significant.
A zero between other significant figures is
100.7 51 Measurement
Initial zeros are NOT significant.
0.001 23 cm3
Final zeros after the decimal point ARE
Final zeros in a measurement with no decimal
point may or may not be significant.
3200 cm (might have 2, 3, or 4
Exact measurements have an infinite number of
significant figures. (They are CERTAIN!!)
CERTAIN!! 52 Measurement Significant Figures in Calculations: In a
measurement, the last significant figure is assumed to
The result of a calculation involving measured values
can be no more certain than the least certain
The number of significant figures in a result depends
on the number of significant figures in the measurement and on the mathematical operation being
performed. 53 Measurement Significant Figures in Calculations:
Addition and Subtraction - A sum or a difference of two or more measurements has the
same number of decimal places as the measuredecimal
ment with the least number of decimal places.
35.2 mL - 0.34 mL = ____ mL
35.2 1.007 94 u + 1.007 94 u + 15.9994 u
= ______ u
u = atomic mass units
http://www.youtube.com/watch?v=ZuVPkBb-z2I&feature=related 54 Measurement
Multiplication and Division - A product or
quotient of two or more measurements has the
same number of significant figures as the measuresignificant
ment with the least number of significant figures.
significant density = (9.5760 g)/(12.2 mL)
= _0.785 g/mL
Round-off Rules - For digits 0 - 4, do not round up.
For digits 5 - 9, round up.Page 24 in text! 55 Measurement 1.6 Round-off the following to two decimal places: 23.044 39 g = 23.04 g 65.89
65.891 mL = _____ mL 30.11
45.106 ms = 45.11 ms 30.1149 kg = _____ kg
_____ 37.995 ng = _____ ng
____ 6. Dimensional Analysis - An extremely useful tool
to help you solve mathematical problems. It is
based on the fact that when doing calculations
involving measured quantities, the units must be
added, subtracted, divided, or multiplied just like
the value of the measurements.
of 56 Dimensional Analysis
1.7 How many meters are in each of the following?
21 km 1023 570 µ m
1023 (21 km)(1 x 103 m) = 21 x 103 m =
2.1 m 10 m
(1023 570 µ m)( 1 m )
= 1.023 570
(1 x 106 µ m) 57 Dimensional Analysis
1.8 How many mL are in 3.0 ft3?
1 ft = 12 in 1 in = 2.54 cm 3 1 cm = 1 mL = _________ mL
1.9 How many ns are in 23.8 s?
(23.8 s)(109 ns) = 23.8 x 109 ns = 2.38 x 1010 ns
(1 s) 58 Properties of Matter
3 Extensive Properties - Properties that depend on the
amount of matter present in a sample.
Heat Capacity 3 Intensive Properties - Properties that do NOT depend
on the amount of matter present in a sample.
Boiling Point 59 Mass and Weight
Mass: the measure of the quantity or amount of
matter in an object. The mass of an object does not
change as Its position changes.
Mass is measured using a BALANCE.
BALANCE Weight: A measure of the gravitational attraction of
the earth for an object. The weight of an object
changes with its distance from the center of the earth.
Weight is measured using SCALES.
SCALES 60 Sample Calculations Involving Masses
1.1 How many µ g are in 2.56 kg?
1.1 = ________ µ g
1.2 How many g are in 2.578 x 1012 ng? = ____ g
____ 61 Sample Calculations Involving Volumes
1.3 How many mL are in 3.456 L? = ____ mL
1.4 How many µ L are in 23.7 cm3?
1.4 = _______ µ L 62 Density
Density - The mass of a unit volume of a material.
Density The density = mass/volume
1.5 What is the density of a cubic block of wood that is
2.4 cm on each side and has a mass of 9.57 g? volume = [2.4 cm x 2.4 cm x 2.4 cm]
density = (9.57 g)/(____ cm3)
= _____ g/cm3 = ____ g/mL
3 63 i>Clicker Question # 9: Which of the following has the
A. B. C. A material that has a
mass of 10.0 g and a
volume of 2.00 L
A material that has a
volume of 5.00 g and a
volume of 10.0 cm3
A material that sinks in
ethanol but floats on
water 64 Temperature and Thermal Energy
Temperature: A measure of the “hotness”and “coldTemperature: measure
ness” of an object; a measure of the average kinetic
energy of the atoms and molecules of the object.
The higher the temperature, the more kinetic energy
the atoms and/or molecules have. This is an
http://www.youtube.com/watch?v=v1zOnyC4RgQ&feature=related Thermal Energy: Often called “heat”, it is the form
of energy toward which all other forms tend to go. This
is an EXTENSIVE property.
EXTENSIVE property. 65 Sample Calculations Involving Temperatures
1.6 Convert 73.6oF to Celsius and Kelvin temperatures. C = (5/9)( F - 32) K = C + 273.15
oC = (5/9)(73.6oF - 32) = (5/9)(41.6)
K = 23.1oC + 273.15 = 296.3 K
273.15 296.3 66 Temperature Scales
Page 15 in
Textbook 67 ...
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This note was uploaded on 07/22/2011 for the course CHEM 1307 taught by Professor Staff during the Fall '10 term at University of Houston - Downtown.
- Fall '10