chemical foundation general chemistry tonya patterson
TRANSCRIPT
Chemical Foundation
General Chemistry Tonya Patterson
Chemistry• Is the study of matter and the changes it
undergoes
• Often referred to as the central science
• Why study chemistry?
Scientific Method• A systematic approach to research• Used by all sciences• Process
o Identify problemo Research and/or observationso Form a hypothesis (tentative explanation or prediction of experimental
observations)o Experiment/Testingo Analysis of datao Draw conclusion
• If hypothesis is correct – finished• If hypothesis is incorrect – start over
Data• Qualitative
o General observations about the system
• Quantitative o Comprising numbers obtained by various measurements of the system
• Theory o Unifying principle that explains a body of facts and the laws based on
themo Capable of suggesting new hypotheseso Can and do change
• Modelo We use many models to explain natural phenomenono When new evidence is found, the model changes!
• Scientific Laws o Summary of observed (measurable) behavioro A theory is an explanation of behavioro Law of Conservation of Masso Law of Conservation of Energy
A law summaries what happens; a theory (model) is an attempt to explain WHY it happens.
Classification of Matter
• Matter – anything that has mass and takes up space. o Includes things we can see and things we cannot seeo Everything in the universe has a “chemical” connection
Substances and Mixtures
• Substance – for of matter that has a definite (constant) composition and distinct properties.
• Mixture- combination of two or more substances in which the substances retain their distinct identities. o Homogeneous mixture – the composition of the mixture is the same
throughout. • Kool-aid• Sweet Tea
o Heterogeneous mixture – composition is not uniform throughout.• Milk• Smog• Chex mix• Sand
o Any mixture can be separated by physical means into pure components without changing the identities of the components.
Elements and Compounds
• Elements – substance that cannot be separated into simpler substances by chemical means. o Coppero Oxygeno Aluminum
• Compound – substance composed of atoms of two or more elements chemically united in fixed proportions. o Watero NaCl
States of Matter
States of Matter• Solid
o Definite volume and shape o Close together
• Liquido Definite volume but no definite shapeo Molecules moving faster than in solids
• Gaso Separated by distance
Physical & Chemical Properties of Matter
• Substances are identified by their properties as well as by their composition. o Physical property – can be measured and observed without changing
the composition or identity of a substance• Color• Melting/boiling points
o Chemical property – a chemical change must occur• Cannot be recovered • Hard-boil egg• Spoiled milk
Measurable Properties of Matter
• Extensive properties- depends on the amount of matter presento Masso Volumeo Length
• Intensive properties- does not depend on the amount of matter presento Coloro Densityo Temperatureo Boiling/melting points
Measurements• Macroscopic properties – can be determined
directly.
• Microscopic properties – are on the atomic or molecular scale and must be determined by an indirect method.
• A measured quantity is usually written as a number with an appropriate unit.
Example• If we say the distance form Katy to San Antonio is
300 by car along a particular path, this is meaningless.
• We have to specify that the distance is 300 km. • The same is true for chemistry; unit are essential
to stating the measurement correctly.
• Note: Points will always be deducted for lack of units.
SI Units• International System of Units (SI) • The table provided (next page) contains the
seven SI base units. • All other units of measurement can be derived
from these base units. • Like metric units, SI units are modified in decimal
fashion by a series of prefixes, shown in table 1.3.
SI Base Units
Prefixes Used with SI Units
Common Measurements in
Chemistry• Time• Mass• Volume• Density• Temperature
Mass and Weight• Video• The terms “mass” and “weight” are generally
used interchangeably, but they are different quantities.
• Mass is a measure of the amount of matter in an object.
• Weight is the force that gravity exerts on an object.
• The SI unit of mass is the kilogram (kg)
Volume• SI Unit of length is meters (m), and the SI-derived
unit for volume is the cubic meter (m3). • Chemist usually work with much smaller volumes,
such as the cubic centimeter (cm3) and the cubic decimeter (dm3)
1cm3 = (1 x 10-2 m)3 = 1 x 10-6 m3
1 dm3 = (1 x 10-1 m)3 = 1 x 10-3 m3
Another common unit of volume is the liter (L). A liter is the volume occupied by one cubic decimeter.
One liter of volume is equal to 1000 mL or 1000 cm3
Volume• 1 L = 1000 mL• 1L = 1000 cm3
• 1L = 1 dm3
• And one milliliter is equal to one cubic centimeter:
• 1mL = 1 cm3
Density• The equation for density is
Density = mass/volume
Example 1• A 74.8 g sample of mercury has a volume of 5.50
mL. Calculate the density of mercury.
Example 2• A piece of platinum metal with a density of 21.5
g/cm3 has a volume of 4.49 cm3. What is the mass?
Specific Gravity • aka Relative Density• Ratio of the density of a substance relative to the
density of water• Specific gravity is dimensionless
Temperature• Three temperature scales:• Fahrenheit (°F) (most common in US outside of
the lab)• Celsius (°C)• Kelvin (K) (SI unit base unit for temperature)• Kelvin is the absolute temperature scale
Temperature Conversions
• Convert from Fahrenheit to Celsius –
• Convert from Celsius to Fahrenheit –
• Convert from Celsius to Kelvin –K = °C + 273.15
• Convert form Kelvin to Celsius –°C = K – 273.15
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Practice• Convert 23°F to Celsius and then to Kelvin.
• Convert -98°C to Fahrenheit and then to Kelvin.
Scientific Notation• Chemist often work with numbers that are
extremely large or extremely small. • For example, 1 mole of any element contains
602,200,000,000,000,000,000,000 atoms• Or 1 hydrogen atom has a mass of
0.00000000000000000000000166g
• Writing in scientific notation makes numbers like these more manageable.
Scientific Notation• In scientific notation all numbers are written in the
form: • a x 10b
• (a times 10 to the power of b)
• For example:• 235,000,000,000 can be written as 2.35 x 1011 or• 0.000000657 can be written as 6.57 x 10-7
Writing Scientific Notation
• The guidelines for writing numbers in scientific notation is:o Count the number of places that the decimal point will have to move to
get one nonzero digit to the left side of the decimal place.o The number of places you moved the decimal point is the number used
as the exponent. • If you moved the decimal to the right, you make the exponent a
negative value.• If you moved the decimal to the left, you make the exponent a
positive number.
Significant Figures• When measurements are made for scientific purposes, we
include one number that is an estimate in all measurement. Although this may seem a bit unusual, but by including an estimated digit (which means that the value may contain an error, our measurement is actually more accurate than if we had just used the values known.
• The last digit is an estimate, and can vary from one person’s observation to another. In reporting measurements we keep all the digits that are known exactly, plus one digit that is an estimate and contains some error.
• This is called significant figures.
Significant Figures• If is also necessary to use measured quantities in
calculations. When this is done, it is necessary to know how many significant figures are in each number involved in the calculation.
• The answer must reflect the proper number of significant figures. Below are rules that are helpful in reporting the correct number of significant figures.
Nonzero Numbers• Nonzero Numbers Rule(s):
All nonzero numbers are significant in a measurement.
• For example: 3476.89 mm, consists of 6 significant figures
Numbers Containing Zeros
• Rules for determining significant figures get more complex when there are zeros in the number. The rules for numbers containing zeros are below:
• Zeros between nonzero numbers are significant. o Example: 509 kg, consists of 3 significant figures
• Leading zeros are zeros that precede all nonzero digits and are not significant.o Example: 0.00078 g, contains only 2 significant figures
• Trailing zeros at the end of a number are only significant if they contain a decimal point.o Example: 56.60 L, consists of 4 significant figures
• Zeros at the end of a number without a decimal point are ambiguous.o For example: 200 has only one significant figures, 200. Has three significant figures,
and 1.00 x 102 has three significant figures.
Practice• Determine the number of significant figures in
each of the following:• A student records a mass of 0.0009860 g in a
laboratory investigation. • 12.098 kg• 470 mm• 87 cm• 0.0034 g
Sig Figs in Multiplication and
Division• In multiplication and division, the significant
figures in the result is the same as the number in the least precise measurement used in the calculation. For example:
2.5 x 45.5 = 113.75• The first number has 2 sig figs and the second
has 3 sig fig, so the answer should have only 2 sig figs. The corrected answer is 110.
Sig Figs in Addition and Subtraction
• In addition and subtraction, the result has the same number of decimal places as the least precise measurement used in the calculation. For example:
10.2215.0
5.10230.322
• Correct answer: is to one decimal place and 30.3
Practice• Using the rules for applying significant figures to
calculations for the following:1. 97.481 + 6.2401 + 0.789112. 21.807 – 12.43 – 3.54093. 4.183 x 200.76 x (56.43 – 24.13)4. 56.09 / 2.9000
Rounding• Rules for Rounding• In a series of calculations, carry the extra digits
through to the final result, then round. • If the digit to be removed:
o Is less than 5, the preceding digit stays the same. For example, 1.23 rounds to 1.2.
o Is equal to or greater than 5, the preceding digit is increased by 1. For example, 2.47 rounds to 2.5.
Accuracy and Precision
• There are two aspects of the results generated in a laboratory investigation that describe the reliability of the measurements; accuracy and precision.
• Accuracy is how close are the experimental results to the true or real value for the quantity being measured. For example, if an oxygen gas sample were 60.0% pure and the lab reported back a value 60.1%, it would be considered accurate because it is close to the true value.
• Precision is how reproducible a measurement is if the same sample is measured multiple times. For example, if a sample of sodium oxide was measure three times as the results were, 32.0%, 31.99%, and 32.3%, those results are numerically close and the would be considered precise.
• In order to be accurate you must also be precise, but you can be precise without being accurate.
•
Practice• Label each of the following sets of data accurate,
precise inaccurate or imprecise. In each case, the true value of the measurement is 13.25.
• 10.22, 14.21, 13.24• 15.24, 15.21, 15.28• 13.24, 13.21, 13.23
Unit Conversion• Solving problems in chemistry often requires
converting from one unit of measurement to another by using a conversion factor.
• The best approach systematic method to convert units is by dimensional analysis.
Conversion Factor• Is an equality by which a quantity is multiplied to
convert from the original units to the quantity to the new units.
• For example: How many items are in a dozen?12 = 1 dozen or
dozen1
12Equivalence Statement
Converting from One Unit to Another – Dimensional
Analysis1. To convert from one unit to another, use the
equivalence statement that relates the two units.
2. Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).
3. Then multiply the quantity to be converted by the unit factor to give the quantity with the desired units.
Group Practice 1• How many grams are in 32 kilograms?
• How many centimeters are in 25.6 inches (1 inch equals 2.54 centimeters)?
• How many seconds are in 25 minutes?
Group Practice 2• How many centimeters are in 1.5 feet?
• A bar of gold has a mass of approximately 18.9 kg. Calculate the mass in pounds, if 1 pound = 454 g.
• Calculate the volume in liters of a 5.6 m3 container.
• • If a truck is traveling 24.6 km/hr, what is the
speed in ft/s? (1 km = 0.621 mile, 1 mile = 5,280 feet)