1. General Chemistry Tonya Patterson
Chemical Foundation
General Chemistry
Tonya Patterson

2. Chemistry Is the study of matter and the changes it undergoes
Often referred to as the central science
Why study chemistry?

3. Scientific Method A systematic approach to research
Used by all sciences
Process
Identify problem
Research and/or observations
Form a hypothesis (tentative explanation or prediction of experimental observations)
Experiment/Testing
Analysis of data
Draw conclusion
If hypothesis is correct – finished
If hypothesis is incorrect – start over

4. Data Qualitative Quantitative General observations about the system
Comprising numbers obtained by various measurements of the system

5. Theory Model Scientific Laws
Unifying principle that explains a body of facts and the laws based on them
Capable of suggesting new hypotheses
Can and do change
Model
We use many models to explain natural phenomenon
When new evidence is found, the model changes!
Scientific Laws
Summary of observed (measurable) behavior
A theory is an explanation of behavior
Law of Conservation of Mass
Law of Conservation of Energy

6. A law summaries what happens; a theory (model) is an attempt to explain WHY it happens.

7. Classification of Matter
Matter – anything that has mass and takes up space.
Includes things we can see and things we cannot see
Everything in the universe has a “chemical” connection

8. 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.
Homogeneous mixture – the composition of the mixture is the same throughout.
Kool-aid
Sweet Tea
Heterogeneous mixture – composition is not uniform throughout.
Milk
Smog
Chex mix
Sand
Any mixture can be separated by physical means into pure components without changing the identities of the components.

9. Elements and Compounds
Elements – substance that cannot be separated into simpler substances by chemical means.
Copper
Oxygen
Aluminum
Compound – substance composed of atoms of two or more elements chemically united in fixed proportions.
Water
NaCl

12. States of Matter

14. States of Matter Solid Liquid Gas Definite volume and shape
Close together
Liquid
Definite volume but no definite shape
Molecules moving faster than in solids
Gas
Separated by distance

15. Physical & Chemical Properties of Matter
Substances are identified by their properties as well as by their composition.
Physical property – can be measured and observed without changing the composition or identity of a substance
Color
Melting/boiling points
Chemical property – a chemical change must occur
Cannot be recovered
Hard-boil egg
Spoiled milk

16. Measurable Properties of Matter
Extensive properties- depends on the amount of matter present
Mass
Volume
Length
Intensive properties- does not depend on the amount of matter present
Color
Density
Temperature
Boiling/melting points

17. 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.

18. 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.

19. 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.

20. SI Base Units

21. Prefixes Used with SI Units

22. Common Measurements in Chemistry
Time
Mass
Volume
Density
Temperature

23. 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)

24. 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 m)3 = 1 x 10-6 m3
1 dm3 = (1 x 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

25. Volume
1 L = 1000 mL
1L = 1000 cm3
1L = 1 dm3
And one milliliter is equal to one cubic centimeter:
1mL = 1 cm3

27. Density
The equation for density is
Density = mass/volume

28. Example 1
A 74.8 g sample of mercury has a volume of 5.50 mL. Calculate the density of mercury.

29. 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?

30. Specific Gravity aka Relative Density
Ratio of the density of a substance relative to the density of water
Specific gravity is dimensionless

31. 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

33. Temperature Conversions
Convert from Fahrenheit to Celsius –
Convert from Celsius to Fahrenheit –
Convert from Celsius to Kelvin –
K = °C
Convert form Kelvin to Celsius –
°C = K –

34. Practice Convert 23°F to Celsius and then to Kelvin.
Convert -98°C to Fahrenheit and then to Kelvin.

35. 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 g
Writing in scientific notation makes numbers like these more manageable.

36. 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
can be written as 6.57 x 10-7

37. Writing Scientific Notation
The guidelines for writing numbers in scientific notation is:
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.
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.

38. 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.

40. 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.

41. Nonzero Numbers Nonzero Numbers Rule(s):
All nonzero numbers are significant in a measurement.
For example: mm, consists of 6 significant figures

42. 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.
Example: 509 kg, consists of 3 significant figures
Leading zeros are zeros that precede all nonzero digits and are not significant.
Example: g, contains only 2 significant figures
Trailing zeros at the end of a number are only significant if they contain a decimal point.
Example: L, consists of 4 significant figures
Zeros at the end of a number without a decimal point are ambiguous.
For example: 200 has only one significant figures, 200. Has three significant figures, and 1.00 x 102 has three significant figures.

43. Practice
Determine the number of significant figures in each of the following:
A student records a mass of g in a laboratory investigation. kg
470 mm
87 cm g

44. 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 =
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.

45. 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:
10.22
15.0
5.102
30.322
Correct answer: is to one decimal place and 30.3

46. Practice
Using the rules for applying significant figures to calculations for the following:
– –
x x (56.43 – 24.13) /

47. 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:
Is less than 5, the preceding digit stays the same. For example, 1.23 rounds to 1.2.
Is equal to or greater than 5, the preceding digit is increased by 1. For example, 2.47 rounds to 2.5.

48. 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.

50. In order to be accurate you must also be precise, but you can be precise without being accurate.

51. Practice
Label each of the following sets of data accurate, precise inaccurate or imprecise. In each case, the true value of the measurement is
10.22, 14.21, 13.24
15.24, 15.21, 15.28
13.24, 13.21, 13.23

52. 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.

53. 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
Equivalence Statement

54. Converting from One Unit to Another – Dimensional Analysis
To convert from one unit to another, use the equivalence statement that relates the two units.
Derive the appropriate unit factor by looking at the direction of the required change (to cancel the unwanted units).
Then multiply the quantity to be converted by the unit factor to give the quantity with the desired units.

55. 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?

56. 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 = mile, 1 mile = 5,280 feet)