1. Introduction to General Chemistry

2. What is Chemistry?
Chemistry is the study of properties of substances and how they react
Chemical substances are composed of matter
Matter is the physical material of the universe; anything with mass that occupies space is matter
Matter can take numerous forms
Most matter is formed by unique arrangements of elementary substances called elements

4. Elements, Compounds, and Molecules
An element can easily be defined as a substance that can not be broken down into simpler substances
Millions of different materials in the world, all comprised of some combination of only 118 elements
Similar to how the alphabet combines 26 letters to yield hundreds of thousands of words, elements bond in unique arrangements to give different molecules
Molecules agglomerate to yield compounds

5. Molecules Are Comprised of Uniquely Arranged Atoms
Molecules Are Comprised of Uniquely Arranged Atoms. Different Molecules Have Different Properties. H O H O

6. acetaldehyde (hangover)
Small Molecular Differences Can Yield Vastly Different
in Terms of Biological Interactions
acetaldehyde (hangover)
acetic acid
carbon dioxide
ethanol
BLINDNESS!!!
methanol

7. Small Molecular Differences Can Yield Vastly Different
in Terms of Biological Interactions
Relief of Morning Sickness
Severe Limb Defects

8. The Properties of Molecules Differ Vastly from those of the
Atoms That Comprise Them
Na (sodium metal)
Cl2 (chlorine gas)
Na+Cl-

9. Different atomic arrangements can change physical properties
Carbon (graphite vs diamond)

10. Spatial Dimensions of Compounds Can Alter Properties
Gold Nanoparticles
Bulk Gold
5 nm
50 nm

11. Spatial Dimensions of Compounds Can Alter Properties
2 nm
12 nm
CdSe quantum dots

12. Phases of Matter: Solids, Liquids and Gases
Atoms tightly bound
Fixed volume and shape (does not conform to container)
A chemical is denoted as solid by labeling it with (s)
S(s)

13. Phases of Matter: Solids, Liquids and Gases
Atoms less tightly bound than solids
Has a definite volume, but not definite shape (assumes the shape of its container)
Denoted by (L)
H2O (L)

14. Phases of Matter: Solids, Liquids and Gases
Free atoms
No shape, no definite volume
Can be expanded or compressed (like engine piston)
Denoted by (g) ; ex. O2 (g)

15. Qualitative and Quantitative Analysis
In chemistry, the scientific method is used to investigate scientific phenomena & acquire new knowledge
Empirical evidence is gathered which supports or refutes a hypothesis
Empirical evidence is either quantitative or qualitative
Quantitative data is numerical, and results can be measured
Qualitative data is NOT numerical, but consists of observations and descriptions

16. Quantitative and Qualitative Analysis
A + B C
Quantitative data
How much C is formed?
How efficient is the reaction?
What is the rate of the reaction?
Qualitative data
What color is it?
Is it solid, liquid, gas?
How does it smell?

17. Units Quantitative measurements are represented by a:
NUMBER and a UNIT
A unit is a standard against which a physical quantity is compared physical quantity
Temperature is measured in Co, Ko,or Fo
Currency is measured in $USD
Distance is measured in meters, miles, ft, etc.
Time is reported in seconds, minutes, hr, etc.
Internationally accepted system of measurements is called the SI unit system

18. SI Unit System: The Units of Physical Science

19. Greek Prefixes Prefixes indicate powers of 10
ex. k= 103; 5 kg = 5 x (103)g
GP1

20. Why Are Units Important? Example #1
In 1999, NASA lost the $125M Mars Orbiter System.
One group of engineers failed to communicate with another that their calculated values were in English units (feet, inches, pounds), and not SI units.
The satellite, which was intended to monitor weather patterns on Mars, descended too far into the atmosphere and melted.

21. Why Are Units Important? Example #2
In 1983, an Air Canada Plane ran out of fuel half way through its scheduled flight. Why?
Airline workers improperly converted between liters and gallons.
Luckily, no one died.

22. Why Are Units Important? Example #3
A case was reported in which a nurse administered 0.5 g of a sedative to a patient.
The patient died soon after
The patient should have only received 0.5 grains (1 grain = 0.065g) but the units were not listed. That was the equivalent of 8 doses!!

23. Derived SI Units: VOLUME
Many measured properties have units that are combinations of the fundamental SI units
Volume: defines the quantity of space an object occupies; or the capacity of fluid a container can hold
expressed in units of (length)3 or Liters (L)
1 L is equal to the volume of fluid that a cube which is 10 cm on each side can hold
10 cm
V = (10 cm)3 = 1000 cm3
1 L = 1000 cm3
1000 mL = 1000 cm3
mL = cm3

24. THE DENSITY OF WATER IS 𝟏 𝒈 𝒄𝒎 𝟑 𝒐𝒓 𝟏 𝒈 𝒎𝑳
Derived SI Units: DENSITY
All matter has mass, and must therefore occupy space. Density correlates the mass of a substance to the volume of space it occupies.
Density = mass per unit volume (mass/volume). Different materials have different densities.
A container filled with bricks does not have the same mass as an equivalent volume of feathers!
GP2
THE DENSITY OF WATER IS 𝟏 𝒈 𝒄𝒎 𝟑 𝒐𝒓 𝟏 𝒈 𝒎𝑳

25. Derived SI Units: ENERGY
What is Energy?
Energy is defined as the capacity to perform “work”
How do we define work?
Work is defined as the action of applying a force acting over some distance
Ex. Pushing an object along a rough surface
In SI units, we use the unit Joule (J) to represent energy.
𝐽= 𝑘𝑔 𝑚 2 𝑠 2

26. Temperature
Temperature: a measure of the tendency of a substance to lose or absorb heat. Temperature and heat are not the same.
Heat always flows from bodies of higher temperature to those of lower temperature
An active stove top is ‘hot’ because the surface is at a much higher temperature than your hand, so heat flows rapidly from the stove to your hand
Ice feels ‘cold’ because it is at a lower temperature than your body, so heat flows from your body to the ice, causing it to melt
Explain wind chill?

27. Temperature
When performing calculations in chemistry, temperature must always be converted to Kelvin (oK) units (unless otherwise stated).
The lowest possible temperature that can ever be reached is 0oK, or absolute zero. At this temperature, all molecular motion stops.
To convert temperatures to celcius and Kelvin:
oK : oC
oC: 5/9 (oF – 32)
GP3

28. Accuracy and Precision
Accuracy defines how close to the correct answer you are. Precision defines how repeatable your result is. Ideally, data should be both accurate and precise, but it may be one or the other, or neither.
Accurate, but not precise. Reached the target, but could not reproduce the result.
Precise, but not accurate. Did not reach the target, but result was reproduced.
Accurate and precise.
Reached the target and the data was reproduced.

29. Measuring Accuracy Accuracy is calculated by percentage error (%E)
We take the absolute value because you can’t have negative error.
GP 4

30. Measuring Precision: Significant Figures
Precision is indicated by the number of significant figures. Significant figures are those digits required to convey a result.
There are two types of numbers: exact and inexact
Exact numbers have defined values and possess an infinite number of significant figures because there is no limit of confidence:
* There are exactly 12 eggs in a dozen
* There are exactly 24 hours in a day
* There are exactly 1000 grams in a kilogram
Inexact number are obtained from measurement. Any number that is measured has error because:
Limitations in equipment
Human error

31. Measuring Precision: Significant Figures
Example: Some laboratory balances are precise to the nearest cg (.01g). This is the limit of confidence. The measured mass shown in the figure is g.
The value has 5 significant figures, with the hundredths place (9) being the uncertain digit. Thus, the (9) is estimated, while the other numbers are known.
It would properly reported as ±.01g
- The actual mass could be anywhere between … g and … g. The balance is limited to two decimal places, so it rounds up or down. We use ± to include all possibilities.

32. Determining the Number of Significant Figures In a Result
All non-zeros and zeros between non-zeros are significant
457 (3) ; 2.5 (2) ; 101 (3) ; 1005 (4)
Zeros at the beginning of a number aren’t significant. They only serve to position the decimal.
.02 (1) ; (1) ; (4)
For any number with a decimal, zeros to the right of the decimal are significant
2.200 (4) ; 3.0 (2)

33. Determining the Number of Significant Figures In a Result
Zeros at the end of an integer may or may not be significant
130 (2 or 3), (1, 2, 3, or 4)
This is based on scientific notation
130 can be written as:
1.3 x  2 sig figs
1.30 x 102  3 sig figs
If we convert 1000 to scientific notation, it can be written as:
1 x 103  1 sig fig
1.0 x  2 sig figs
1.00 x  3 sig figs
1.000 x  4 sig figs
*Numbers that must be treated as significant CAN NOT disappear in scientific notation

34. Calculations Involving Significant Figures
You can not get exact results using inexact numbers
Multiplication and division
Result can only have as many significant figures as the least precise number
𝑐𝑚 𝑥 𝟓.𝟖𝟐 𝑐𝑚= 𝑐 𝑚 2 =36.2 𝑐 𝑚 2
(3 s.f.)
𝑚 0.𝟗𝟖 𝑠 = 𝑚 𝑠 =110 𝑚 𝑠 𝑜𝑟 1.1 𝑥 𝑚 𝑠
(2 s.f.)
𝑘𝑔 𝑥 𝟒 𝑚 𝑠 2 = 𝑘𝑔 𝑚 𝑠 2 = 𝑘𝑔 𝑚 𝑠 2 𝑜𝑟 2 𝑥 𝑘𝑔 𝑚 𝑠 2
(1 s.f.)

35. Calculations Involving Significant Figures
Addition and Subtraction
Result must have as many digits to the right of the decimal as the least precise number
20.4
1.322 83
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36. Limit of certainty is the ones place
Mixed Operations
H= cm
W = .40 cm
L = cm
Volume of rectangle ?
V = LWH = 124 cm3 ≈ 120 cm3 or 1.2 x 102 cm3
Surface area (SA = 2WH + 2LH + 2LW) ?
note: constants in an equation are exact numbers
=2 4.0𝑐 𝑚 𝑐 𝑚 2 +2(1𝟐.4𝑐 𝑚 2 )
Limit of certainty is the ones place
=8.0𝑐 𝑚 𝑐 𝑚 2 +2𝟒.8𝑐 𝑚 2
=65𝟐.8𝑐 𝑚 2 =653𝑐 𝑚 2

37. Dimensional Analysis
Dimensional analysis is an algebraic method used to convert between different units
Conversion factors are required
Conversion factors are exact numbers which are equalities between one unit set and another.
For example, we can convert between inches and feet. The conversion factor can be written as:
In other words, there are 12 inches per 1 foot, or 1 foot per 12 inches.

38. Dimensional Analysis conversion factor (s)
Example. How many feet are there in 56 inches?
Our given unit of length is inches
Our desired unit of length is feet
We will use a conversion factor that equates inches and feet to obtain units of feet. The conversion factor must be arranged such that the desired units are ‘on top’
𝟓𝟔 𝑖𝑛𝑐ℎ𝑒𝑠 𝑥 1 𝑓𝑜𝑜𝑡 12 𝑖𝑛𝑐ℎ𝑒𝑠 = 𝑓𝑡
4.7 ft
GP 6

39. High Order Exponent Unit Conversion (e.g. Cubic Units)
As we previously learned, the units of volume can be expressed as cubic lengths, or as capacities. When converting between the two, it may be necessary to cube the conversion factor
Ex. How many mL of water can be contained in a cubic container that is 1 m3 3
Must use this equivalence to convert between cubic length to capacity
1 𝑚 3 𝑥 𝑐𝑚 1 0 −2 𝑚 𝑥 𝒎𝑳 𝒄 𝒎 𝟑
Cube this conversion factor
=1 𝑚 3 𝑥 𝒄𝒎 𝟑 𝟏 𝟎 −𝟔 𝒎 𝟑 𝑥 𝑚𝐿 𝑐 𝑚 3
=𝟏 𝒙 𝟏 𝟎 𝟔 𝒎𝑳
GP 7