1. Chemistry: The Study of Change

2. Chemistry is the study of matter and the changes it undergoes
Matter is anything that occupies space and has mass.
Matter may exist as a (pure) substance?
1.4

3. What is a pure substance?
Group Discussion
What is a pure substance?

4. Pure Substance
A (pure) substance is a form of matter that has a definite composition and distinct properties.
Examples: water, ammonia, sucrose, gold, oxygen

5. soft drink, milk, salt water
A mixture is a combination of two or more substances in which the substances retain their distinct identities.
Homogenous mixture – composition of the mixture is the same throughout.
soft drink, milk, salt water
Heterogeneous mixture – composition is not uniform throughout.
Oil and water, wood,
iron filings in sand, sand in water
1.4

6. Physical means can be used to separate a mixture into its pure components.
distillation
magnet
1.4

7. 115 elements have been identified
An element is a substance that cannot be separated into simpler substances by chemical means.
115 elements have been identified
83 elements occur naturally on Earth
gold, aluminum, lead, oxygen, carbon
32 elements have been created by scientists
technetium, americium, seaborgium
1.4

8. A compound is a substance composed of atoms of two or more elements chemically united in fixed proportions.
Compounds can only be separated into their pure components (elements) by chemical means.
Water (H2O)
Glucose (C6H12O6)
Ammonia (NH3)
1.4

9. 1.4

10. Three States of Matter
1.5

11. hydrogen gas burns in oxygen gas to form water
Physical or Chemical?
A physical change does not alter the composition or identity of a substance.
ice melting
sugar dissolving
in water
A chemical change alters the composition or identity of the substance(s) involved.
hydrogen gas burns in oxygen gas to form water
1.6

12. TA p9

13. Matter - anything that occupies space and has mass.
mass – measure of the quantity of matter
SI unit of mass is the kilogram (kg)
1 kg = 1000 g = 1 x 103 g
weight – force that gravity exerts on an object
A 1 kg bar will weigh
2.2 lb on earth
0.4 lb on moon
1.7

14. Measurements

15. All measured quantities make known three pieces of information.
Measurements
All measured quantities make known three pieces of information.
The quantity or number
The unit
The uncertainty in the measurement.

16. Table 1.2 SI Base Units Base Quantity Name of Unit Symbol Length meter
Mass
kilogram kg
Time
second s
Temperature
kelvin K
Amount of substance
mole
mol
1.7

17. Derived SI Units Quantity Definition of Quantity SI unit
Area Length squared m2
Volume Length cubed m3
Density Mass per unit volume kg/m3
Force Mass times acceleration of object kg * m/s2
( =newton, N)
Pressure Force per unit area kg/(ms2)
( = pascal, Pa)
Energy Force times distance traveled kg * m2/s2
( = joule, J)
Speed Distance traveled per unit time m/s
Acceleration Speed changed per unit time m/s2

18. Table 1.3 Prefixes Used with SI Units
Symbol
Meaning
Tera- T
1012
Giga- G
109
Mega- M
106
Kilo- k
103
Deci- d
10-1
Centi- c
10-2
Milli- m
10-3
Micro-
10-6
Nano- n
10-9
Pico- p
10-12
1.7

19. Volume – is length cubed or L3
SI derived unit for volume is cubic meter (m3)
1 L = 1 dm3 (Definition)
1 mL = 1 x 10-3 L or 1L = 1 x 103 mL
1 dm3 = (10 cm)3 = 1000 cm3
1 L = 1000 mL = 1000 cm3 = 1 dm3
1 mL = 1 cm3
1.7

20. Density – SI derived unit for density is kg/m3
1 g/cm3 = 1 g/mL = 1000 kg/m3
density =
mass
volume
d = m V
A piece of platinum metal with a density of 21.5 g/cm3 has a volume of 4.49 cm3. What is its mass?
d = m V
m = d x V
= 21.5 g/cm3 x 4.49 cm3 = 96.5 g
1.7

21. K = 0C
273 K = 0 0C
373 K = 100 0C
0F = x 0C + 32 9 5
32 0F = 0 0C
212 0F = 100 0C
1.7

22. Convert 172.9 0F to degrees Celsius.
0F = x 0C + 32 9 5
0F – 32 = x 0C 9 5
x (0F – 32) = 0C 9 5
0C = x (0F – 32) 9 5
0C = x (172.9 – 32) = 78.3 9 5
1.7

23. Scientific Notation The number of atoms in 12 g of carbon:
602,200,000,000,000,000,000,000
6.022 x 1023
The mass of a single carbon atom in grams:
1.99 x 10-23
N x 10n
N is a number
between 1 and 10
n is a positive or
negative integer
1.8

24. Scientific Notation Addition or Subtraction 568.762 0.00000772
move decimal left
move decimal right
n > 0
n < 0
= x 102
= 7.72 x 10-6
Addition or Subtraction
Write each quantity with the same exponent n
Combine N1 and N2
The exponent, n, remains the same
4.31 x x 103 =
4.31 x x 104 =
4.70 x 104
1.8

25. Scientific Notation Multiplication Division
(4.0 x 10-5) x (7.0 x 103) =
(4.0 x 7.0) x (10-5+3) =
28 x 10-2 =
2.8 x 10-1
Multiply N1 and N2
Add exponents n1 and n2
Division
8.5 x 104 ÷ 5.0 x 109 =
(8.5 ÷ 5.0) x =
1.7 x 10-5
Divide N1 and N2
Subtract exponents n1 and n2
1.8

26. Chemskillbuilder
Register under:
Chipola College
Dr. Buffone
CHM1045

27. Significant Figures

28. SIGNIFICANT FIGURES
Except when all numbers are integers it is impossible to measure the exact value of a quantity.
The uncertainty in a measurement is indicated by the number of significant figures, which are the meaningful digits in a measured or calculated quantity. The last digit is understood to be uncertain by + or - 1.

29. Measurement Uncertainty
6 mL (+/- 1 mL)
6.0 mL (+/- 0.1 mL)
6.00 mL (+/ mL)
As a rule of thumb estimate 1 figure beyond the smallest subdivision on the scale.

30. The Number of Significant Figures in a Measurement Depends Upon the
Measuring Device
Fig 1.15A
As a rule of thumb estimate 1 figure beyond the smallest
subdivision on the scale.

31. Fig. 1.6

32. Significant Figures Any digit that is not zero is significant
1.234 kg significant figures
Zeros between nonzero digits are significant
606 m significant figures
Zeros to the left of the first nonzero digit are not significant
0.08 L significant figure
If a number is greater than 1, then all zeros to the right of the decimal point are significant
2.0 mg significant figures
If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant
g 3 significant figures (4.20 mg)
1.8

33. How many significant figures are in each of the following measurements?
24 mL
2 significant figures
3001 g
4 significant figures m3
3 significant figures
6.4 x 104 molecules
2 significant figures
560 kg
2 significant figures
1.8

34. Do this practice exercise with your group
For numbers that do not contain decimal points,
the trailing zeros may or may not be significant.
Use scientific notation to avoid ambiguity.
Example 1.3
478 cm b cm c m
d kg e x 1022 atoms f mL
Do this practice exercise with your group

35. Significant Figures Addition or Subtraction – use decimal places
The answer cannot have more digits to the right of the decimal
point than any of the original numbers.
89.332
1.1 +
90.432
one significant figure after decimal point
round off to 90.4
3.70
0.7867
two significant figures after decimal point
round off to 0.79
1.8

36. Significant Figures Multiplication or Division – use # of sig. figs.
The number of significant figures in the result is set by the original number that has the smallest number of significant figures
4.51 x =
= 16.5
3 sig figs
round to
3 sig figs
6.8 ÷ =
= 0.061
2 sig figs
round to
2 sig figs
1.8

37. Significant Figures Exact Numbers
Numbers from definitions or numbers of objects are considered
to have an infinite number of significant figures
The average of three measured lengths; 6.64, 6.68 and 6.70? 3
= = 6.67
= 7
Because 3 is an exact number
1.8

38. Example 1.4 Practice Exercise
L L
9.1 g – g
7.1 x 104 dm x x 102 dm
6.54 g / mL
7.55 x 104 m – 8.62 x 103 m

39. Accuracy – how close a measurement is to the true value
Precision – how close a set of measurements are to each other
accurate
&
precise
precise
but
not accurate
not accurate
&
not precise
1.8

40. Factor-Label Method of Solving Problems
Write down starting number with units.
Determine which unit conversion factor(s) are needed
Carry units through calculation
If all units cancel except for the desired unit(s), then the problem was solved correctly.
How many mL are in 1.63 L?
1 L = 1000 mL 1L
1000 mL
1.63 L x
= 1630 mL 1L
1000 mL
1.63 L x = L2 mL
1.9

41. The speed of sound in air is about 343 m/s
The speed of sound in air is about 343 m/s. What is this speed in miles per hour?
meters to miles
seconds to hours
1 mi = 1609 m
1 min = 60 s
1 hour = 60 min
343 m s x
1 mi
1609 m
60 s
1 min x
60 min
1 hour x
= 767 mi
hour
1.9

42. Common SI-English Equivalent Quantities
Quantity English to SI Equivalent
Length mile = 1.61 km
1 yard = m
1 foot (ft) = m
1 inch = 2.54 cm (exactly!)
Volume cubic foot = m3
1 gallon = dm3
1 quart = dm3
1 quart = cm3
1 fluid ounce = 29.6 cm3
Mass pound (lb) = kg
1 pound (lb) = g
1 ounce = g

43. Example 1.5 Practice exercise
A roll of Al foil has a mass of 1.07 kg. What is the
mass in pounds.
Example Practice exercise
The density of silver is 10.5 g/cm3. Convert this
to kg/m3.

44. Sample Problem
The volume of an irregularly shaped solid can be determined
from the volume of water it displaces. A graduated cylinder
contains mL water. When a small piece of Pyrite, an ore
of Iron, is submerged in the water, the volume increases to
315.8 mL. What is the volume of the piece of Pyrite in cm3
and in liters.
Vol (mL) = mL mL = 70.8 mL
Vol (cm3) = 70.8 mL x 1 cm3/ 1 mL = 70.8 cm3
Vol (liters) = 70.8 mL x 10 –3 L / mL = 7.08 x L

45. Density Density = mass/volume (an intensive quantity)
A small rectangular slab of lithium has a mass of
1.49 g and measures 2.09 cm by 1.11 cm
by 1.19 cm. Find its density in g/mL and lb/in3.
What is the volume of 15.3 g of Li?
What is the mass of 134 mL of Li?

46. Densities of Some Common Substances
Substance Physical State Density (g/cm3)
Hydrogen Gas
Oxygen Gas
Grain alcohol Liquid
Water Liquid
Table salt Solid
Aluminum Solid
Lead Solid
Gold Solid