Introduction to Chemistry

1
Introduction to Chemistry
2
The world is full of obvious things. Which
nobody by any chance ever observe.-- Sherlock
Holmes

• Careful observation is the foundation of
  chemistry as an experimental science, leading us
  to question what we have observed how, what
  why?
• The answers to these questions are sought in
  experiments, which may be described as
  observations made under controlled conditions
• Observation and experimentation the twin
  pillars of the scientific method

3
Scientific Method

• The scientific way of knowing often called the
  scientific method is sometimes presented as a
  rigid sequence of events
• It is not however a rigid path it is a process
  of discovery!
• Discovery begins when we make observations and
  then try to understand what we have observed by
  asking key questions and proposing possible
  answers
• This process of discovery begins as we design and
  conduct experiments to test whether our answers
  to these questions are valid!!

4
What are the steps of the Scientific Method?

1. Make an observation(s)
2. Propose a hypothesis
3. Design and conduct a controlled experiment
4. Analyze the results
5. Form conclusions

5
Controlled Experiment

• Experiments should be designed so that the
  effects of different variables on the behavior of
  a substance can be studied independently
• A controlled experiment is when only one variable
  at a time is changed
• There are two types of variables in an
  experiment
• Independent variable the one that is
  deliberately changed
• Dependent variable the thing that changes
  because of the independent variable

6
An Example

• Students were trying to determine if the amount
  of a Sodium chloride added to calcium carbonate
  effected the amount of heat given off
• Give 3 possible variables for this experiment
• Give the independent and dependent variables

7
What is chemistry?

• Chemistry is the study and investigation of the
  structure, composition and property of matter and
  the changes it undergoes
• The properties of materials are always related to
  their structure
• Hence, structure determines properties

8
Measuring and Calculating in Science

• Chemistry is a quantitative science because it
  involves measuring and calculating
• A measurement must have a number and a scale
  (called a unit) to be meaningful
• It can also be a qualitative science because it
  can involve describing what is happening in a
  reaction

9
What makes a measurement?

• In order to make measurements, we must meet three
  requirements
• Know what we are trying to measure
• Have a standard with which to compare whatever we
  are measuring
• Have a method for making comparisons

10
Exact Numbers

• A number that is the result from a definition or
  an exact count
• For example there are 12 apples or p 3.14
• All are significant
• Do not limit the number of sig figs

11
Uncertainty in Measurement

• A measurement always has some degree of
  uncertainty
• The amount of uncertainty depends on the
  precision of the measuring device
• In science it is customary to report a
  measurement by recording all certain digits plus
  the first uncertain (estimated) digit these
  numbers are called the significant figures of a
  measurement

12
Estimating Uncertainty in a measurement

• Remember all measurements are a result of known
  values and one estimated number
• When finding the uncertainty in a measurement, we
  look at the estimated number
• For example 0.023
• The 3 is the estimated number
• It is in the 1000th place
• The estimated uncertainty is written 0.001
• This measurement has a very small uncertainty

13
Rules for Counting Significant Figures

• Nonzero integers are always significant figures
• Zeros there are three classes of zeros
• Leading zeros precede all the nonzero digits and
  are not significant figures
• Captive zeros are between nonzero digits and
  always count as significant figures
• Trailing zeros are at the right end of the number
  and are only significant if the number contains a
  decimal point

14
Try a few

• Tell how many sig figs are in each measurement
  and tell the uncertainty in each measurement
• 1508 cm
• 300.0 ft
• 20.003 g
• 0.00705 L

15
Bell Ringer

• Tell the number of sig fig
• 94300
• 0.000400670
• 100000.
• 56.00
• Tell the uncertainty of each measurement
• 3.45
• 6.0
• 12
• 4.725

16
Bell Ringer

• Each of the following are statements from
  different labs tell if they are quantitative or
  qualitative
• Bubbling
• Heat given off
• 23.6 cm wide
• A strong odor
• pH of 5.4
• 273 K

17
Sig figs in Mathematical Operations

• To this point we have learned to count the
  significant figures in a given number, but we
  must also consider how uncertainty accumulates as
  calculations are carried out

18
Rules for Sig figs in Mathematical Operations

• For multiplication or division the number of sig
  figs in the answer is the same as the number in
  the least precise measurement used in the
  calculation
• 4.56 x 1.4 6.38 ? 6.4 (correct answer)
• For addition or subtraction the number of sig
  figs in the answer has the same number of decimal
  places as the least precise measurement used in
  the calculation
• 12.11 18.0 1.013 31.123 ? 31.1 (correct
  answer)

19
Bell Ringer

• Carry out the following mathematical operations
  and give each result with the correct number of
  significant figures
• 1.05 x 10-3 / 6.135
• 21 13.8
• 20 X 23.00
• 14.75 34.25

20
Bell Ringer

• The actual length of a certain plank is 26.782
  cm. Which of the following measurements is the
  most accurate? Are the measurements precise?
• 26.5 cm
• 26.8 cm
• 26.202 cm
• 26.98 cm

21
Rules for Rounding

• In most calculations you will need to round
  numbers to obtain the correct number of sig figs
• When rounding, use only the first number to the
  right of the last significant figure
• 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, then the preceding digit stays
  the same
• 1.33 ? 1.3
• Is equal to or greater than 5, the preceding
  digit is increased by 1
• 1.36 ? 1.4

22
Precision and Accuracy

• Two terms often used to describe the reliability
  of measurements are precision and accuracy
• Precision the degree of agreement among several
  measurements of the same quantity. It also is
  known as the degree of reproducibility of the
  measurement
• Accuracy the agreement of a particular value
  with the true value

23
Bell Ringer

• Decide if the following lab data is accurate or
  precise or both
• 13.2 mL, 13.3 mL, 13.1 mL, 13.2 mL
• The actual value is 13.0 mL
• Make each of the following have 3 sig figs
• 34098
• 0.0003219
• 7154
• 76.78

24
Types of errors

• There are two types of errors in measurements
• Random error (indeterminate error) a
  measurement that has an equal probability of
  being high or low. This type of error occurs in
  estimating the value of the last digit of a
  measurement
• Systemic error (determinate error) occurs in
  the same directions each time. The measurement
  is either always too high or too low

25
In groups

1. There are 365 days/year, 24 hours/day, 12
   months/year and 60 minutes/hr. Use this data to
   determine how many minutes are in a month.
2. Now use the following data to calculate the
   number of minutes in a month 24 hours/day, 60
   minutes/hour, 7 days/week, and 4 weeks/month.
3. Why are these answers different? Which, if any,
   is more correct and why?

26
Dimensional Analysis

• It is often necessary to convert a given result
  from one system of units to another
• The best way to do this is by a method called
  unit factor method OR dimensional analysis

27
Converting from One Unit to Another

• 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
  unwanted units)
• Multiply the quantity to be converted by the unit
  factor to give the quantity with desired units

28
Bell Ringer

• A marathon race is 26 miles and 385 yards.
• What is the distance in rods
• What is the distance in meters
• What is the distance in furlongs?
• 5.5 yards 1 rod
• 40 rods 1 furlong
• 8 furlongs 1 mile
• 1 meter 39.37 inches
• 1 yard 36 inches

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What if there is more than one unit present?

• When more than one unit is present, decide which
  unit you want to convert first
• Convert it first
• Then convert the second unit
• do not get confused!!!
• EX
• How fast is a car going 35 miles/hour going in
  yards/second?
• 1 mile 1760 yards1 hour 60 minutes 1 minute
  60 seconds

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A few problems

1. How many doughnuts can one purchase for 123 if
   doughnuts cost 3.25/doz?
2. Convert 9.85 L to gal. 1.06 qt 1.00 L and 4 qt
   1 gal
3. A certain size of nail cost 1.25/lb. What is
   the cost of 3.25 kg of these nails? 1kg 2.2 lb

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Metric System Review

• Scientists recognized that long ago a standard
  system of units had to be adopted if measurements
  were to be useful
• The system agreed upon in 1960 was the
  International System or le Systeme International
  (SI system)
• The SI system is based on the metric system and
  units derived from the metric system
• Because fundamental units are not always
  convenient, the SI system employs prefixes to
  change the size of the unit

32
The Fundamental SI Units
Physical Quantity Name of Unit Abbreviation
Mass gram g
Length Meter m
Time Second s
Temperature Kelvin K
Amount of Substance Mole mol
Electric current Ampere A
Luminous intensity Candela cd
33
Derived units

• Many SI units are combinations of quantities
• These units are produced by multiplying or
  dividing standard units

34
Derived SI units
Quantity Name of Unit Abbreviation Derivation
Area (A) square meter m2 Length x width
Volume (V) Cubic meter m3 Length x width x height
Density (D) Kilograms per cubic meter kg/m3 Mass / volume
Molar mass (M) Kilograms per mole kg/mol Mass / amt of substance
Concentration (c) Moles per liter M Amt of substance/ volume
Molar volume (Vm) Cubic meters per mole m3/mol Volume / amt of substance
Energy (E) joule J Force x length
35
Dimensional Analysis with metric units

• When converting with metric, always use that
  value of the unit as compared to the base unit
• Convert 35.4 mm to m
• Convert 2327.9 cg to kg
• How many grams are in 53.24 dg?

36
Bell Ringer

• Why do we use the metric system?
• Convert 35.4 mm to m
• Convert 2327.9 cg to kg
• How many grams are in 53.24 dg?
• Convert the following
• How many inches are in 3.0 meters?
• A baby weighs 8.5 lbs. How many grams is that?
• How many gallons of Coke would you drink if you
  drank entire 2 liter?

37

• Science fiction often uses nautical analogies to
  describe space travel. If the starship U.S.S.
  Enterprise is traveling at warp factor 1.71, what
  is its speed in knots?
• Warp 1.71 5.00 times the speed of light
• The speed of light 3.00 x 108 m/s
• 1 knot 2000 yd/hr

38
Mass

• The measure of the resistance of an object to a
  change in its state of motion OR the amount of
  stuff in an object
• A scale is used to mass an object

39
Mass vs. weight

• An important point concerning measurements is the
  relationship between mass and weight
• Weight is the force gravity exerts on mass,
  therefore weight varies with the strength of the
  gravitational field
• Therefore if you went to the moon your weight
  would change but not your mass
• Many times the terms mass and weight are
  sometimes used interchangeably, although this is
  incorrect!

40
Volume

• The derived SI unit of volume is cubic meters
  (m3)
• Many times this unit is way too large to be a
  practical way of expressing volume in a chemistry
  lab
• Instead, a smaller unit cubic centimeters (cm3)
  is used
• When dealing with the volumes of liquids and
  gases, the non-SI unit liter (L) is often used
• Again the liter is often too large so the unit
  milliliter (mL) is used
• This means 1 cm3 1mL

41
Review

• Round the following to 3 sig figs
• 96747210
• 91
• 0.0006589
• How many sig figs are in each in 1?
• What is the uncertainty of each measurement?
• 34.09
• 6.0222
• 12
• What is the difference between precision and
  accuracy?

42
Review

• Convert the following
• How many grams are in 548.9 mg?
• How many feet are in 34.2 m?
• How many liters are in 2 gallon and 3.4 quarts?

43
It can be tricky with volume conversions

1. How many mL are in 14.65 kL?
2. How many L are in 48.6 cm3?
3. How many dm3 are in 29100 mL?

44
Bell Ringer

1. A piece of metal has the mass of 3.45 kg. What
   is its mass in g?
2. A container has 2.3 L of gas in it? What is its
   volume in mL?
3. A container has 750.00 mL of liquid in it. What
   is its volume in m3?

45
What is Temperature?

• A measure of the AVERAGE kinetic energy
• When looking at the different temperature scales,
  all are talking about the same height of mercury

46
Temperature Conversions

• There are three systems used to measure
  temperature
• Degrees Fahrenheit (F)
• Degrees Celsius (C)
• Kelvin (K)
• Each has a different way of converting between
  the values

47
How the equation for F to C was derived

• Notice -- 0C 32F and 100C 212F
• If we subtract these values then
• 100C 180F
• Find the value of 1C
• 1C (180/100) F
• 1C 9/5 F

48
Converting

• 1. Converting from C to Kelvin
• TC TK 273.15
• TK TC 273.15
• Converting from C to F
• TF TC x 9F 32F
• 5C

49
More Converting

• Converting F to C
• TC (TF - 32F)5C
• 9F

50
Try these

• 1. Normal body temperature is 98.6F. Convert
  this to the Celsius and Kelvin scales.
• 2. Liquid nitrogen, which is often used as a
  coolant for low-temperature experiments has a
  boiling point of 77 K. What is this temperature
  on the Fahrenheit scale?

51
Density

• Ratio of mass to volume
• D m/V
• Unit usually g/mL or g/cm3
• Useful for predicting mass
• Does not depend on the amount of material in a
  compound

52
Density

• A characteristic physical property of a substance
• It does not depend on the size of the sample
  because as the samples mass increases, its
  volume increases proportionally
• Density varies with temperature V ? T? D?

53
Bell Ringer

• What is the density of an object that has a mass
  of 14 g and a volume of 2 ml?
• What is the mass of an object with a volume of 2
  cm3 and a density of 1.5 g/mL?
• What is the volume of an object with the mass of
  20 kg and a density of 2.5 g/mL?

54
Bell Ringer

• An empty container weighs 121.3 g. Filled with
  carbon tetrachloride (density 1.53 g/cm3) the
  container weighs 283.2 g. What is the volume of
  the container?
• A 55.0 gal drum weighs 75 lbs when empty. What
  will the total mass be when filled with ethanol?
• density of ethanol 0.789 g/cm3
• 1 gal 3.78 L
• 1 lb 454 g

55

• In the opening scenes of the movie Raiders of the
  Lost Ark, Indiana Jones tries to remove a gold
  idol from a booby-trapped pedestal. He replaces
  the idol with a bag of sand of approximately the
  same volume. (density of gold 19.32 g/mL
  density or sand 2 g/mL)
• Did he have a reasonable chance of not activating
  the mass sensitive bobby-trap?
• In a later scene he and an unscrupulous guide
  play catch with the idol. Assume the volume of
  the idol is 1.0 L. If it were solid gold, what
  mass would the idol have?
• Is playing catch plausible?

56
Its all that is Matter

• Matter is anything that has mass and takes up
  space
• All matter, regardless of form, has some
  properties in common
• Volume the amount of 3-D space an object
  occupies
• Mass a measure of the amount of matter in an
  object

57
Basic Building Blocks of Matter

• The most fundamental parts of matter are atoms
  and molecules, which make up elements and
  compounds
• Atom the smallest unit of an element that has
  all the properties of that element
• Element a pure substance made of only one kind
  of atom
• Compound a substance that is made from the
  atoms of two or more elements that are chemically
  bonded
• Molecule the smallest unit of an element or
  compound that retains all the properties of that
  element or compound

58
Classification of Matter

• Matter is classified according to how it is
  organized
• Matter is complex and has different levels of
  organization
• Mixtures
• Pure substances

59
Mixtures

• Most of the matter around us consists of mixtures
  of pure substances
• Mixture consists of materials with variable
  compositions
• Two types of mixtures
• Homogeneous
• heterogeneous

60
Homogeneous mixture

• Homogeneous mixture having visibly (to the
  naked eye) indistinguishable parts
• Has one phase present
• It also called a solution
• There are two parts of a solution
• 1. the solute the part that is dissolved
• 2. the solvent the part that does the
  dissolving
• WATER IS THE UNIVERSAL SOLVENT!!

61
Alloy

• A homogeneous mixture of metallic elements with
  one solid phase

62
Alloy ProblemsConversion of Total Mass to Mass
of a Component

• Manganese steel is very strong and finds use as
  railroad rails. It is composed of 86.0 iron,
  13.0 manganese, 1.0 carbon. What is the mass
  of each of the three elements in a 254-kg sample
  of manganese steel?

63
One more try

• A sample of brass is composed of 72 copper and
  the remainder zinc. What mass of brass can be
  made from 25-kg of zinc?

64
Examples of solutions

• Air a gaseous mixture of various gases
• Brass a solid mixture of various metals
• Iced Tea a liquid mixture of various materials

65
Heterogeneous mixtures

• Heterogeneous mixture having visibly
  distinguishable parts
• Can usually be separated into two or more
  homogeneous mixtures or pure substances
• Has 2 or more phases present

66
Examples of heterogeneous mixtures

• Sand and water
• Iced tea and ice cubes
• Pepperoni pizza

67
Pure substances

• Mixtures can be separated into pure substances by
  physical methods
• Pure substance contains materials with a
  constant composition, such as compounds and
  elements
• Have a definite composition and definite
  unchanging properties (both chemical and physical)

68
Classify the following as a mixture or a pure
substance

• Salt water
• Smog
• Water
• 10 karat gold
• Sugar
• Diamond
• Coffee
• Chex mix
• Chef Salad

69
Properties

• We use two type of properties to describe people
• Physical properties what we look like
• Emotional/Personality how we interact with
  other people
• We also use properties to describe matter
• Physical properties what the matter looks like
• Chemical properties how matter interacts with
  other matter

70
Physical Properties

• A physical property is a characteristic that can
  be observed or measured without changing the
  identity of the substance
• It describes a substance
• Color
• State at room temperature
• Melting point
• Boiling point
• Density
• Specific gravity

71
Physical Changes

• A physical change is a change in a substance that
  does not involve a change in the identity of the
  substance
• A change of state is a physical change of a
  substance from one state to another
• Solid
• Liquid
• Gas
• Plasma

72
Chemical Properties

• A chemical property relates to a substances
  ability to undergo changes that transform it into
  different substances
• Chemical properties are easiest to see when
  substances react to form new substances which has
  different properties than the original substances

73
Chemical change

• A chemical change is a change in which one or
  more substances are converted into different
  substances
• It can also be called a chemical reaction
• Reactants original substances
• Products the new substances formed
• Chemical changes or reactions form products whose
  properties differ greatly from the reactants
• Chemical changes do not affect the total amount
  of matter (and also the mass) present before and
  after a reaction

74
Methods for Separating Mixtures

• There are many different ways to separate
  mixtures based on their physical properties
• Distillation uses the volatility of the
  components
• Filtration used when a mixture consists of a
  solid and a liquid
• Chromatography separates based on speed of
  movement of the components of the mixture
• Solubility uses the amount of solute that
  dissolves in water at a given temperature
• Density
• Melting point

75
Energy Transfers

• Physical and Chemical changes are always
  accompanied by energy changes
• One way energy can be transferred is through a
  temperature difference and is called heat (q)
• The quantitative measurements in energy changes
  are expressed in joules (J)

76
Energy

• Energy is an important concept in chemistry and
  is a property of all matter
• The ability to do work
• All objects possess energy
• Forms of energy
• 1. Chemical energy released as heat energy
• 2. Nuclear energy
• 3. Mechanical energy
• 4. Electrical Energy
• 5. Light energy
• 6. Radiant energy (not a property of objects)
  the transfer of
• 
  energy through empty space

77
Categories of Energy

• There are two main categories of all energy we
  see
• Potential energy the energy stored in the
  object
• Kinetic energy the energy due to motion

78
Laws of Matter and Energy

• Law of Conservation of Matter
• Law of conservation of Energy
• Law of Conservation of Mass-Energy
• Mass and energy are interconvertible
• Matter can change forms
• Energy can change forms
• Mass can change into energy
• However, never can mass or energy be created or
  destroyed

79
Review Questions

• What is the Law of Conservation of Energy?
• What are the two main types of energy?
• Classify either as a chemical change or a
  physical change
• Digestion of a candy bar
• Melting of ice
• Formation of clouds
• Growth of plants
• Fading of dye in a cloth

80
A little more about the Properties of Matter

• Every substance, whether it is an element or
  compound, has characteristics properties
• These properties are used to distinguish between
  substances and to separate them

81
Properties of Matter cont

• Properties can also be characteristics of an
  entire group as with metals
• Properties can help identify unknown substances
  notice the plural not just one property can
  identify a substance
• Properties are either intensive or extensive
• Intensive does not depend on the amount of
  matter present
• Extensive does depend on the amount of matter
  present

82
States of Matter

• Matter exists in three physical states
• Gas also known as vapor
• has no fixed volume or shape
• it takes the shape of the container by either
• compression or expansion
• Liquid has a definite volume but no specific
• shape
• Solid is rigid and has a fixed volume and a
  fixed shape
• Neither liquids or solids are compressible to any
  appreciable extent

83
Phase Change

• The change of matter from one state to another
  state
• It is a PHYSICAL change
• Energy changes ALWAYS accompany a phase change
  between the three states

84
Changes in state

• Melting/freezing solid liquid
• Sublimation/deposition solid gas
• Vaporization/condensation liquid gas

85
A Graphical Representation
gas
liquid
Energy of the system
solid
86
Some other terms associated with phase changes

• The melting point is the same temperature as the
  freezing point differing only in the direction
  from which the phase change is approached
• Boiling point of a liquid is related to the
  pressure if you increase the pressure, you
  increase the boiling point and visa versa!

87
Compare and Contrast the Phases of Matter
Solid
Liquid
Gas
Similarities
88
Review Questions

• Describe the three phases of matter?
• Classify the following as extensive or intensive
  color, mass, length, melting point, ductility
• Classify the following as a chemical or physical
  property
• Reactivity
• Odor
• Rusting
• Stability
• Expansion
• Porosity

89
Heat Energy

• Although there are different forms and types of
  energy that are important to chemistry, heat
  energy has the most relevance to chemical changes
• The most obvious thing about heat energy is that
  it causes changes in the temperature of matter
• Just like matter, it too can be quantified how
  much heat?

90
How much heat?

• Heat can be lost or gained by a reaction during a
  chemical change
• Exothermic when a chemical reaction releases
  heat
• Endothermic when a chemical reaction absorbs
  heat

91
Measuring Energy Changes

• Experimentally, energy changes of chemical
  reactions are measured in a calorimeter
• To change the temperature of a substance, heat
  must be added or removed some substances
  require little heat to cause a change while
  others require a great deal of heat

92
Specific Heat

• The heat needed to raise the temperature of one
  gram of a substance by one Celsius degree is
  called the specific heat (Cp) of a substance
• Every substance has its own specific heat
• Water has a specific heat of 4.18 J/gC while
  aluminum has a specific heat of 0.900 J/gC

93
Calculating Specific Heat

• In calculating specific heat, some items must be
  understood the heat lost by one substance is
  gained by another until equilibrium is reached
• The equation used
• q (m)(?T) (Cp)
• q energy change
• m mass
• ?T Tfinal T initial
• Cp specific heat

94
Some examples

• How much heat is lost when a solid aluminum ingot
  with mass 4100 g cools from 660.0 C to 25 C?
• q ?
• m 4100g
• ?T (660 25) 635C
• Cp 0.900 J/gC
• q 4100g 635C 0.900J 2.35 x 106 J
• gC

95
Try these

• How much heat is required to raise the
  temperature of 54.5 g PCl3 from 18.6C to 79.1C?
  Cp 0.874
• How much heat is required to raise the
  temperature of 7.90 x 102 g H2O from 38.4C to
  85.4C? Cp 4.18 J/gC