Why Study Chemistry

1. Why Study Chemistry?
• To be better informed
• To be a knowledgeable consumer
• To make better decisions for yourself and
society
• To learn problem-solving skills
• To enhance analytical thinking

2. Chemistry as the “Central
Science”
• Chemistry = the study of matter and the
transformation it undergoes
• EVERYTHING is a CHEMICAL
– Table salt = sodium chloride, NaCl
– Table sugar = sucrose, C12H22O11
– Clothes: Wool? Cotton? Polyester?
– Body: lipids, Proteins, Carbohydrates, DNA/RNA
– You name it– it’s a chemical!

3. Chemistry as the “Central
Science”
• Chemistry is the driving force behind many
“liberal arts”
– Composition of paints? Colors?
– Economies of industrial nations
• #1 commercial chemical is sulfuric acid– LOTS of uses!
• All idustry involves chemical processes
– Economies of Developing Nations
• Agriculture depends on chemicals as fertilizers, pesticides
– Politics and Natural Resources

4. The Study of Chemistry
• Chemistry is everywhere!
• Matter is everywhere!
• Thus, chemistry matters!
• Chemistry involves the study of
matter – its properties and behavior.
• Macroscopic observations are
rooted in microscopic structure.

5. Assignment: Chemistry in your
major
• Find a current news story or historical
example that demonstrates the importance
of chemistry to your major
– For example: chemical resource as a key
issue in a political / economic rift; wars fought
over chemical resources; etc
• Write a 2 paragraph summary on issue
and its relevance to your studies

6. Chemistry as the “Central
Science”

7. Classification of Matter
4 Physical States: solid, liquid, gas, plasma
Solid:
Fixed shape and fixed volume;
Atoms tightly packed together

8. Classification of Matter
Liquid:
No fixed shape but maintains a fixed volume
Atoms loosely packed together, slide around
each other

9. Classification of Matter
Gas:
No fixed shape or volume
Atoms not really associated with neighbors at
all

10. Classification of Matter
Plasma:
mix of subatomic particles with not
organization
(sun)

11. States of Matter

12. Microscopic view of a
gas.
Microscopic view of a
liquid.
Microscopic view of a
solid.
States of Matter
Gases, liquids and solids are all made up of microscopic particles, but the
behaviors of these particles differ in the three phases. The following figure
illustrates the microscopic differences.

13. States of Matter
State Shape Volume Compress Flow
Solid Keeps
Shape
Keeps
Volume
No No
Liquid Takes
Shape of
Container
Keeps
Volume
No Yes
Gas Takes
Shape of
Container
Takes
Volume of
Container
Yes Yes

14. Properties of Matter
Physical Properties =
characteristics of a
material
Color
Mass
Temperature
Odor
Density
Hardness
Solubility
Conductivity (heat or
electrical)
Freezing/boiling point
Chemical Properties =
describe how a material
reacts with another type
of matter
Ability to burn
Ability to rust / corrode
Ability to make a solution
acidic or basic
Lack of ability to react with
something

15. Properties of Matter
• physical – measured without changing
substance, e.g. physical state, color, odor,
density, boiling point
• chemical – describes a substance’s
reactivity, e.g. flammability, corrosiveness
• extensive – depends on the amount of
matter present, e.g. mass, volume
• intensive – does not depend on the amount
of matter present, e.g. density, color,
temperature

16. • Properties: “ The characteristics that give each
substance its unique identity “
• Physical Properties: “ Properties that can be observed
without changing the identity of a substance “
Color
Melting Temperature - a physical change of state
Electrical conductivity
Density
Boiling Temperature - a physical change of state
Solubility
Hardness

17. Chemical Properties: “ Properties that result
in changes in the identity of one or more
reactants “
The rusting of iron
Hydrogen and oxygen burning to form water
The baking of bread
The absorption of oxygen by hemoglobin

18. Changes in Matter
Physical Changes =
a change in a
physical property;
does NOT change
the chemical
composition or
atomic arrangement
of the material
– Increase in
temperature
– Phase changes
– Cutting into smaller
pieces
Chemical Changes =
changes that alter the
identity of a material, a
change in the chemical
composition or atomic
arrangement of the
material
– Wood burns in air to
produce CO2 and H2O
– Cooking an egg (change
molecular structure of the
proteins, loss of water)
– Formation of rust (iron to
iron oxide)

19. • Properties: “ The characteristics that give each
substance its unique identity “
• Physical Properties: “ Properties that can be observed
without changing the identity of a substance
Color
Melting Temperature - a physical change of state
Electrical conductivity
Density
Boiling Temperature - a physical change of state
Solubility
Hardness continue…..
Changes in Matter: Is it Physical or Chemical?

20. • Chemical Properties: “ Properties that
result in changes in the
identity of one or more reactants “
The rusting of iron
Hydrogen and oxygen burning to form water
The baking of bread
The absorption of oxygen by hemoglobin
continue…..
Changes in Matter (cont)

21. Changes in Matter (cont)
Reactants Products
Chemical Reactions: “ Process in which one or more pure substances are
converted to one or more different pure substances “
Reactants: “ Substances that undergo change in a chemical reaction “
Products: “ Substances formed as the result of a chemical reaction “
Hydrogen + Oxygen Water
 Reactants are on the left side of the chemical equation
 Products are on the right side of the chemical equation

22. Changes in Matter - Physical &
Chemical
• Physical Change: “ A change that alters the
physical form of matter without changing its
chemical identity “
• Chemical Change: “ A change which changes the
chemical identity of the substance and creates one
or more new substances “
continue…..

23. Changes in Matter - Physical Change
A Melting Ice Sickle
Solid Water
Liquid Water
continue…..
•Example of a Physical Change:

24. 24
• Example of a Chemical Change: The Electrolysis of Water (H2O)
Hydrogen Gas
Oxygen Gas
Negative Electrode Positive Electrode
Particulate
Viewpoint
The Chemical Identity of Water ( H2O ) is changed
into the elements Hydrogen ( H2 ) and Oxygen ( O2 )
2H2O 2H2 + O2 continue…..
Changes in Matter - Chemical Change

25. 2.7 Using Chemical Symbols (cont)
Chemical Equations: “ Representations of chemical reactions by the
formulas of reactants and products “
2 C (s) + O2 (g) 2 CO
At the Macroscopic Level: “ Carbon, a solid plus oxygen gas yields
carbon monoxide “
At the Particulate Level: “ Two atoms of carbon plus one diatomic
molecule of oxygen yields two molecules
of carbon monoxide “
Equation Coefficients: “ Gives the relative amount of each compound
involved in the chemical equation “
Balanced Chemical Equations: “ The number of each kind of atom on
the reactant side must equal the
number of each kind of atom on
the product side “

26. 26
Classification of Matter
Matter - Anything that
occupies space and has
mass (solid, liquid or gas)
Heterogeneous Mixture:
Non-uniform composition
Homogeneous Matter:
Uniform composition
Pure Substances: Fixed
composition; cannot be
further purified
Physically Separable Into
Solution:
Homogeneous
mixture
Physically Separable Into
Compounds:
Elements united in
fixed ratios
Elements:
Cannot be subdivided by
chemical or physical
changes
Chemically decomposable Into
Combine Chemically to

27. The Chemical View of Matter
continue….
 What are elements and chemical compounds made of?
 What is the difference between a mixture and a pure substance?
 What is the difference between a chemical and a physical process?
 What is the basic theme of chemistry?
 How are symbols for the elements used in formulas and equations
to communicate chemical information?

28. Macroscopic, Microscopic & Particulate Matter
• Matter: - “ Anything that has mass and takes up space
• (occupies volume) “
• Matter can be studied on three levels:
• Macroscopic Level: “ Matter that can be seen with the human
eye “
 Beach Sand, Trees, Cars, Pen, CD, Mountains,
 Planets, Galaxies, etc
• Length: 101 to 109 meters
continue…..

29. • Microscopic Level: “ Matter that is too
small to be seen by the
naked eye, but can be seen under a
• microscope
 Very small plants, individual bacteria, cellular
 structures, DNA Molecule, Semiconductors, etc
• Length: 10- 6 meters
continue…..
Macroscopic, Microscopic & Particulate Matter (cont)

30. • Particulate Level: “ Matter too small to be seen
with even the most powerful optical
microscope “
 Particulate matter consists of the tiny particles
 that make up all matter
 Molecules, atoms, protons & electron
• Length: 10 - 10 meters (1 Angstrom = 10 - 10
meters )
continue…..
Macroscopic, Microscopic & Particulate Matter (cont)

31. Elements - The Most Simple Kind of Matter
Pure Substance: “Something that with a uniform, fixed
composition at the submicroscopic level”
 Recognized by the unchanging nature of their properties
Element: “A pure substance composed of only one kind of atom”
Atom: “The smallest particle of an element”
 Atoms of different elements are different and are shown
on the periodic table
Each element has a one or two letter abbreviation
 Hydrogen - H
 Helium - He
 Sodium - Na
 Lithium - Li

32. Microscopic view of the atoms of the
element argon (gas phase).
Microscopic view of the molecules of the
element nitrogen (gas phase).

33. Elements

34. 34
The Periodic Table and the Elements (cont)
Transition Metals
continue….
Main Group
Elements
Main Group Elements
Inner Transition Elements

35. Chemical Compounds - Atoms in Combination
Chemical Compounds: “ Pure substances made of atoms of different
elements combined in definite ways”
Examples:
 H2O Water
 NaCl Sodium Chloride
 C2H6O Ethanol
 C6H12O6 Sugar

36. Chemical Compounds (cont)
• Compound: “ Any pure substance that can
be decomposed by a
• chemical change into two
or more pure substances
• is a compound “ -
(another definition)
• Compounds are made up of elements
• Examples of Compounds: continue…..

37. Using Chemical Symbols
Chemical Formulas: “ Combinations of the symbols for the elements
that represent the stable combinations of atoms
in molecules “
Examples:
Water H2O
Carbon dioxide CO2
Ammonia NH3
Methane CH4
Carbon Tetrachloride CCl4
Subscripts: “ Indicate the relative numbers of atoms of each kind “

38. Using Chemical Symbols (cont)
Structural Formulas: “ Formulas that show the connections between
atoms in molecules “
H N H
H
Ammonia
H O H
Water
H C H
H
Methane
H

39. Microscopic view of the molecules of the compound
water (gas phase). Oxygen atoms are red and hydrogen
atoms are white.

40. Mixtures and Pure Substances
• Homogeneous Sample: “ Matter that has a
uniform appearance and
• composition throughout “
A mixture of water and alcohol
Sugar dissolved in water
Gold blended with silver (18 karat gold)
The air we breathe - a mixture of oxygen and
nitrogen
continue…..
Solutions: “Homogeneous mixtures, either liquid, solid or gaseous”

41. 41
Table salt is stirred into water (left), forming a
homogeneous mixture called a solution (right)
Mixtures and Pure Substances (cont)
continue…..

42. Heterogeneous Sample: “ Matter that does not have a uniform
appearance and composition
throughout “
 A mixture of cooking oil and water (two phases develop)
 Concrete (sand, rock, cement, etc)
 A mixture of sand, sawdust, iron fillings and water
Mixtures and Pure Substances (cont)
continue…..

43. Sand and water do not mix to form a
uniform mixture
Mixtures and Pure Substances (cont)
continue…..

44. Mixtures:
Homogeneous
• Same composition
throughout sample
• Ex- milk, tea,
others?
Heterogeneous
• Different samples of
the same mixture
have different
compositions
• Ex- air in the room
others?

46. Microscopic view of a gaseous mixture containing two
elements (argon and nitrogen) and a compound (water).

47. Classification of Matter

48. Substances vs Mixtures
Substance
– has a definite or fixed
composition
– Composition does not
vary from sample to
sample
Mixture
– Has a varied
composition
– Each individual
component can be
separated by physical
means
– Ex: salt and pepper,
sugar in water, sea
water

49. Energy
• The “fuel” of the universe
• The capacity of something to do work
– chemical, mechanical, thermal, electrical,
radiant, sound, nuclear
• The SI unit of energy is the Joule (J)
– Other common units are
• Calories (cal)
• Kilowatt-hour (kW.hr)
• Types of energy:
– Potential
– Kinetic
– Heat
• Energy cannot be created nor destroyed (but it
does change from one type to another!)

50. Changes in Matter - Energy
Energy: “ The ability to cause change or, in formal terms of physics,
the ability to do work “
Potential Energy: “ Energy in storage “
 There is potential energy in gasoline called chemical energy
 Chemical energy is release as heat and light when it burns
 Chemical energy can also be released as electrical energy
Kinetic Energy: “ Energy in motion “
 Examples are - Muscle in movement, a rocket in flight,
inflation of a car air bag during collision

51. Heat & Temperature
• Temperature is _____.
– how hot or cold something is (a physical property)
– related to the average (kinetic) energy of the
substance (not the total energy)
– Measured in units of
• Degrees Fahrenheit (oF)
• Degrees Celsius (oC)
• Kelvin (K)
• Heat is energy that _____.
– flows from hot objects to cold objects
– is absorbed/released by an object resulting in its
change in temperature
• Heat absorbed/released is measured by
changes in temperature

52. Substances
Elements
• Fundamental
substances from
which all things are
constructed
• Only one type of
atom is present
• Can not be broken
down any further

53. Substances
Compounds
• Substances made
up of two or more
elements in distinct
ratios
• Molecules: smallest
characteristic part of
a compound;
composed of a
distinct and unique
arrangement of
elements

54. Temperature Scales
• Fahrenheit Scale, °F
– Water’s freezing point = 32°F, boiling point = 212°F
• Celsius Scale, °C
– Temperature unit larger than the Fahrenheit
– Water’s freezing point = 0°C, boiling point = 100°C
• Kelvin Scale, K
– Temperature unit same size as Celsius
– Water’s freezing point = 273 K, boiling point = 373 K

55. Temperature of ice water and boiling
water.

56. Heat
• Heat is the flow of energy due to a temperature
difference
– Heat flows from higher temperature to lower
temperature
• Heat is transferred due to “collisions” between
atoms/molecules of different kinetic energy
• When produced by friction, heat is mechanical
energy that is irretrievably removed from a
system
• Processes involving Heat:
1. Exothermic = A process that releases heat energy.
• Example: when a match is struck, it is an exothermic
process because energy is produced as heat.
2. Endothermic = A process that absorbs energy.
• Example: melting ice to form liquid water is an endothermic
process.

57. Heat (cont.)
• The heat energy absorbed by an object is
proportional to:
– The mass of the object (m)
– The change in temperature the object undergoes
(DT)
– Specific heat capacity (s) (a physical property unique to
the substance)
• To calculate heat (Q):
Q = c . m . DT

58. Specific Heat Capacity (c)
• The amount of heat energy (in J or Cal) required to
increase the temperature of 1 gram of a
substance by 1oC (or 1K)
• The Units of Specific Heat Capacity:
1. J/goC (SI)
2. cal/goC (metric & more useful in the lab)
• Specific Heat Capacity is a unique physical
property of different substances
– Metals have low specific heat capacity
– Non-metals have higher specific heat capacity
– Water has an unusually large specific heat capacity
c = Q/(mDT)