This document provides an overview of chemistry concepts including the definition of chemistry, major branches of chemistry, early theories of matter, and important figures in the development of modern chemistry such as Aristotle, Democritus, Boyle, Priestley, and Dalton. It also discusses the classification of matter as elements, compounds, and mixtures. Key chemistry concepts like physical and chemical properties, physical and chemical changes, energy, heat, and phase changes are introduced.

1. Science, Chemistry and You

2. Chemistry
• Definition – study of the composition and
properties of matter and the energy
transformations accompanying changes in the
structure of matter

3. Major Branches of Chemistry
• Inorganic Chemistry – Study of all the
elements other than Carbon
• Organic Chemistry – Study of compounds
containing carbon
• Biochemistry – study of chemical processes in
living things
• Nuclear Chemistry – study of radioactivity,
the nucleus and the changes that the nucleus
undergoes

4. Aristotle
Early Greek Theories
• 400 B.C. - Democritus thought matter
could not be divided indefinitely.
• 350 B.C - Aristotle modified an earlier
theory that matter was made of four
“elements”: earth, fire, water, air.
Democritus
• Aristotle was wrong. However, his
theory persisted for 2000 years.
• This led to the idea of atoms in a
void.

5. The Rise of Modern Chemistry
• The Greek idea of the 4 basic elements
was not disputed until the mid 1600s
• Robert Boyle proposed that elements
are substances that cannot be
chemically decomposed into simpler
substances. Earth, air, fire and water
could not be called elements
• In 1774 Joseph Priestly discovered a
gas in which substances burned easily,
Antoine Lavoisier named the gas
Oxygen
Boyle
Priestly

6. John Dalton
• 1800 -Dalton proposed a modern atomic model
based on experimentation not on pure reason.
• All matter is made of atoms.
• Atoms of an element are identical.
• Each element has different atoms.
• Atoms of different elements combine
in constant ratios to form compounds.
• Atoms are rearranged in reactions.
• His ideas account for the law of conservation of
mass (atoms are neither created nor destroyed)
and the law of constant composition (elements
combine in fixed ratios).

7. Reaction of the Day
Table sugar + sulfuric acid  Carbon + H20
H2SO4
C12H22011 (s)  12 C (s) + 11 H2O (g)

8. Ch 2 - Matter
Matter – anything that takes up space and has mass

9. Chemical and Physical
Properties of Matter
Physical properties – color, shape, texture, odor,
taste, electrical conductivity, and density
density – how closely packed the molecules are
malleable – substances that can be easily
hammered into shapes
ductility – substances that can be stretched into
wires
conductivity – substances that can transfer heat
or electricity
Chemical properties – describe how matter acts in
the presence of other materials

10. What is each picture modeling?
Density, malleability, ductility, conductivity

11. Physical or Chemical Change

12. Physical vs. Chemical Change
Physical Change
• Atoms do not rearrange
• Only physical properties change. Chemical properties do not
change.
• Physical changes are generally easy to reverse.
• No energy is produced by the substance.
Chemical Change
• Atoms are rearranged into different molecules
• Both physical and chemical properties are changed
• Changes are not reversible without another reaction
• Energy is often produced ( fire or heat, for example)

13. Identify each of the following as a Physical or Chemical Change.
Put a P next to Physical Changes and a C next to Chemical Changes
1. A piece of wood burns to form
ash.
2. Water evaporates into steam.
3. A piece of cork is cut in half.
4. A bicycle chain rusts.
5. Food is digested in the
stomach.
6. Water is absorbed by a paper
towel.
7. Hydrochloric Acid reacts with
zinc.
8. A piece of an apple rots on the
ground.
9. A tire is inflated with air.
10. A plant turns sunlight, CO2,
and water into sugar and
oxygen.
11. Sugar dissolves in water.
12. Eggs turn into an omelette.
13. Milk sours.
14. A popsicle melts.
15. Turning brownie mix into
brownies.

14. Demonstration of the day
Vinegar + baking soda
Acetic acid + sodium bicarbonate  carbon dioxide +
water + sodium acetate
Heterogeneous mixture containing, solid, liquid and gas phases

15. The Division of Matter
Two major categories:
1) pure substances - consists of only one type
of matter, which cannot be separated into
other kinds of matter by any physical
processes. Ex: Olive oil
2) mixtures – material that can be separated
by physical means into two or more pure
substances. Ex: Oil and vinegar salad dressing

16. Two Types of Mixtures
• Heterogeneous – a mixture in which
the substances are not uniformly
mixed
Ex: oil & vinegar dressing, granite
has quartz & mica
• Homogeneous – a substance in
which the particles are uniformly
mixed
Ex: dough & air

17. Elements and Their Symbols
Element - pure substance that cannot be
broken down into simpler substances

18. Elements and Their Symbols
• Atoms – smallest particles that maintain the
physical and chemical characteristics of an
element
• Monoatomic elements – elements that do not
naturally combine or bond together. Ex: Ne,
He, Ar
• Diatomic elements - elements that bond into
two-atom units. Ex: O2, H2
• Polyatomic elements – elements composed
of multi-atom units. Ex: S8

19. Elements and Their Symbols
Symbol – letter given to represent the name of
each element
Hydrogen
Oxygen
Calcium
Magnesium
Manganese
Sodium

20. Compounds and Their Formulas
• Compounds are made up of atoms from two
or more different elements, chemically
bonded together
• Formulas tell the type and number of atoms
that are present in compounds
Common Compounds and Their Formulas
Compound Formula Atoms
Ammonia NH3 1 nitrogen, 3 hydrogen
Rust Fe2O3 2 iron, 3 oxygen
Salt NaCl 1 sodium, 1 chlorine
Sucrose C12H22O11 12 carbon, 22 hydrogen, 11 oxygen

21. Sample Problems
How many atoms of each element are present in
each of the following groups?
a.Na2S2O3
b.Mg(NO3)2
c. 5 Fe2O3

22. Molecule
• The smallest independent units of
compounds
• Consist of two or more atoms that are
chemically bonded together
• Ex: H20, NH3, H2SO4
• Homework: Read pgs 21-28
Section Review Questions 2A, pg 29, #1-3

23. Tuesday September 14, 2010
• Go over homework problems

24. 2B Energy in Matter
• Every chemical reaction either releases or
absorbs energy
• Exothermic reactions – release energy (get hot)
Ex: lighting a match
• Endothermic reactions – absorb energy (get
cold) Ex: ice pack

25. Energy – the ability to do work
• There are many forms of energy
• Chemistry is concerned with the relationship among
chemical, thermal, electrical and nuclear energy

26. Energy Conservation
• Thermodynamics – the study of energy flow
• First Law of Thermodynamics or Law of
Conservation of Mass-Energy –matter and
energy can neither be created nor destroyed,
simply changed from one form to another
• Second Law of Thermodynamics – during any
energy transformation, some energy goes to
an unusable form

27. Energy Conservation
• Entropy –
randomness or
disorder of a
system
• There is a tendency
for all natural
processes to
increase in entropy
(disorder)

28. Heat, Energy & Temperature
• Kinetic Energy – energy of motion
All matter contains particles that are moving
• Thermal Energy – sum of all the kinetic energy
of an object
• Temperature measures the average kinetic
energy of all the particles in a sample
• Heat – thermal energy that is transferred from
one object to another
• Amount of heat transferred between objects is determined
by the temperature difference between them and the mass of
the hotter object

29. Which contains more thermal energy?
A teaspoon of boiling water or a bathtub full of
lukewarm water
Which has a higher temp?

30. The Measurement of Energy
• Joule – standard unit of measurement for
energy
• BTU – English unit of measurement for
thermal energy, the amount of heat required
to raise one pound of water by one degree
Fahrenheit
• Calorie – amount of energy required to raise
the temperature of one gram of water one
degree Celsius
• 1 cal = 4.184 J

31. Temperature Scales
Celsius scale – freezing point of water is 0◦
C
boiling point of water is 100◦
C
Kelvin scale – uses absolute zero (point at which
molecules no longer move) as the zero point
freezing point of water is 273 K
boiling point of water is 373 K
Fahrenheit scale – freezing point of water is
32◦
F
boiling point of water is 212◦
F

33. Conversion between scales
K = ◦
C + 273 ◦
C = K - 273
◦
F = (1.8 x ◦
C) ◦
C = (◦
F-32)/1.8
Sample Problem: The weatherman announces
that the high for the day is expected to be 33◦
C
What is this temperature on the Kelvin scale
and the Fahrenheit scale?

35. Phase Changes of Matter
• Condensation –gas to liquid
• Vaporization – liquid to gas
• Freezing – liquid to solid
• Melting –solid to liquid
• Sublimation – solid to gas
• Deposition – gas to solid

36. Tuesday Homework
Read pgs 29 – 39
Section Review Questions 2B
Pg 36, questions 1 - 4

37. Wednesday
• Do Review Questions pg 40 & 41

38. Thursday Go Over Review

39. Friday
• Test Ch 1&2