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MATTER
States of Matter
Gas
- no fixed volume or shape
- easily compressed or expanded
Examples: air, neon, acetylene
Liquid
- has fixed volume
- shape of liquid is shape of container
- not easily compressed
Examples: water, gasoline, blood
Solid
- has fixed volume and shape
- not easily compressed
Examples: salt, steel, wool
Plasma
- the fourth state of matter where some electrons are separated from the rest of
the matter
- a gas of charges created from high voltage.
Examples: sun, neon light
CLASSIFICATIONS OF MATTER
I. Homogeneous – same composition throughout sample
A. Pure substances
1. Elements
- cannot be decomposed into simpler substances
- building blocks of all matter
- only 115 elements, listed in periodic table
- individual constituents are atoms
Examples: oxygen, iron, helium
2. Compounds
- made of two or more elements
- millions of compounds
- individual constituents are molecules or ions
- atoms (elements) bind together to form molecules (compounds)
Examples: sugar, water, rust
B. Solutions
- perfect blending of two or more pure substances with each other
- no region of matter has only one type of molecule
- examples
a) air N2 in O2
b) sea water salt in water
c) brass
Copper (Cu) in Zinc (Zn)
d) soda pop CO2 and sugar in water
e) cocktail alcohol dissolved in water
f) millions of others possible
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II. Heterogeneous
-must be imperfect mixture
-properties are different throughout sample
-examples
a) sand
b) wood
c) meat
d) potato
e) butter
PROPERTIES OF MATTER
Physical properties – identity of substance does not change when measured
intensive properties
– independent of amount of substance
– Examples: color, density, melting point, temperature
extensive properties
– dependent on amount of substance
– Examples: mass, volume, energy
Chemical properties – how substance reacts with other substances
Consider how wood changes when we mix it with other substances
- When we heat it with oxygen (burning), the wood emits smoke (carbon
dioxide, water vapor, etc.) and leaves behind charcoal and ash (carbon).
- When we heat it with lead, no chemical change occurs. (Lead melts on
surface of wood.)
- When we put a drop of concentrated sulfuric acid on it, smoke is released and
smells nasty (hydrogen sulfide) and wood turns brown.
- When we put a drop of water on it, no chemical change occurs. (It only gets wet.)
CHANGES IN MATTER
Physical change – composition remains the same
- ice melting
- adjustable tint sunglasses
- dissolving sugar in water
Gas
g
ltin
me zing
e
fre
de
po
su
s
bli ition
ma
tio
n
Phase changes
- when one phase of matter changes to another
- a phase change is a physical change
Solid
evaporation
condensation
Liquid
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Melting point – temperature where solid melts
- same as freezing point
Tm(H2O) = 0 C
Tm(O2) = -223 C
Tm(Fe) = 1536 C
Boiling point – temperature where liquid boils
- same as condensation point
Tb(H2O) = 100 C
Tb(N2) = -196 C
Tb(ethanol) = 78 C
Chemical change – composition is altered
Examples:
- burning of paper
- baking cake
- cooking egg
NOTES ON DENSITY
Riddle: Which is heavier, a pound of feathers or a pound of lead?
Answer: Neither, they both weigh one pound.
Question: Which has more volume, a pound of lead or a pound of feathers?
Answer: A pound of feathers
Since feathers take more space than lead, the feathers are less dense than lead.
Definition of density
density 
mass
volume
- in chemistry, density often has the units of g/mL or g/cm3.
Example: What is the density of motor oil? A 755 mL sample of motor oil weighs
733 g.
Density can be considered a conversion factor between mass and volume.
d(H2O) = 1.00 g/mL  1.00 g = 1.00 mL

100
. g H2O
mL H 2 O
or
100
. mL H 2O
g H 2O
d(Hg) = 13.6 g/mL  13.6 g = 1.00 mL
13.6g Hg 1.00mL Hg

or
mL Hg
13.6g Hg
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Example: The density of kerosene is 0.82 g/mL. How much mass does 250 mL
contain?
Example: How much space does 4.76 kg of lead occupy? The density of lead is
11.3 g/mL.
Note: we will have to include a conversion from kg to g in our calculation.
ENERGY IN CHEMICAL REACTIONS
Energy is the capacity to do work (to move an object through a distance).
Forms of Energy
Kinetic – energy which causes motion
- A moving ball has kinetic energy.
- Electrical energy is kinetic energy. (Electrical charge is moving.)
- Hot water has kinetic energy. (Water molecules are moving.)
Potential – energy which has ability (i. e. potential) to cause motion
- Apple on a branch has potential energy. (Gravity is pulling on it.
When it drops, it has motion.)
- Battery has potential energy. (When circuit is complete, electricity will flow.)
TYPES OF ENERGY
Light Energy – as in energy from a light bulb
Heat Energy – as in energy from a fire
Electrical Energy – as in putting your finger in an electrical socket Ouch!
Mechanical Energy – as in energy from an engine or water wheel
Chemical Energy – as in energy from burning gasoline (chemical energy is
converted to heat and light)
- Chemical energy is stored inside a substance until a substance undergoes a
chemical change. Then it can be released or sometimes absorbed.
- Chemical energy is stored in chemical bonds between atoms. (More about
bonds later)
Energy exchange in chemical reactions
Exothermic
- energy is released when reaction occurs
- most reactions are exothermic
Endothermic
- energy is absorbed when reaction occurs
- example is an instant cold pack.
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Calculation of heat energy
Units of energy
Joule – SI unit (J)
1 kg  m
 1 m
Definition: 1J  1 kg  
 
 s 
s2
2
2
A 0.250 kg baseball going 60 mi/hr (27 m/s) has 90 J of energy (kinetic).
The amount of heat energy needed to start 1 quart of water boiling is about 10,000 J.
Calorie – traditional (English) unit (cal)
Definition
a) original – amount of heat needed to raise 1 gram of water 1 C
b) modern – 1 cal = 4.184 J (exactly)
Note: 1 dietary calorie (Cal) = 1 kilocalorie (kcal)
HEAT CHANGES AND SPECIFIC HEAT
Heat energy within a substance is proportional to both mass and temperature
difference.
- when mass increases, more heat is needed to raise temperature to same level.
- when the temperature difference increases (i. e. to make temperature hotter), the
amount of heat needed increases
Heat Equation
q = m c T
q – heat
m – mass
T – temperature difference = Tfinal - Tinitial
c – specific heat
Specific Heat
- shows how much heat is needed to raise temperature of a substance
- compare specific heats of Cu and H2O.
c(Cu) = 0.3844 J/gC
c(H2O) = 4.184 J/gC
- specific heats show that copper heats and cools quickly while water heats and
cools slowly.
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Example: How much heat is lost when 3780 g of H2O (1 gallon) is cooled from 61.4
C to 30.0 C? c(H2O) = 4.184 J/gC
Negative sign means heat is lost. (More about this later.)
Example: 98.3 g of Aluminum at 25 C is given 15,500 J of heat. What is the final
temperature of the metal? c(Al) = 0.895 J/gC
LAW OF CONSERVATION OF MASS
In a chemical reaction, mass is not created or destroyed.
We use the law of conservation of mass later to account for all the mass in a
chemical reaction.
LAW OF CONSERVATION OF ENERGY
In a chemical reaction, energy is not created or destroyed.
Question: If this is true, then where does energy from burning gasoline come from?
Answer: From the chemical energy inside the gasoline