Unit 3 - Properties of Matter
Part 1 - The Basic Properties of Matter
Matter is all the physical things in the universe. All
the stars in the galaxies, the sun and planets in
our solar system, the Earth, and everything on it
and in it are matter.
What is matter???
MATTER – anything that takes up space and has
mass
Big ideas which are always true when talking
about matter:
1. Matter (solid, liquid, and gas) is made up of
tiny particles called atoms and molecules.
2. The atoms or molecules that make up matter
are always in motion.
There are 3 common states or phases of Matter Solid, Liquid, Gas
It is the motion of the molecules in matter that
determines which state or phase of matter a
substance is in
SOLID
·
·
·
has a definite shape (rigid)
definite volume
particles vibrate around fixed
positions
LIQUID
·
No definite shape (takes the shape of
its
·
·
container)
Has definite volume
Particles are free to move over each
other, but are still attracted to each
other
GAS
·
No definite shape (takes the shape of
its
·
·
·
container)
No definite volume
Particles move in random motion with
little or no attraction to each other
Highly compressible
THE MOVEMENT OF MATTER’S PARTICLES The Kinetic Theory of Matter
- all matter is composed of tiny particles: atoms,
molecules, ions, or some combination of these
- they are always in motion
-the various states of matter differ from each
other on the basis of their motion.
Solids - move very slowly Liquids - move more rapidly
Gases - move much more rapidly than either
solid or liquid particles
PLASMA - The fourth state of matter. A very hot
gas like mixture of electrons and positive ions, the
atoms that are left after their electrons have
been removed.
99% of Universe
Ex. Stars, Lightening, Fluorescent lights
Plasma is electromagnetic - it can carry a charge
CLASSIFICATION OF MATTER - MATTER IS EITHER
PURE OR MIXED
Matter can be sub-divided into two categories:
mixtures and pure substances.
1. MIXTURE
-has variable composition
-the proportions that make up the mixture vary
-can be separated by physical methods
Examples: gasoline, wine, soil, air, chocolate
chip cookie dough
Mixtures can be
A.
B.
HOMOGENEOUS – having visibly (to
the naked eye) indistinguishable
parts Example: wine, air
HETEROGENEOUS – having visibly
distinguishable parts
Example: Chocolate Chip Cookie
dough, soil
*SOLUTION – a homogeneous mixture
Examples: air (a gaseous solutions), wine (a liquid
solution), brass (a solid solution of copper and
zinc)
Which one of the following is heterogeneous
mixture?
I. Coke
II. Sea Water
III. Water+Sand
IV. Natural Gas
PURE SUBSTANCES – have constant composition
and can only be separated by chemical
reactions
2 Types: Elements and compounds
ELEMENTS - substances that cannot be
broken into simpler substances by chemical
or physical means
COMPOUNDS – substance with constant
composition that can be broken down into
elements by chemical processes
So how are Compounds and Mixtures different if
they are both made of more than one element?
1. Ratio between matters forming compound is
constant but ratio between matters forming
mixture is variable.
2. Matters forming compounds loose their
properties but matters forming mixtures preserve
their properties.
3. We cannot breakdown compounds with
physical methods - only chemical. Mixtures can
breakdown physically.
http://studyjams.scholastic.com/studyjams/jams/
science/matter/mixtures.htm
PHYSICAL AND CHEMICAL PROPERTIES
All substances have Physical and Chemical
properties that we can use to identify them.
PHYSICAL PROPERTIES: Properties that do not
change the chemical nature of matter
Measuring each of these properties will not alter
the basic nature of the substance.
Density is an important physical property of
matter.
The density of a substance can often be used to
help identify it.
CHEMICAL PROPERTIES: Properties that change
the chemical nature of matter
Measuring each of these properties will alter the
basic nature of the substance.
PHYSICAL
PROPERTIES
CHEMICAL PROPERTIES
Color
Melting Point
Boiling Point
Solubility
Hardness
Strength
Elasticity
Heat
Conductivi
ty
Electrical
pH
Reaction with oxygen
(flammability or corrosion)
Reaction with water
Reaction with acids and bases
Reaction with metals
Conductivi
ty
Ability to
transmit
light
Lustre
('shininess'
or dullness)
Magnetic
attraction
'
PHYSICAL AND CHEMICAL CHANGES
PHYSICAL CHANGE - occurs when no new
substance is made, and the change is usually
easy to reverse.
Examples of a Physical Change - dissolving
common salt in water - still salty and wet!
ANOTHER EXAMPLE OF PHYSICAL CHANGES:
CHANGES OF STATE - Matter changes from one
state to another
Names of phase changes
Starting
phase:
solid
Change to:
Name
Liquid
melting
FreezingSolidification
Boiling evaporation
liquid
Solid
liquid
Gas
gas
solid
gas
Liquid
gas
(skipping
liquid phase)
solid
(skipping
liquid phase)
Condensation
Sublimation
Deposition
A phase change depends upon whether heat is
being added or heat is being taken away
- Heat is applied - change in state typically goes
from solid to liquid to gas.
- A material is cooled - change in state typically
goes from gas to liquid to solid.
Part 3 - Heat, Temperature and Chemical
reactions
Temperature and heat are not the same thing!
Temperature
- a measure of how hot something is
- measured in ºC
- physical property of matter.
- The temperature of matter = motion of the
molecules
- The greater the motion the higher the
temperature
http://www.iun.edu/%7Ecpanhd/C101webnotes/
matter-and-energy/specificheat.html
Heat
- a measure of the thermal energy contained in
an object
- measured in J (Joules)
Heat energy flows from a hot object to a cooler
one. This causes:
- hot objects to cool down
- cool objects to warm up
When heat energy is transferred to an object, its
temperature increase depends upon:
- the mass of the object
- the substance the object is made from
- the amount energy transferred to the object
- For a particular object, the more heat energy
transferred to it, the greater its temperature
increase.
Chemical reactions give off and take in heat
Endothermic vs. Exothermic
When trying to classify a process as exothermic or
endothermic, watch how the temperature of the
surroundings changes.
exothermic process - releases heat, and causes
the temperature of the immediate surroundings
to rise.
endothermic process - absorbs heat and cools
the surroundings.
Exothermic processes
making ice cubes
formation of snow in
clouds
condensation of rain from
water vapor
Endothermic
processes
melting ice cubes
conversion of frost
to water vapor
evaporation of
water
a candle flame
mixing sodium sulfite and
bleach
rusting iron
burning sugar
forming a cation
from an atom in
the gas phase
baking bread
cooking an egg
producing sugar
by photosynthesis
Specific Heat and Specific Heat Capacity
Changing temperature
Specific heat - energy required to raise the
temperature of 1g of a substance by 1C
Specific heat capacity - how much heat energy
a substance can hold. It is the energy needed to
increase the temperature of 1 kg of the
substance by 1 ºC.
Different substances have different specific heat
capacities.
Water has a particularly high specific heat
capacity.
This makes water useful for storing heat energy this means:
1. It can transport it around the home using
central heating pipes.
2. It moderates the Earth's climate by causing
temperatures to change slowly in areas around
large bodies of water.
Because of the high specific heat of water, water
and land near bodies of water are heated more
slowly than land without water. More heat
energy is necessary to heat up the area because
the water absorbs the energy. A similar
amount of heat energy would increase the
temperature of dry land to a much higher
temperature, and the soil or dirt would keep the
heat from going into the ground. Deserts reach
extremely high temperatures specifically
because of their lack of water.
Thermal energy (heat energy) moves from one
place to another because of the difference in
temperature between them. The energy transfer
is always from hot to cold.
Heat can be transferred (moved) three ways:
1. Conduction
The movement of heat from one molecule to
another
Needs direct contact
Heat flows from a higher-temperature area to a
lower-temperature one
2. Convection
The movement heat by currents in liquids or
gases
circulation through a fluid
cool air sinks down, while warmer air rises to the
top
3. Radiation
Energy movement through electromagnetic
waves
A way in which energy is transferred from place
to place in the form of a wave
THE KINETIC THEORY OF MATTER: atoms and
molecules (particles) are in constant motion - the
higher the temperature - the higher the speed increased heat energy make atoms and
molecules move faster
Gases have:
more kinetic energy - higher temperatures more heat
particles that move quickly and randomly
no fixed shape or volume
Liquids have:
less kinetic energy - lower temperatures less heat
particles move quickly
a fixed volume, but take the shape of
containers
Solids have:
the least kinetic energy - lowest
temperatures - less heat
particles vibrate in place
a fixed volume and shape
Phase changes (changes of state)
require a gain or loss of energy to occur.
Gas to
liquid
Liquid
to solid
Solid to Liquid
liquid
to gas
Gas
to
solid
Solid
to gas
Change in temperature
When a material reaches the temperature at
which a change in state occurs, the temperature
will remain the same until all the energy is used to
change the state.
Some Boiling and freezing temperatures
Material
H2O
(water)
Fe (iron)
O
(oxygen)
Hg
(mercury
Boiling (°C)
becomes a
gas
Freezing (°C)
becomes a solid
100° C
0° C
2750° C
1535° C
-183° C
-218° C
357° C
-39° C
)
Ethyl
Alcohol
78° C
-114° C
Using the data given in the chart above, explain
how butter can be considered frozen at room
temperature.
CHEMICAL CHANGES (ALSO CALLED CHEMICAL
REACTION)
1. occurs when a new substance is made
2. often the change is difficult or impossible to
reverse
Examples of a Chemical Change - cooking a
cake, cooking a roast chicken, burning wood
Signs that a Chemical Change or Chemical
Reaction has taken place:
1. Formation of a Precipitate - solid formed during
a chemical reaction between two liquids. The
precipitate differs from both of the reactants and
generally occurs when solutions containing ionic
compounds are mixed.
2. Color Change- Ex. changing color of leaves
and food that has spoiled
3. Production of Gases
4. Temperature Change
5. Light or Sound-. Ex. A fireworks
MATTER CAN BE CLASSIFIED AS METAL, NONMETAL
OR METALOID
METALS
Most elements are metals.
Physical Properties of Metals:
Luster (shininess)
Good conductors of heat and electricity
High density (heavy for their size)
High melting point
Ductile (drawn out into thin wires)
Malleable (hammered into thin sheets)
Chemical Properties of Metals:
Easily lose electrons
Corrode easily. Corrosion is a gradual wearing
away. (Example: silver tarnishing and iron rusting)
NONMETALS
Nonmetals’ characteristics are opposite of
metals.
Physical Properties of Nonmetals:
No luster (dull appearance)
Poor conductor of heat and electricity
Brittle (breaks easily)
Not ductile
Not malleable
Low density
Low melting point
sulfer
Chemical Properties of Nonmetals:
Tend to gain electrons
Metals and nonmetals like to form compounds
with each other. These compounds are called
ionic compounds.
When two or more nonmetals bond with each
other, they form a covalent compound.
METALLOIDS
Have properties of both metals and nonmetals.
Physical Properties of Metalloids:
Solids
Can be shiny or dull
Ductile
Malleable
Conduct heat and electricity better than
nonmetals but not as well as metals
Silicon
Chemical Formulas
Ammonia has one
Nitrogen atom and
three Hydrogen atoms
in its molecule.
So, its chemical
formula must show
that with Subscripts
Chemical Reactions
2 Parts of a chemical reaction
1. Reactants - break down or combine
chemically to form a new substance (what you
start with)
2. Products - what is formed
Conservation of Mass
Mass is neither created nor destroyed
during the course of a chemical reaction
The mass of reactants must equal the
mass of the products in a chemical reaction
Iron + Oxygen -----> Rust
100 g + ?g ------> 143g
mass reactants = mass products
mass products = 143g = mass reactants
= 100 + mass of oxygen
mass oxygen is??
This can be visualized by considering the
formation of water from oxygen and hydrogen
molecules:
Note that the hydrogen and oxygen atoms
simply rearrange themselves but are not
destroyed. Therefore mass is conserved.
Chemical equations - show the reactants and
products. An arrow is used to indicate that a
chemical reaction has taken place - like this:
Reactants→Products
Example 1
You light your gas range. The reaction is Methane
reacting with the oxygen in the atmosphere to
produce carbon dioxide and water vapor.
The chemical equation that represents this
reaction is written like this:
This is what it would look like if we could see the
molecules:
You can read the equation like this: One
molecule of methane gas reacts with two
molecules of oxygen gas to form one molecule
of carbon dioxide gas and two molecules of
water vapor. The 2 in front of the oxygen gas and
the 2 in front of the water vapor are called the
reaction coefficients.
Label the products and reactants in this
chemical reaction?
In this reaction, all reactants and products are
invisible. The heat being given off is the clue that
tells you a reaction is taking place. This is a good
example of an exothermic reaction.
Example 2: The reaction between hydrogen and
oxygen to form water is represented by the
following equation.
2 H2 + O2 2 H2O
Balancing Chemical Equations
http://www.youtube.com/watch?v=XPFUnQ1X1_
0&feature=related
http://misterguch.brinkster.net/eqnbalance.html
There is no sequence of rules that can be
followed blindly to get a balanced chemical
equation!
You must change the COEFFICIENTS written in
front of the formulas of the reactants and
products until the number of atoms of each
element on both sides of the equation are the
same.
Think of them like number puzzles
It is usually a good idea to tackle the easiest part
of a problem first.
Example: Consider what happens when propane
(C3H8) burns in air to form CO2 and H2O. The first
thing to look for when balancing equations are
relationships between the two sides of an
equation. This means - whatever is one side must
be on the other side too. You must always
remember to read both the coefficients and the
subscripts in the equations.
_____ C3H8 + _____ O2 _____ CO2 + _____ H2O
1 C3H8 + _____ O2 3 CO2 + _____ H2O
1 C3H8 + _____ O2 3 CO2 + 4 H2O
1 C3H8 + 5 O2 3 CO2 + 4 H2O
C3H8(g) + 5 O2(g) 3 CO2(g) + 4 H2O(g)
Practice Problem:
Write a balanced equation for the reaction that
occurs when ammonia burns in air to form
nitrogen oxide and water.
_____ NH3 + _____ O2  _____ NO + _____ H2O
Types of Chemical Reactions
4 Basic Types: synthesis, single replacement,
double replacement or decomposition
Synthesis: two or more elements or compounds
may combine to form a more complex
compound.
Basic form: A + X →
AX
Examples of synthesis reactions:
Metal + oxygen → metal oxide
EX. 2Mg(s) + O2(g) →
2MgO(s)
Decomposition
: A single compound breaks down into its
component parts or simpler compounds.
Basic form: AX → A + X
Examples of decomposition reactions:
Metallic carbonates, when heated, form metallic
oxides and CO2(g).
EX. CaCO3(s) →
CaO(s) + CO2(g)
Replacement: a more active element takes the
place of another element in a compound and
sets the less active one free.
Basic form: A + BX →
AY + X
AX + B or AX + Y →
Examples of replacement reactions:
Replacement of a metal in a compound by a
more active metal.
EX. Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
Single vs. Double:
Single Replacement Reactions
A substance, C, exchanges itself for a
component in a compound, AB, to form another
compound, AC, and a single free molecule, B, as
shown.
Double Replacement Reactions
A double replacement reaction is based on the
same principle as the single replacement
reaction, but, You start with two substances, AB
and CD. Here A replaces C to form AD, but C
joins with B, to form BC, as seen below.
AB +
CD 
AD + BC
Handy Checklist for figuring out what type of
reaction is taking place:
Follow this series of questions. When you can
answer "yes" to a question, then stop!
1) Does your reaction have one large molecule
falling apart to make several small ones? If yes,
then it's a decomposition reaction
2) Does your reaction have two (or more)
chemicals combining to form one chemical? If
yes, then it's a synthesis reaction
3) Does your reaction have any molecules that
contain only one element? If yes, then it's a
single replacement reaction
4) If you haven't answered "yes" to any of the
questions above, then you've got a double
replacement reaction
Sample Problems
List what type the following reactions are:
1) NaOH + KNO3 --> NaNO3 + KOH
2) 2 Fe + 6 NaBr --> 2 FeBr3 + 6 Na
3) CaSO4 + Mg(OH)2 --> Ca(OH)2 + MgSO4
4) Pb + O2 --> PbO2
5) Na2CO3 --> Na2O + CO2
1)
2)
3)
4)
5)
double displacement
single displacement
double displacement
synthesis
decomposition
Part 2 - Bonding
Dissolving - the process in which molecules
interact and attract each other to form a solution
Dissolving depends on:
1. The molecules of the substance doing the
dissolving, called the solvent, and
2. The molecules of the substance being
dissolved, called the solute.
Solids, liquids, and gases can all dissolve.
The mutual attraction between water molecules
and other substances with positive and negative
charges causes these substances to dissolve.
SOLUTES, SOLVENTS, AND SOLUTIONS
Solution: homogeneous mixture of a solute
dissolved in a solvent
solute: component(s) present in smaller amount substance being dissolved
solvent: component present in greatest amount substance doing the dissolving
– Unless otherwise stated, assume the solvent is
water
Solubility - The extent to which a substance
dissolves
MOLECULAR BONDS
In a chemical reaction, bonds between atoms
are broken; the atoms rearrange, and then form
new bonds.
But how are atoms bonded together in the first
place?
The electrons (negative charge) and protons
(positive charge) from each atom are attracted
to the oppositely charged protons and electrons
in the other atom
ELECTRONS ARE THE GLUE THAT HOLDS GROUPS OF
ATOMS TOGETHER.
1. Ionic Bonding
- an atom feels a stronger attraction toward
another atom
- one or more electrons may leave that atom
and join the other
- the atom that lost an electron becomes a
positively charged ion
-the atom that gained an electron becomes a
negatively charged ion
The oppositely charged ions attract each other
and form an ionic bond.
2. COVALENT BONDING
- forms when one or more electrons from each
atom end up moving around the nuclei of both
atoms
- atoms share electrons
http://www.ider.herts.ac.uk/school/courseware/
materials/bonding.html
http://www.inquiryinaction.org/chemistryreview/
dissolving/
COVALENT = SHARING
TRANSFERING
IONIC =
http://people.seas.harvard.edu/~jones/es154/lec
tures/lecture_2/covalent_bond/bond_class.jpg
WACKY WATER
The molecular level-
http://sayee.ca/image/water%20molecule.gif
Hydrogen atom
- one proton (positive charge)
- one electron (negative charge)
Oxygen atom
- 8 protons
- 8 electrons
In Water: Protons = Electrons Thus, a water
molecule is neutral
It is the way a water molecule is bonded
together covalently that helps to make water a
good dissolver.
This means that the electrons from the oxygen
and hydrogens go around all three atoms
instead of just the atom they started with. To form
a covalent bond, the oxygen and hydrogen
atoms share several electrons.
But this sharing is not equal
WATER’S COVALENT BOND MAKES IT POLAR
-Oxygen has a stronger attraction for electrons
than the hydrogens do
- The electrons spend more time near the oxygen
than near the hydrogens
-the area near the oxygen is slightly negative
-the area near the hydrogens slightly positive
This makes water a polar molecule. A polar
molecule has no overall charge, but has an area
of positive charge separated from an area of
negative charge
WATER MOLECULES ARE ATTRACTED TO ONE
ANOTHER
http://www.physicalgeography.net/fundamental
s/images/water.JPG
The negative area of one water molecule is
attracted to the positive area of another, and
vice versa.
Attractions between the water molecules are
constantly breaking and forming with other water
molecules.
http://www.inquiryinaction.org/chemistryreview/
water/
Polar nature of water - makes water so good at
dissolving many substances.
SOLIDS DISSOLVING IN LIQUIDS
HTTP://WWW.INQUIRYINACTION.ORG/CHEMISTRYR
EVIEW/SOLIDS/
SALT dissolving in water
Salt (NaCl) on the molecular level:
a positive sodium ion (Na+) and a negative
chloride ion (Cl-)
Since positive and negative attract, the sodium
ion and the chloride ion form an ionic bond,
which results in NaCl.
When salt crystals are placed in water:
1. The positive ends of the water molecules
attract the negative chloride ions.
2. The negative ends of the water molecules
attract the positive sodium ions
3. When attraction between the water and the
ions overcomes the attraction the ions have for
each other, the salt dissolves.
SUGAR dissolving in water
Sucrose on the molecular level:
has positive and negative areas (polarity)- a
covalent bond
When sugar crystals are placed in water:
1. Water has positive and negative areas
(polarity)
2. Water molecules are attracted to the
oppositely charged area on the sucrose
molecules
3. When the attraction that water molecules
have for sucrose molecules overcomes the
attraction that sucrose molecules have for each
other, the water separates the sucrose molecules
from each other, and they dissolve.
So, Sugar and Salt dissolve differently because
they are bonded differently:
The positive and negative areas of water
molecules are attracted to
1. The oppositely charged ions in salt
2. The positive and negative areas on sugar
Because salt and sugar are made up of different
atoms that bond together differently, water is
attracted to them differently.
LIQUIDS DISSOLVING IN LIQUIDS
Just like solids, liquids can also dissolve in liquids
Alcohol - oxygen atom is bonded to a hydrogen
atom. Because of the characteristics of the O—H
bond, the oxygen is slightly negative and the
hydrogen is slightly positive
The mutual attractions between the water
molecules and the alcohol molecules cause the
alcohol to dissolve in water.
Corn syrup - Corn syrup is made up of glucose
and fructose molecules. These two molecules
also have O—H bonds and areas of positive and
negative charge.
The mutual attraction between the water
molecules and the corn syrup molecules cause
the corn syrup to dissolve in water.
Vegetable oil - Vegetable oil does not dissolve in
water because it is made of long chains of
carbon and hydrogen atoms bonded together.
These C—H bonds produce only a very small
amount of positive and negative charge.
There is little attraction between the molecules of
the oil and water, so the water does not dissolve
the vegetable oil.
GASES DISSOLVING IN LIQUIDS
Depends on the interaction between water
molecules and the gas molecules
CO2 - The positive and negative areas on water
molecules are attracted to the oppositely
charged areas of the carbon dioxide molecules.
This results in carbon dioxide gas (CO2) dissolving
in water (H2O) to make carbonated water
SOLUTIONS CAN BE SEPARATED
A solution is a mixture that can be separated by
physical means.
Some of the ways that mixtures can be
separated include evaporation, precipitation,
distillation, filtration and chromatography.
SOLUTIONS ARE MIXTURES
Remember:
A solution is a mixture of two or more substances
in which one or more of these substances
(solutes) are dissolved in another substance such
as water or another sort of liquid (solvent).
A solute can be a gas, solid or liquid.
CONCENTRATION OF THE SOLUTION: The amount
of the solutes compared to the solvent in a
solution
Dilute Solution: a solution containing a relatively
small quantity of solute as compared with the
amount of solvent
Concentrated Solution: a solution containing a
relatively large quantity of solute as compared
with the amount of solvent
SOLUBILITY OF A SOLUTION: The greatest
concentration of a solute in a solvent
Saturated: When the solvent contains the most
solute it can hold
Unsaturated: When the solvent contains less
solute than it can hold
SUSPENSIONS AND COLLOIDS
Suspension: like a solution, but the particles that
are in the suspension can be seen by the naked
eye
Colloid: a type of mixture in which one substance
is split up into tiny particles and spread
throughout another substance
Liquids, solids, and gases all may be mixed to
form colloidal dispersions.
Aerosols: solid or liquid particles in a gas.
Examples: Smoke is a solid in a gas. Fog is a liquid
in a gas.
Sols: solid particles in a liquid. Example: Milk of
Magnesia is a sol with solid magnesium hydroxide
in water.
Emulsions: liquid particles in liquid. Example:
Mayonnaise is oil in water.
Gels: liquids in solid. Examples: gelatin is protein in
water. Quicksand is sand in water.
Telling Them Apart
You can tell suspensions from colloids and
solutions because the components of
suspensions will eventually separate. Colloids can
be distinguished from solutions using the Tyndall
effect.
Tyndall effect - A beam of light passing through a
true solution, such as air, is not visible. Light
passing through a colloidal dispersion, such as
smoky or foggy air, will be reflected by the larger
particles and the light beam will be visible.
FACTORS THAT AFFECT RATE OF DISSOLVING AND
SOLUBILITY
Rate of Dissolving
1) Temperature
Rate increases with increase in temperature
Reason: At high temp, solvent molecules have
more kinetic energy → collide with undissolved
solid molecules more often thus dissolving
2) Size of substance
Rate increases by decreasing size of solute
Reason: surface area is increased (more contact
with solvent)
3) Stirring
→ increase with stirring or shaking
Reason: Brings fresh solvent into contact with
undissolved solid
FACTORS THAT AFFECT SOLUBILITY
1) Temperature
Solid: Solubility increases when temp increases
Reason: Energy is required to break tough bonds
in solid molecules (i.e.: ionic bond in salt).
Liquid: not much difference in solubility of when
temp increases
Reason: Bonds between particles in a liquid are
not as strong as bonds between particles in solid.
Not much happens when temp is increased.
Gas: solubility of decreases when temp
increases
Reason:When gas is dissolved in a liquid, it loses
some of its energy (slows down). At higher
temperature, gas picks up energy and thus picks
up speed and leaves solution (becomes less
soluble)
2) Size of Molecule
Solubility of solid and liquid increases when size of
molecule of solute decreases
3) Pressure and Solubility
Solid and liquid : not much difference in solubility
of with increased pressure
Gas: solubility increases with increased pressure
ie: pressure inside sealed bottle of pop is very
high → therefore, CO2 is dissolved in water.
When bottle is opened, pressure decreases, CO2
comes out of solution (that’s why you get
bubbling)
CONDUCTIVITY
Solutions composed of ions are conductors of
electricity.
Strong electrolyte: A solution with a large
concentration of ions
Weak electrolyte: A solution that has a small
amount of ions and is composed of mostly
molecular substances
Nonelectrolyte: A solution without ions that does
not conduct electricity
Strong acids, strong bases, and salts are all strong
electrolytes.
SOLUBILITY CURVES
A graph showing the concentration of a
substance in its saturated solution in a solvent as
a function of temperature.
Reading a Solubility Chart
1) The curve shows the # of grams of solute in a
saturated solution containing 100 mL or 100 g
of water at a certain temperature.
2) Any amount of solute below the line indicates
the solution is unsaturated at a certain
temperature
3) Any amount of solute above the line in which
all of the solute has dissolved shows the
solution is supersaturated.
4) If the amount of solute is above the line but
has not all dissolved, the solution is saturated
and the # grams of solute settled on the
bottom of the container = total # g in solution
– # g of a saturated solution at that
temperature. (according to the curve)
5) Solutes whose curves move upward w/
increased temperature are typically solids b/c
the solubility of solids increases w/ increased
temperature.
6) Solutes whose curves move downward w/
increased temperature are typically gases
b/c the solubility of gases decreases with
increased temperature.
Part 4 - Acids and Bases
Acids - a solution that has an excess of H+ ions. It
comes from "sharp" or "sour". The more H+ ions,
the more acidic the solution.
Properties of an Acid
- Tastes Sour
- Conduct Electricity
- Corrosive, which means they break down
certain substances. - Many acids can corrode
fabric, skin, and paper
- Some acids react strongly with metals
- Turns blue litmus paper red
Uses of Acids
- Acetic Acid = Vinegar
- Citric Acid = lemons, limes, & oranges. It is in
many sour candies such as lemonhead & sour
patch.
- Ascorbic acid = Vitamin C which your body
needs to function.
- Sulfuric acid is used in the production of
fertilizers, steel, paints, and plastics.
- Car batteries
Bases - a solution that has an excess of OH- ions.
- Another word for base is alkali.
- Bases are substances that can accept
hydrogen ions
Properties of a Base
- Feel Slippery
- Taste Bitter
- Corrosive
- Can conduct electricity. (Think alkaline
batteries.)
- Do not react with metals.
- Turns red litmus paper blue.
Uses of Bases
- Bases give soaps, ammonia, and many other
cleaning products some of their useful
properties.
- The OH- ions interact strongly with certain
substances, such as dirt and grease.
- Chalk and oven cleaner are examples of
familiar products that contain bases.
-) Your blood is a basic solution.
pH Scale - a measure of how acidic or basic a
solution is.
- The pH scale ranges from 0 to 14.
- Acidic solutions have pH values below 7
- A solution with a pH of 0 is very acidic.
- A solution with a pH of 7 is neutral.
- Pure water has a pH of 7.
- Basic solutions have pH values above 7.
Acid – Base Reactions
-A reaction between an acid and a base is
called neutralization. An acid- base mixture is not
as acidic or basic as the individual starting
solutions.