Matter
Chapter 2
1
Definition
• Matter is anything that
has mass and takes up
space.
• Mass is a measure of
the amount of matter.
• Volume is a measure of
the amount of space.
2
Types of Matter
• An element is a substance
that cannot be separated
or broken down into
simpler substances by
chemical means.
• An atom is the smallest
unit of an element that
maintains the properties
of that element.
3
Combinations of Matter
• A compound is a substance
made up of atoms of two
or more different elements
joined by chemical bonds.
• A molecule is the smallest
unit of a substance that
keeps all of the physical
and chemical properties
of that substance.
4
Chemical Formulas
• Chemical formula shows
how many atoms of each
element are in a unit of a
substance.
• Chemical symbols
represent the element.
• Subscripts indicate how
many atoms of each
element are in one
molecule of the
compound.
5
Mixtures
• Pure substances is matter
that has a fixed
composition and definite
properties.
• A mixture is a combination
of two or more substances
that are not chemically
combined.
6
Types of Mixtures
• Heterogeneous mixture
aren’t mixed uniformly
and are not evenly
distributed.
• Homogeneous mixture
are evenly distributed
and the mixture is the
same throughout.
7
What is used to define matter?
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1. mass
2. volume
3. Both mass and
volume.
4. Neither mass nor
volume.
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What is the smallest part of an
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What do we call a physical combination of two
or more substances?
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What is milk an example of?
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1. compound
2. Heterogeneous
mixture
3. Homogeneous
mixture
4. element
10
Physical Properties
• Physical Properties are
characteristics of a
substance that does not
involve a chemical change,
such as density, color or
hardness.
• Examples are melting
point, boiling point,
strength, hardness,
conductivity, magnetism,
heat.
12
Chemical Properties
• Chemical properties are
characteristics of matter
that describes a
substance's ability to
participate in chemical
reactions.
• Examples are reactivity
with oxygen, acid, water,
or other substances;
flammability.
13
Physical Changes
• Physical changes are
changes of matter from
one form to another
without a change in
chemical properties.
• Examples are cutting,
dissolving, mixing, melting,
evaporating, subliming.
14
Chemical Changes
• Chemical changes are
changes that occur
when a substance
changes composition by
forming one or more
new substances.
• Examples are burning,
rusting, and ripening of
fruit.
15
Triple Point of Carbon Dioxide
• Under certain conditions,
you can see dry ice as a
solid, liquid and gas all at
the same time.
• Dry ice under pressure will
exhibit all three states of
matter.
• Under normal conditions,
dry ice will sublime from a
solid to a gas.
16
Kinetic Theory
• All matter is made of atoms and molecules that act like tiny
particles.
• These tiny particles are always in motion. The higher the
temperature of the substance, the faster the particles
move.
• At the same temperature, more massive particles move
slower than less massive particles.
17
Different States of Matter
• Solids
– Definite volume.
– Definite shape.
– Low energy, molecules
close together.
• Liquids
– Definite volume.
– No definite shape.
– More energy, molecules
farther apart.
18
Different States of Matter
• Gases
– No definite volume.
– No definite shape.
– Lots of energy,
molecules are very far
apart.
• Plasma
– No definite shape.
– Particles have been
broken apart.
19
Energy in States of Matter
• Energy is the ability to
change or move matter.
• Thermal energy is the total
kinetic energy of the
particles that make up an
object.
20
Change in States of Matter
• Evaporation is the change
of a substance from a
liquid to a gas.
• Boiling point is the
evaporation of a liquid at a
certain temperature.
• Sublimation is the process
by which a solid turns
directly into a gas.
21
Changes is States of Matter
• Condensation is the
changes of state from a gas
to a liquid.
• Melting is the process of a
solid changing into a liquid.
• Freezing is the reverse of
melting.
22
Conservation of Mass & Energy
• The law of conservation of mass states that mass cannot be
created or destroyed.
• The law of conservation of energy states that energy
cannot be created or destroyed.
23
Fluids
• Buoyant force is the
upward force that fluids
exert on matter.
• All fluids exert pressure,
which is the amount of
force exerted on a given
area.
24
Archimedes’ Principle
• Archimedes Principle
states that the buoyant
force on an object in a
fluid is an upward force
equal to the weight of
the fluid that the object
displaces.
25
Fluid & Pressure
• Pressure is the amount
of force applied over an
area.
• Pressure=force/area
26
Pascal’s Principle
• Pascal’s principle states
that a change in
pressure at any point in
an enclosed fluid will be
transmitted equally to
all part of the fluid.
27
Fluids in motion
• Viscosity is a liquid’s
resistance to flow.
• Bernoulli’s principle
states that as the speed
of a moving fluid
increases, the pressure
of the moving fluid
decreases.
28
Properties of Gases
• Gases have no definite shape or volume, and
they expand to completely fill their container.
• Gas particles move rapidly in all directions.
• Gases are fluids.
29
Properties of Gases
• Gases have no definite shape or volume, and
they expand to completely fill their container.
• Gas particles move rapidly in all directions.
• Gases are fluids.
30
Properties of Gases
• Gas molecules are in constant motion, and
they frequently collide with one another and
with the walls of their container.
• Gases have a very low density because their
particles are so far apart. Because of this
property, gases are used to inflate tires and
balloons.
31
Properties of Gases
• Gases are compressible.
• Gases spread out easily and mix with one
another. Unlike solids and liquids, gases are
mostly empty space.
32
Gas Laws
• Gas laws explain the relationship between
volume, temperature and pressure of gases.
• Different Laws:
– Boyle’s Law
– Charles’ Law
– Gay-Lussac’s Law
33
Boyle’s Law
• Boyle’s Law states
that for a fixed
amount of gas at a
constant temperature,
the volume of a gas
increases as its
pressure decreases.
• P1V1=P2V2
34
Charles’ Law
• Charles Law states that
for a fixed amount of
gas at a constant
pressure, the volume of
the gas increases as its
temperature increases.
• K=V/T
35
Gay-Lussac’s Law
• Gay-Lussac’s Law states
that the pressure of a
gas increases as the
temperature increases
if the volume of the gas
does not change.
36
Atoms and the Periodic Table
Chapter 4
37
Atomic Structure
• Atoms are the simplest unit of
a substance that still maintain
the properties of the
substance.
• John Dalton proposed that
atoms could not be divided.
• Dalton also stated that atoms
of different elements could
join to form compounds.
38
Parts of an Atom
• Atoms are composed of
subatomic particles.
• Protons and neutrons are
found in the nucleus of
the atom.
• Electrons are found in
orbitals around the
nucleus of the atom.
39
Protons
• Protons have a positive
charge.
• Protons have the mass of one
amu (atomic mass unit).
• The number of protons in an
atom is characteristic of that
element.
• Each element has different
number of protons.
40
Neutrons
• Neutrons have no charge.
• Neutrons have the mass
of one amu.
• Neutrons add to the mass
of an atom.
• Atoms of the same
element with different
number of neutrons have
different masses
(isotopes).
41
Electrons
• Electrons have a negative
charge.
• Electrons have the mass of
1/2000 the mass of a proton.
• The mass of an electron is
insignificant.
• The outer most electrons
determine how elements
combined in chemical
compounds.
42
Quarks
•
A quark (IPA: /kwɔrk/) is a generic
type of physical particle that forms
one of the two basic constituents of
matter, the other being the lepton.
• Various species of quarks
combine in specific ways to form
protons and neutrons, in each
case taking exactly three quarks
to make the composite particle in
question.
43
Models of Atoms
• Niels Bohr suggested that
electrons in an atom move
in set paths around the
nucleus.
• Electrons can only be in
certain energy levels.
• Number of electrons=2n2
44
Electron Cloud Model
• This model suggest that
electrons orbit the
nucleus in a cloud.
• The regions in an atom
where electrons are likely
to be found are called
orbitals.
• The four different kinds of
orbitals are the s, p, d and
f orbitals.
45
Valence Electrons
• An electron in the
outermost energy level of
an atom is called a
valence electron.
• Valence electrons
determine an atom’s
chemical properties and
its ability to form bonds.
46
Periodic Table
• The Periodic Law states
that when elements are
arranged this way,
similarities in their
properties will occur in a
regular pattern.
47
Structure of the Periodic Table
• Horizontal rows in the
periodic table are called
periods.
• Atoms of elements in the
same group, or column,
have the same number of
valence electrons, so
these elements have
similar properties.
48
Ions Formation
• Atoms that gain or lose
electrons form ions.
• Elements that lose
electrons have a positive
charge (cation).
• Elements that gain
electrons have a negative
charge (anion).
49
Periodic Information
• The atomic number is the
number of protons in an
element.
• The mass number of an
atom equals the number
of protons plus the
number of neutrons.
50
Isotopes
• Isotopes are atoms of the
same element that have
different number of
neutrons.
• Isotopes of the same
element have different
atomic masses.
• The three isotopes of
hydrogen are protium,
deuterium & tritium.
51
Mass of Atoms
• An atomic mass unit (amu) is
equal to one-twelfth of the
mass of a carbon 12 atom.
• The average atomic mass for
an element is a weighted
average, so the more
commonly found isotopes
have a greater effect on the
average than rare isotopes.
52
Squares on the Periodic Table
• The chemical symbol is
abbreviation of the
chemical name.
– The first letter of the
chemical symbol is
capitalized.
– If there is more than one
letter, the other letters are
lower case.
• Atomic number.
• Average atomic mass.
53
Classifications of Elements
• Metals are on the left side
of the periodic table.
• Nonmetals are on the
right side of the periodic
table.
• Metalloids are located
between the metals and
nonmetals.
54
Alkali Metals
• This is the most reactive
group of metals.
• Has only 1 valence
electrons.
• Reacts violently with
water.
• Elements include Li, Na,
K, Rb, Cs & Fr.
55
Alkaline Earth Metals
• Not as reactive as Alkali
Metals.
• Has two valence
electrons.
• Members include Be, Mg,
Ca, Sr, Ba & Ra.
56
Transition Metals
• Members have a wide
variety of properties.
• Has 1, 2 or 3 valence
electrons.
• Contain the coin metals:
gold, silver & copper.
• Contain the iron triad of
Fe, Co & Ni.
57
Halogens
• The most reactive group
of nonmetals.
• Has 7 valence electrons.
• All members are
poisonous.
• Members include F, Cl, Br,
I & At.
58
Noble Gases
• Members contain a stable
octet with 8 valence
electrons.
• These elements do not
react with other
elements. Don’t form
compounds.
• Members include He, Ne,
Ar, Kr, Xe & Rn.
59
Semiconductors
• Have the properties of
both metals & nonmetals.
• Used in the electronic
industry.
60
Using Moles to Count Atoms
• A mole is a collection of a
very large number of
particles.
• Avagadro’s constant is the
number of particles in a
mole of a pure substance.
• Avagadro’s constant is
6.022 x 1023/mol.
61
Molar Mass
• Molar mass is the mass in
grams of 1 mol of a
substance.
• The molar mass of an
element in grams is the
same as its average
atomic mass in amu on
the periodic table.
62
Converting Moles to Grams
• Amount(mol) x molar mass of element/1 mol of element = mass(g)
• Amount(mol)=1 mole of element/molar mass of element x mass(g)
63
Converting Amount to Mass
• Determine the mass in grams of 5.50 mol of
iron.
• Given: amount of iron=5.50 mol mol Fe
molar mass of iron=55.85 g/mol Fe
• Unknown: mass of iron=?g Fe
• 55.85 g Fe/1 mol Fe
• 5.50 mol Fe x 55.85 g Fe/1 mol Fe=307 g Fe
64
Converting Amount to Mass
• Determine the mass in grams of 5.50 mol of
iron.
• Given: amount of iron=5.50 mol mol Fe
molar mass of iron=55.85 g/mol Fe
• Unknown: mass of iron=?g Fe
• 55.85 g Fe/1 mol Fe
• 5.50 mol Fe x 55.85 g Fe/1 mol Fe=307 g Fe
65
The Structure of Matter
Chapter 5
66
Compounds
• Compounds are made from
two or more elements.
• The compound has
properties that are different
from those of the elements
that make it.
• Compounds always have
the same chemical formula.
67
Chemical Bonds
• The attractive forces that
hold different atoms or ions
together in compounds are
called chemical bonds.
• A bond length gives the
distance between the nuclei
of the two bonded atoms.
• Bond angles tell how these
atoms are oriented in space.
68
Structural Formulas
• Structural formulas can
show how the atoms are
arranged in a compound.
69
What do we call the force that holds
atoms in a molecule together?
Atoms
Chemical bond
Nucleus
Valence electrons
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The properties of a compound are
different than the properties of the
elements that make up the compound.
1. True
2. False
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71
Types of Bonds
• Ionic bonds are formed
between oppositely charged
ions.
• Metallic bonds are formed
between atom’s nucleus
and a neighboring atom’s
electrons.
• Covalent bonds are made
between atoms of
nonmetals.
72
Ionic Bond
• Bond between a metal ion
and a nonmetal ion.
• The oppositely charged ions
are attracted to each other.
• Forms ionic solids.
• High melting point.
• Conduct electric current.
73
Metallic Bond
• Bond that occur between
atoms of metal.
• Compounds are malleable
and ductile.
• Conduct electricity & heat.
• Mixed to form alloys.
74
Covalent Bond
• Bond formed between two
nonmetals.
• Share electrons.
• Lower melting and boiling
points than ionic
compounds.
• Covalent compounds are
soft and squishy.
• Covalent compounds tend
to be more flammable than
ionic compounds.
75
Polyatomic Ions
• Polyatomic ions are groups
of covalently bonded atoms
that have either lost or
gained electrons.
• The group of atoms act as a
single ion.
• Polyatomic ions are put in
parenthesis when you need
to have more than one in a
ion.
HCO3−
76
Writing Chemical Formulas for Ionic
Compounds
• List the symbols for each ions.
• Write the symbols for the ions
with the cation first.
• Find the least common multiple
of the ions’ charges.
• Write the chemical formula,
indicating with subscripts how
many of each ion are needed to
make a neutral compound.
77
Naming Covalent Compounds
• Name the element that is
farthest to the left on the
periodic table.
• Add a prefix if there is more
than one atom.
• Name the element that is
farthest to the right on the
periodic table and change
the ending to an –ide.
78
Writing Chemical Formulas for Covalent
Compounds.
• List the chemical symbols of
the elements.
• Add subscripts to indicate
how many atoms of each
element there is in a
molecule of the compound.
BaF2
79
Empirical vs. Molecular Formulas
• Empirical formula shows the
smallest whole-number
ratio of atoms that are in a
compound.
• Molecular formula tells how
many atoms are in one
molecule of the compound.
C6H12O6
80
Organic Compound
• Organic compound is a covalently
bonded compound made of
molecules.
• Organic compounds contain
carbon and most of the time
hydrogen.
• Carbon can form four covalent
bonds.
• Compounds made of only
hydrogen & carbon atoms are
called hydrocarbons.
81
Types of Organic Compounds
• Alkanes only have single
covalent bonds.
• Alkenes have double
covalent bonds.
• Alcohols have hydroxly or –
OH group attached.
82
Alkanes
•
•
•
•
•
•
•
•
•
•
Methane
Ethane
Propane
Butane
Pentane
Hexane
Heptane
Octane
Nonane
Decane
83
Alkenes
•
•
•
•
•
•
•
•
•
•
Methene
Ethene
Propene
Butene
Pentene
Hexene
Heptene
Octene
Nonene
Decene
84
Alcohols
•
•
•
•
•
•
•
•
•
•
Methanol
Ethanol
Propanol
Butanol
Pentanol
Hexanol
Heptanol
Octanol
Nonanol
Decanol
85
Polymers
• Polymers are large molecule
that is formed by more than
five monomers or small
units.
• Polymers can be natural or
synthetic.
• Polymers elasticity varies
depending on the structure.
86
Biochemical Compounds
• Biochemical compounds are
naturally occurring organic
compounds that are very
important to living things.
• Type of biochemical
compounds:
– Carbohydrates (simple sugars)
– Proteins (amino acids)
– Lipids (fatty acids)
87
Chemical Reactions
8 West Science Investigations
Types of Chemical Reactions
• The left side of a
chemical equations
are the reactants.
• The right side of a
chemical equation
are the products.
• The arrow is the yield
sign.
CH4 + 2 O2 → CO2 + 2 H2O
Energy in a Chemical Reactions
• Chemical energy is the
energy that comes
from the bonds
between two atoms.
• Exothermic reaction
releases energy during
the reaction.
• Endothermic reaction
absorbs energy during
the reaction.
Reaction Types
• Synthesis reaction
involves combining
substances to
produce a new
substance.
• A + B → AB
2Na + Cl2 → 2 NaCl (formation of table salt)
S + O2 → SO2 (formation of sulfur dioxide)
4 Fe + 3 O2 → 2 Fe2O3 (iron rusting)
Reaction Types
• Decomposition
reactions involves
substances that are
broken apart.
• AB → A + B
2H2O2 → 2H2O + O2
H2CO3 → H2O + CO2
CaCO3 → CaO + CO2
2KClO3 → 2KCl + 3O2
Reaction Types
• Combustion
reactions involves
reactions that uses
oxygen.
• Oxygen is always in
one of the reactants
of the reaction.
CH4 + 2O2 → CO2 + 2H2O + energy
CH2S + 6F2 → CF4 + 2HF + SF6
Reaction Types
• Single displacement
reaction involves atoms
of one element that
2AgNO3(aq) + Zn(s) → 2Ag(s) + Zn(NO3)2(aq)
move into a compound,
and atoms of the other
Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g)
element appear to move
out.
• A + BX → AX + B
Reaction Types
• Double displacement
reaction involves two
compounds that
appear to exchange
ions.
• AX + BY → AY + BX
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)
AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3
(aq)
Reaction Types
• Oxidation-reduction
reaction involves one
atom losing electrons
(oxidized) and another
atom gains electrons
(reduced).
• Also called redox
reaction.
Balancing Chemical Equations
1. Identify the reactants and
products.
2. Write a word equation for
the reaction.
3. Write the equation using
formulas for the elements
and compounds in the
word equation.
4. Balance the equation one
element at a time by
adding coefficients before
each chemical formula.
Rates of Change
• Most reactions go faster
at high temperatures.
• A larger surface area
speeds up reactions.
• Concentrated solutions
react faster.
• Reactions are faster at
higher pressure.
• Massive , bulky
molecules react slower.
Catalyst & Inhibitors
• Catalysts are added to
reactions to make them
go faster.
• Inhibitors are catalysts
that cause a reaction to
go slower.
Chemical Equilibrium
• Most chemical
reactions are not
reversible.
• Some changes are
reversible.
• Systems in equilibrium
respond to minimize
change.
Le Chatelier’s Principle
• If a change is made to a
system in chemical
equilibrium, the
equilibrium shifts to
oppose the change until
a new equilibrium is
reached.
Homework
• Chapter 6 Review, page
214-215, questions 128.
Mixtures
• Heterogeneous
mixtures is composed
of dissimilar
components.
• Homogeneous mixtures
describes something
that has a uniform
structure or
composition
throughout.
Types of Mixtures
• Suspension is a mixture,
that separates upon
standing.
• Colloid is a mixture
where the particles are
permanently suspended
and never separate.
• Emulsion is a colloid
made of liquids that
usually never mix.
Solution
• Solution is a mixture
where something
dissolve in something.
• Solute is the part of a
solution that is being
dissolved.
• Solvent is the part of a
solution that is doing
the dissolving.
Factors that Affect the Rate of
Dissolving.
• Stirring
• Surface area
• Temperature
Solubility
• Concentration is the
quantity of solute
dissolved in the given
volume of solvent.
• Unsaturated solution
contains less than the
maximum amount of
solute in the solvent
• Saturated solution can
not dissolve any more
solute.
• Supersaturated solution
holds more solute than
it normally can.
Molarity
• Concentrations can be
expressed in molarity.
• Molarity is expressed as
moles per liter.
• Molarity = moles of
solute ÷ liters of
solution.
• M=mol ÷ L
Acid & Bases
Acids
• Acid are substances
that contain hydronium
ions.
• Properties:
– High concentration of
hydronium ions.
– Sour taste
– Turns blue litmus red
– Reacts with metals
Bases
• Bases are substances
that contain hydroxide
ions.
• Properties:
– High concentrations of
hydroxide ions.
– Bitter taste.
– Turn red litmus blue.
– React with animal
matter.
pH
• pH is the measure of
the concentration of
hydronium ions.
• pH below seven is an
acid.
• pH above seven is a
base.
• pH of seven is neutral.
Neutralization Reaction
• Neutralization reaction
occurs when an acid
and a base are
combined.
• The reactants are acid
and base.
• The products are water
and salt.
• HCl+NaOH → H2O+NaCl
Indicators
• Indicators are chemicals
that indicate the pH
range of a solution
• Examples:
–
–
–
–
–
Phenolphthalein
Litmus
pH paper
Phenol red
Methylene blue
Acids vs. Bases
• Acids dissolve in water
ionize to form ions.
• Bases dissolve in water
dissociates to separate
hydroxide ions.
Task of the day.
• Close your notes.
• Write an essay
comparing acid & bases
based on what you
know.
• Minimum of 127.5
words!