Transcript Section 2.1 - Harlan Christian School

Nature of Matter
SECTION 2.1
Matter
 Defined as anything that occupies space and
has mass
 Everything we see around us is matter
 Soil
 Air
 Water
 Rocks
 Plants
 Animals
States of Matter
 Kinetic Energy—energy of motion; all
particles of matter are in constant motion
and therefore have kinetic energy
 The temperature of matter is an indirect
measurement of the particles’ average kinetic
energy
 At low temperatures, the particles have very
little kinetic energy and move very little
 At high temperatures, the particles can have
enough kinetic energy to fly about at
thousands of kilometers per hour
States of Matter
 Solid
 Has a definite shape
 A definite volume
 Very difficult to compress
 Particles are packed close together and held
rigidly
States of Matter
 Liquid
 Has the shape of its container
 Has a definite volume
 Difficult to compress
 Particles close together but free to move
States of Matter
 Gas
 Has no definite shape
 Has no definite volume
 Relatively easy to compress
 Particles far apart
 Plasma
 A hot gas in which atoms are partially broken
down to form charged particles, or ions
States of Matter
 When the substance reaches a certain
temperature, called the melting point, the
particles have enough energy to break loose
from their rigid positions and form a liquid
 When a liquid is heated to the boiling point,
the particles gain enough kinetic energy to
break away from each other forming a gas
Classification of Matter
 Substance: a form of matter with unique
properties that make it different from every
other substance
 Examples: Water, copper, baking soda, table
salt, aluminum, etc.
 Atoms: the smallest particles of an element
(not the smallest particles of matter)
 Elements: substances that cannot be broken
down into simpler substances by ordinary
chemical means
 Examples:
 Oxygen, hydrogen, gold, iron, lead, potassium,
nitrogen, silicon (other elements listed in the
periodic table of the elements)
 Not all elements are found as single atoms
 Molecules: groups of two or more atoms that
are linked by chemical bonds to form distinct
units
 These two-atoms elements are called
diatomic elements; other elements exist as
larger molecules
 Examples of diatomic elements: hydrogen,
oxygen, nitrogen, flourine, chlorine, bromine, and
iodine
Atomic Symbols
 Memorize tables 2.2 and 2.3
Abundance of Elements
 118 elements identified thus far
 88 occur naturally in significant amounts
 More than 99% of the earth’s crust, oceans,
and atmosphere is made up of only eight
elements
 Oxygen-the most abundant element in the
earth’s crust
 Iron-the most abundant element in the earth
as a whole
 Hydrogen-the most abundant element in the
universe
Compounds
 A substance that can be decomposed into
simpler substances
 Examples: water (H2O) and table salt (NaCl),
 A compound has properties quite different
from those of the elements from which it is
formed
 Example: hydrogen-highly flammable;
oxygen-flammable; H2O-commonly used to
put out fires
Formula
 A molecule of a compound is composed of
different types of atoms linked together
 The compound may be represented by a
formula-a grouping of symbols that tells what
types of atoms compose the compounds and
the number of each type of atom in one
molecule of the compound
 Ex. H2O; CO; CO2; C6H12O6
Pure Substances
 Elements and compounds
 Definite composition (relative amounts of
each element in a given compound are
unchangeable, no matter where the
compound may be found)
 Homogeneous (same kind)-composed of the
same kind of matter throughout the sample
Mixtures
 A substance consisting of two or more pure
substances that are physically mixed but not
chemically combined
 Each component retains its own properties
 Each can be present in various proportions
and can be separated from the mixture by
physical means
Homogenous Mixtures
 Solutions; same throughout
 Examples: air, vinegar, sweetened coffee,
steel
Heterogeneous Mixtures
 Consists of pure substances that are
incompletely mixed
 Each component in a heterogeneous mixture
forms a distinct phase—a homogeneous part
of a system that is in contact with but
physically distinct from other parts of the
system
 Examples:
Dalton’s Atomic Theory
 1. Every element consists of tiny, indivisible,
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indestructible particles called atoms
2. All of the atoms of a particular element have
the same size, mass, and chemical behavior
3. Differences in properties of elements result
from differences in the atoms of the elements
4. The atoms of the elements combined in a
compound are combined in a definite ratio
5. A chemical reaction is the result of
rearrangement, combination, or separation of
atoms
The Law of Definite
Composition
 States that the ratios of the masses of each
element in a given compound are always the
same
 Example on pg. 22
The Law of Multiple
Proportions
 States that when two elements can combine
to form more than one compound, the
masses of one element that combine with a
fixed amount of the other element are in a
ratio of small whole numbers
 Example on pg. 22
Section 2.2: Properties and
Changes of Matter
 One type of matter may be described,
identified, or distinguished from another by
its properties.
Physical Properties
 Properties that describe a substance’s
appearance
 Include: color, odor, density, hardness,
solubility, taste, state, boiling point, and
melting point
 Can be measured without changing the
identity or composition of the substance
Chemical Properties
 Properties that describe how matter reacts to
change into other chemically different
substances having different properties
 That iron rusts easily in moist air and that
propane burns readily are chemical properties
of these substances
 Can only be determined through chemical
reactions, which change the identity of a
substance
Physical Changes
 Changes in the physical appearance of matter that
do not change the identity or chemical
composition of a substance
 Physical changes may often be reversed by
physical processes
 Change of state, conversion of heterogeneous to
homogeneous matter, metal hammered into a
sheet, tearing a sheet of paper, crushing a large
rock into gravel, cutting your hair, etc.
 No change in chemical identity results, only
changes in shape or size
Chemical Changes
 A change in which a substance becomes a
different substance with a different
composition and properties
 Sodium–a shiny, soft, poisonous metal that
reacts violently with water
 Chlorine—a greenish, poisonous gas with an
irritating odor
 When sodium and chlorine react chemically, a
white crystalline solid with a salty taste is
formed, namely sodium chloride
Chemical Changes
 Other examples:
 Rusting of iron
 The digestion of food
 The grilling of steaks
 The burning of gasoline
 The cooking of eggs
 The burning of a match
Chemical Changes
 Chemical changes can be reversed only by
other chemical changes
 Not all chemical changes can be reversed
 The cooking of an egg
Chemical Changes
 A number of observations can be made that
will suggest that a chemical change has
occurred
 Formation of a gas
 Formation of a precipitate (an insoluble
substance)
 Liberation or absorption of heat, light, or some
other form of energy
 A distinct change in color
Separation of Mixtures
 Distillation—used to separate homogeneous
mixtures in liquid form if the components
have widely differing boiling points
 Simple Distillation
 When a mixture has components with boiling
points that are close together, fractional
distillation is used
Separation of Mixtures
 Fractional Crystallization
 Used to purify a solid containing a relatively small
amount of solid impurity
 Chromatography
 Used to separate complex mixtures
 Paper Chromatography
 Column Chromatography
 Gas Chromatography
Section 2.3: Subatomic
Particles
 Subatomic Particles
 Protons, neutrons, and electrons
 Electrons—has a negative charge
 Proton—has a positive charge, located in the
nucleus
 Neutron—has no charge, located in the
nucleus
 Rest of the notes in this section will come
from the videos
Section 2.4: Atomic Number,
Mass Number, Isotopes, and Ions
 The number of protons in the nucleus
determines the identity of an element
 Atomic number—the number of protons in
the nucleus
 Denoted by the symbol Z
 All atoms of the same element have the same
number of protons and hence the same
atomic number, Z
Mass Number
 Denoted by the symbol A
 The sum of the number of protons and the
number of neutrons in a nucleus
 A=Z+N
 Nuclear particles, whether protons or
neutrons, are known collectively as nucleons
 The mass number is not really a mass; it is a
number of nucleons and must be a whole
number
Isotopes
 Atoms of the same element that differ in
their mass numbers
 Examples:
 Pg. 31
Ions
 Atoms or molecules that have a different number
of electrons than protons
 An ion with more electrons than protons is called
and anion
 Has a negative charge
 An ion with fewer electrons than protons is called
a cation
 Has a positive charge
 In neutral atoms, there are equal numbers of
electrons and protons
Examples: pg. 32
 Example 2.1
 Application 1-3
Section 2.5: Atomic Mass
 An atom’s atomic mass measured in atomic
mass units approximately equals the atom’s
mass number
 Atomic mass units denoted a u
 Mass Spectrometer
 Used to measure atomic masses with high
precision
 Mass Spectrum
 A record of the mass distribution of particles in a
sample
Average Atomic Mass
 Calculated by finding the weighted average
 Example: pg. 34
 Track team
 Example 2.2 pg. 35
Chapter 2
 Homework and Groupwork will be due Friday
 They will be handed out at the beginning of class
on Tuesday
 Chapter 2 Quiz on Friday