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Unit 2
An Introduction to Chemistry
Chapter 1 Section 1
What is Matter?
• Anything that takes up space (volume)
• And has mass.
• Matter can be living or non-living.
– People, plants, animals, air, water, rocks are
examples.
Sound, light, gravity, thoughts are not matter.
What makes matter different?
• All matter has general properties of mass,
volume, density along with other properties.
• We can also describe matter using our senses
or by how matter reacts.
– Physical properties
– Chemical properties
What is volume?
• Volume is the amount of space that an object
occupies.
– Volume of irregularly shaped objects and liquids
are measured in mL or liters.
• Water displacement can be used to determine a solids
shape.
– Volume of a regular solid object is measured in
cubic centimeters.
Mass vs Weight
• Mass refers to the amount of matter that
makes up an object.
– Mass is measured in grams.
• Weight is the gravitational pull upon that
object, which varies depending on the force of
gravity.
– Weight is measured in Newtons.
Mass and Inertia
• Inertia is the resistance of an object to change
its motion.
• Objects with mass resist change.
• The more massive the object the more inertia
the object has and a greater force will need to
be applied to move the object.
Chapter 1 Section 2
Physical Properties
In general, what physical properties does
all matter have?
• Physical properties are characteristics of a
substance that can be observed without
changing the identity of the substance.
– Mass
-Color
– Weight -Size
– Inertia
-Shape
– Volume
-Texture
– Density-Phase (State)
What is density?
• Density is the amount of matter in a given
volume of an object.
• Density allows us to compare different types
of matter.
• All matter has density and the same type of
matter has the same density.
• Density= Mass/ Volume
Chapter 1 Section 3
Chemical Properties
What are the chemical properties of
matter?
• Chemical properties describe a substances
ability to change into another substance.
• Chemical reactions take place.
• Examples: Flammability, Reacting with acids,
Oxidation, Reacting with electricity.
Physical vs. Chemical Change
What are physical changes?
• A physical change is the process that alters the
physical properties of the substance and no
new substance is formed.
• Physical changes include phase changes.
What are chemical changes?
• Chemical change is the process by which a
substance becomes a new substance.
• Burning, rusting, corroding, tarnishing result
in a chemical change.
Chapter 2 Section 1
Three States of Matter
What are the Phases (states) of Matter
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•
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Solid
Liquid
Gas
Plasma
Liquid
Solid
Gas
What is a solid?
• Solids have a definite shape and volume.
• Particles that make up an solid are closely
packed together and only vibrate in place
– If they are arranged in regular repeating patterns
they are called-crystals or cyrstalline solids
• Some solids can lose their shape over a long
period of time-candle wax, glass
– Amorphous solids
What is a liquid?
• Liquids have no definite shape but have a
definite volume.
– Take the shape of the container
– Have surface tension
• Means it acts at the surface of the liquid (causes liquids
to form spherical drops).
• Ex: gasoline has low tension and forms flat drops
– Particles are free to move around each other
• The slower they move is the viscosity of the liquid. For
example, honey vs. Water
What is a Gas?
• A gas has no definite shape or volume.
• Particles are in no set pattern, and are far apart.
Particles are moving constantly and collide into the
sides of the container.
– Helium is a gas.
– Empty space between gas particles can change.
• Ex: The particles of helium in the balloons are farther
apart than the particles of helium in the tank.
What is Plasma?
• Plasma is the most common phase of matter
in the universe.
• Plasma is a very high energy state.
• Stars and the Sun are made of matter in their
plasma phase.
Chapter 2 Section 2
Behavior of Gases
What factors affect how a gas
behaves?
• Temperature
– Measures how fast the particles are moving in an
object. Warmer something is, the faster the particles
move.
• Volume
– Amount of space that an object takes up. Gas will fill
the shape of an object.
• Pressure
– Amount of force exerted on a given area of surface.
Increased force=Increased pressure
Gas Behavior Laws
• Boyles Law
• Charles Law
What does Boyle’s Law say?
• Volume of a fixed amount of gas varies
inversely with the pressure of the gas
• In other words, if you increase the pressure
the volume decreases.
What does Charles Law tell us?
• The volume of a fixed amount of gas varies directly
with the temperature of the gas.
• In other words, an increase in temperature will
increase the volume
• Boyle’s & Charle’s Law are called the gas laws which
describe the behaviors of gases as temperature,
pressure and volume change.
Chapter 2 Section 3
Changes of State
What are phase changes?
• As energy is absorbed or released by a
substance the particles move faster or slower.
• When particles have sufficient energy they will
go from one phase to another.
Types of phase changes include:
– Melting: solid to liquid
– Freezing: liquid to solid
– Evaporation: liquid to gas
– Condensation: gas to liquid
– Sublimation: solid to gas
Liquid
Solid
Gas
Chapter 3 Section 1
Elements
Elements
• The simplest pure substance that cannot be
broken down any farther.
• Cannot be changed by heating or chemical
reactions.
• An atom is the smallest particle of an element
that exhibits the properties of the element.
• “Building Block of matter”
Examples: Cobalt, Iron, Nickel
Chemical Symbol
• Shorthand way to represent an element
– Consists of 1 or 2 letters usually taken from the
name of the element
• 1st letter is ALWAYS CAPITILIZED
• 2nd letter is always lower case used when the first
letter has already been used
– Sometimes the symbol is based on the Latin term
for the element
Pure Substances
• When all particles are alike or made of one
kind of material
• A pure substance has definite and unique
identifiable properties.
The atoms of the element iron are alike whether they are in a meteorite or in a
common iron skillet.
Properties of Elements
• Characteristic properties include some
physical properties, such as boiling point,
melting point, and density and a chemical
property, such as reactivity with acid.
• These properties allow you to tell the
elements apart.
• Elements can share a property, but will never
have all of the exact properties as another
element.
Classifying Elements by Their Properties
• Categories of Elements
– Metals
• high luster, high density, ductile, malleable, good conductors,
high melting points, tend to lose electrons
– Metalloids
• Properties of metals and nonmetals
– Nonmetals
• little luster, brittle, lower densities and melting points, tend
to gain electrons
 By knowing the category of an unfamiliar
element, you can predict some of its properties.
Chapter 3 Section 2
Compounds
Compounds
• Pure substances that can be broken down into
simpler substances
– The simpler substances can only be separated by a
chemical change (heat/electricity)
– made of more than one type of element that are
chemically bonded
• A molecule is the smallest particle of a
compound that exhibits all the properties of
the compound.
Example: Sodium
+
Chloride
Sodium Chloride
Ratio of Elements
• Elements do not randomly join to form
compounds.
– Join according to the mass
– Ex: hydrogen=mass of 1 & oxygen=mass of 8
Ratio would be 1:8
Chapter 3 Section 3
Mixtures
What is a mixture?
• Matter is two or more substances mixed together
but not chemically combined.
• Substances keep most of their identities and own
properties
• Vary in amounts
• Can be separated by physical means based on the
physical properties of the various parts of the
mixture
• No chemical change will be found in a mixture.
Heterogeneous vs Homogeneous
Mixtures
• Heterogeneous mixtures do not appear the
same throughout and can be easily
separated. (least mixed)
– ie trail mix, salad dressing, pizza, granite
• Homogeneous mixtures appear the same
throughout and are harder to separate.
(well mixed)
– ie lemonade, jello,
Common Ways to Separate Mixtures
• Distillation-separates a mixture based on the
boiling pts. of the components.
• Magnet
• Centrifuge-separates according to the
densities of the components.
• Evaporation
• Filtration
Solutions
• Solutions are homogeneous mixtures formed
when one substance dissolves in another.
(best mixed)
• Particles are evenly spread out and cannot be
separated by physical means.
Parts of a Solution
– Solute: the substance that is dissolved
– Solvent: the substance that does the dissolving
• Universal solvent: Water
• Concentrated vs dilute
– A dilute solution contains less solute than the
concentrated solution.
• Solubility is the ability for a substance to dissolve
in a solvent at a specific temperature, in general
as temperature increases so does the solubility.
Suspensions & Colloids
• Suspension
– A mixture in which particles of a material are
more or less evenly dispersed throughout a liquid
or gas.
– Ex: Snow Globe
• Colliod
– A mixture in which the particles are dispersed
throughout but are not heavy enough to settle
out.
– Ex: Milk, mayo, stick deodorant
Chapter 4 Section 1
Development of the Atomic Theory
Models of the atom
• Since the individual atom can not be seen,
scientist have tried to describe the atom using
models based on observations (indirect
evidence).
• The first description of what an individual
particle of matter looks like was made by
Democritis, a greek philosopher over 2000
years ago.
Who has contributed to what we
believe the atom looks like?
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Democritis
John Dalton
J.J. Thomson
Ernest Rutherford
Neils Bohr
Chadwick
Schrodinger
What did Democritis say?
• He theorized that at some point you would no
longer be able to get a smaller particle and
the smallest particle called the atom would
be:
– small, hard spheres made of the same material
that are constantly moving
– that different materials came in different sizes and
shapes
– infinite in number and capable of joining together
(uncuttable).
Who was John Dalton ?
• English chemist of the 1800s.
• Studied the weather and the composition of
the air.
• Performed experiments on decomposition of
water into hydrogen and oxygen
• Developed Atomic theory in 1803 which
became the foundation of Modern Chemistry.
Who was Thomson and what did he
say?
• English scientist that first proved there were
particles smaller than the atom in 1897.
• Discovered the electrons or negatively charged
subatomic particles.
• Called them “corpuscles” and reasoned that if
there were negatively charged particles that
there must be positively charged particles as well.
• He conducted the cathode-ray experiment
(page 84).
Thomson’s Plum Pudding Model
• Stated that the atom was made of a pudding
like positively charged material throughout
which negatively charged particles were
randomly scattered.
– Named after a popular desert in Thomson’s day.
– Today you would call it the chocolate chip ice
cream model.
Who was Ernest Rutherford?
• English Physicist who worked with Thomson.
• In 1908, he performed his famous gold foil
experiment (page 85).
• In 1911, he revised the atomic theory. In the
center of the atom is a tiny, extremely dense,
positively charged part called the nucleus.
Rutherford’s Model
• His model of the atom stated that all the
positively charged particles were in the center
of the atom, called the nucleus with the
negatively charged particles scattered around
the edge of the atom.
• Most of the atom was empty space.
• Most of the atom’s mass is in the nucleus.
Who was Neils Bohr?
• Danish Scientist
• In 1913, he refined Rutherford’s model
• Identified, various energy levels of electrons.
Electrons jump between levels from path to
path.
What did Bohr say?
• Electrons move in definite orbits around the
nucleus much like the planets around the sun.
• These orbits, or energy levels, are located at
certain distances from the nucleus.
– 1st level holds 2 electrons
– 2nd level holds 8 electrons
– 3rd level holds 18 electrons
Who was Chadwick?
• In 1932, he discovered the neutron.
Who was Schrodinger?
• He is one of a group of scientists that developed a
model of the atom based on the principles of wave
mechanics (mathematics).
• Stated it was impossible to locate the exact position
of an electron but could find the probable location
based on the energy and called the location
“orbitals”
• Electrons are found in electron clouds, not paths.
Modern Atomic Theory
• States that the atom has a small positively
charged nucleus surrounded by a large region
in which there are enough electrons to make
the atom neutral.
• Theory is changing
– evidence supporting even smaller particles called
“Quarks” that make up the three subatomic
particles
• quarks are made up of different “flavors” and “colors”
Chapter 4 Section 2
The Atom
What are atoms made up of?
• Protons-positively charged particles in the
nucleus. 1 proton=1 amu
• Neutrons-particles of the nucleus that have no
charge. 1 neutron=1 amu
• Electrons-negatively charged particles found
outside the nucleus in the electron cloud.
– Electrons with the lowest energy are closest to the
nucleus.
How do we tell the difference
between different atoms?
• Atomic number = # of protons
• Atomic number identifies the element and
does not change for that element.
• Indicates the number of valence electrons in
the highest energy level that are available to
chemically bond.
What is an isotope?
• Isotopes are the same element that vary in
the number of neutrons they have in the
nucleus.
– Mass number is the number of protons and
neutrons located in the nucleus.
• The atomic mass on the periodic table reflects
the average mass of all known isotopes for a
given element.
4 Forces at Work in an Atom
• Gravitational Force-pulls objects toward one
another.
• Electromagnetic Force-opposite charges are
attracted to each other.
• Strong Force-holds the protons and neutrons
together.
• Weak Force-in unstable atoms, a neutron can
change into a proton and an electron.
Chapter 5 Section 1
Arranging the Elements
Who developed the Periodic Table?
When? Why?
• Dmitri Mendeleev-Russian scientist
• mid 1800s
• He wanted to show that there was a
relationship between the various known
elements.
What did Mendeleev do?
• He took the 63 known elements at the time
and worked with their physical and chemical
properties.
– Atomic mass
– density
– melting point
– bonding power
• valence number indicates the process by which an
element loses, gains or shares the valance electrons
What did Mendeleev’s Periodic Table
do for chemistry?
• He left spaces where he did not have a known
element with the correct properties.
• He predicted the empty spaces would be filled
with yet to be discovered elements.
– Even predicted their physical and chemical
properties.
– He was correct and 3 new elements were
discovered before he died.
Mendeleev’s Prediction vs. Actual
Properties of Element 32
Ekasilicon
Germanium
Date predicted: 1871
Atomic Mass: 72
Density: 5.5g/cm3
Bonding power: 4
Color: Dark gray
Date discovered:1886
Atomic mass: 72.6
Actual density: 5.32g/cm3
Bonding Power: 4
Color: Grayish white
How did the Periodic Table change?
• Mendeleev’s periodic table had some flaws.
• Henry Moseley a British scientist in the early
1900s determined the atomic number of the
elements.
– Atomic number tells how many protons are in the
nucleus of the atom.
– He rearranged the periodic table in order of
increasing atomic number and all elements fell
into place with no exceptions/ flaws.
How is the Modern Periodic Table
Organized?
• Based on increasing atomic number.
• Rows
– Each row is called a “period”
– the properties across a row change in a repeated
pattern
• Columns
– Each column is called a “group or family”
– the properties of down a column are similar but
not identical
Why understand the Periodic Table?
• Understanding the organization of the
Periodic table helps you know basic
information about each of the known
elements without having to memorize each
elements facts.
• It also, allows you to predict the physical and
chemical properties of an element just by its
location within the table.
Chapter 5 Section 2
Grouping the Elements
Looking at the Chemical Families
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Family 1: Alkali metals
Family 2: Alkaline-Earth metals
Family 3-12: Transition Metals
Family 13: Boron family
Family 14: Carbon family
Family 15: Nitrogen family
Family 16: Oxygen family
Family 17: Halogens
Family 18: Noble gases
Rare-Earth Elements: Lanthanide and Actinide series
Family 1: Alkali Metals
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1 valence electron to lose
soft, silver white and shiny
good conductors of heat and electricity
never found uncombined in nature-bond
readily
• reacts violently with water to produce
hydrogen gas which can burn/explode and
extreme heat.
Family 2: Alkaline-Earth Metals
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2 valence electrons to lose
never found uncombined in nature
not quite as reactive as alkali metals
Have many uses
– Ex: Magnesium can be mixed with other metals to
make low-density materials used in airplanes.
– Ex: Compounds of calcium are found in cement,
chalk, and even the human.
Family 3-12: Transition Metals
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1 or 2 valance electrons to lose or share
good conductors of heat and electricity
Less reactive than groups 1 & 2
bright colors
have properties similar to one another and
other metals, but different from other families
Rare-Earth Elements
• Appear in 2 rows at the bottom of the table to
keep from making the table too wide.
• Go with periods 6 & 7
– Lanthanides series is made up of soft, malleable
metals that have a high luster and conductivity.
– Actinides series is made up of all radioactive
elements and most must be made within the
laboratory.
Family 13: Boron family
• 3 valence electrons
• Boron is a metalloid, rest of the family are
metals
– metalloids show some properties of both metals
and non-metals
Family 14: Carbon Family
• 4 valence electrons to share
• carbon is a nonmetal, silicon & germanium are
metalloids and rest metals
• Carbon is often referred to as the “basis of
life” element
– carbon compounds make up living organisms
– it forms millions of compounds
• branch of chemistry known as Organic Chemistry
Family 15: Nitrogen Family
• 5 valence electrons which tend to be shared.
• Nitrogen and Phosphorus are nonmetals,
Arsenic a metalloid, and the rest metals.
• Nitrogen does not readily combine with other
elements even as it is the most abundant
element in the atmosphere.
Family 16: Oxygen Family
• 6 valence electrons to share
• Oxygen, Sulfur, Selenium and Tellurium are
nonmetals, Polonium is a rare radioactive
metal.
• Oxygen is most reactive element that readily
combine with almost every other element.
Family 17: Halogen Family
• 7 valence electrons which they will share or
want to gain electrons.
• Most active nonmetals that readily combine
with the alkali and alkaline metals to form
salts.
• Fluorine is the most active then Chlorine both
are never found uncombined in nature and
both are gases.
Family 18: Noble gases
• All have 8 valence electrons except for Helium
which has two.
• Inert gases: means they do not normally react
with other elements.
• They have a stable number of valence
electrons.
• Were some of the last elements to be
discovered and make up less than 1% of the
atmospheric gases.
Hydrogen
• Its properties do not match the properties of
any single group.
• It is above Group 1 b/c atoms of the alkali
metals also have only one electron in their
outer level.
• However, the physical properties are more like
a nonmetals.
• Hydrogen really is in a group of its own.