Transcript Unit 1 Matter review

1.2 Investigating Matter
• Matter is anything that has mass and volume.
• Mass is the amount of matter in a substance or object.
• Mass is often measured in grams or kilograms.
• Volume is the amount of space a substance or an object
occupies.
• Volume is often measured in litres.
• More information and practice on pages 480 - 483
See page 16
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1.3 Atomic Theory
• Early ideas about matter
• Greek philosophers believed that matter was made of atomos
that were the smallest pieces of matter.
• Aristotle believed matter was made of
different combinations of earth, air,
fire, and water.
• Alchemists experimented with matter and
tried to turn common metals into gold.
Their activities marked the beginning of
our understanding of matter.
See pages 28 - 29
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Development of Atomic Theory I
• John Dalton (1766 - 1844)
• Credited with developing a theory that was a new way of
explaining matter.
• He studied gases that make up Earth’s atmosphere.
Based on his studies, he suggested that:
• matter is made of small, hard spheres
that are different for different
elements
• the smallest particle of an element
is called an atom
• This is the basis for Dalton’s Atomic Theory.
See page 29
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Dalton’s Atomic Theory
1. All matter is made of small particles called atoms.
2. Atoms cannot be created, destroyed, or divided into
smaller particles.
3. All atoms of the same element are identical in mass
and size, but they are different in mass and size from
the atoms of other elements.
4. Compounds are created when atoms of different
elements link together in definite proportions.
See page 30
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Atomic Theory II
• J. J. Thomson (1856 - 1940)
• Thomson studied electric currents in gas discharge tubes (like
today’s fluorescent lights). From his studies, he determined that the
currents were streams of negatively charged particles. These were
later called electrons.
• He hypothesized that atoms are made
of smaller particles. He proposed
the “raisin bun” model of the atom.
• This model is best visualized as a
positively charged bun with negatively
charge particles spread out in it like raisins.
See page 30
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Atomic Theory III
• Ernest Rutherford (1871 - 1937)
• After experimenting with charged particles, he
found that some particles were deflected in
directions not originally predicted.
• He suggested that the deflection of the charged
particles was because the atom contained a tiny
dense centre called a
nucleus, and electrons
moved around the nucleus.
See page 31
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Atomic Theory IV
• Niels Bohr (1885 - 1962)
• He studied gaseous samples of atoms, which were
made to glow by passing an electric current
through them.
• Based on his observations, Bohr proposed
that electrons surround
the nucleus in specific
“energy levels” or “shells.”
See page 31 - 32
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Inside the Atom
• An atom is the smallest particle of an element that
retains the properties of the element.
• All atoms are made up of three kinds of particles called subatomic
particles. These particles are:
• electrons
• protons
• neutrons
Take the Section 1.3 Quiz
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See pages 32 - 33
Chemical Change
• A chemical change is a change in matter
that occurs when substances combine to form
new substances.
• For example, fireworks
• Find Out Activity 1-2A Bag of Change
See page 17
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Physical Change and Changes
of State
• When a physical change occurs, there may be a
change in appearance, but no new substances
are formed.
• For example, when ice or snow
melts to water, this physical change
is a change of state. No new
substances are formed.
See page 18
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The Particle Model of Matter
• Describes the behaviour of matter
• Matter is made of small particles.
• There are spaces between the particles.
• Gases have more space than liquids. Liquids have
more space than solids.
• Particles are always moving.
• Particles are attracted to each other. The
strength of attraction depends on the type of
particle.
See page 18
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The Kinetic Molecular Theory
• Describes what happens to matter when
the kinetic energy of particles changes.
The main points in the theory are:
• Matter is made of small particles.
• There is empty space between particles.
• Particles are constantly moving.
• Solid particles are packed together and
cannot move freely. They can only vibrate.
• Liquid particles are farther apart and can slide
past each other.
• Gas particles are far apart and move around
quickly.
• Energy makes particles move.
See page 19
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The Kinetic Molecular Theory
and Changes of State
Solid
Particles are close together, fixed in position and vibrating.
Melting
As temperature increases, particles’ kinetic
energy increases.
Liquid
Particles are still close, but slide past one another.
Boiling
As temperature increases, particles’ kinetic energy
continues to increase, creating more space.
Gas
Particles are highly energetic and moving freely.
See page 20
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Temperature and Changes of
State
See page 21
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Describing Matter
• Physical Properties
• Qualitative - state, colour, malleability
• Quantitative - conductivity, viscosity, density
• Pure Substances
• Element - a pure substance that cannot be broken
down or separated into simpler substances (e.g.,
gold)
• Compound - a pure substance composed of at
least two elements (e.g., water)
Take the Section 1.2 Quiz
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See pages 22 - 23
2.1 Elements
• Why are elements studied in chemistry?
• Chemistry is the study of matter and its changes.
• Elements make up an incredible variety of different substances.
• An element is a pure substance that cannot be broken down or
separated into simpler substances. Each element is one kind of
atom.
• By studying elements, we can learn more about the structure of
matter.
See page 42
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Chemical Symbols
• Element names and symbols
• Because elements have different names in different
languages, chemists use international symbols for
them
• Chemical symbols consist of one or two letters.
• Ancient names are used as the source of many of the
symbols. Example:
• Mercury - Hg - Hydragyrum (Latin for liquid silver)
See pages 43 - 44
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Chemical Symbols
All elements are represented by symbols.
Here are just a few element symbol examples:
See page 44
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Common Elements
• Hydrogen
• Colourless, odourless, tasteless, and highly flammable gas.
• Makes up over 90 percent of the
atoms in the universe
• Used in producing fertilizers
• Lighter than air
• Can be separated from water or
gasoline and be used as a
source of fuel
See page 45
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Common Elements
• Iron (Fe) - mixed with
carbon to make steel
• Good structural material,
but can rust when mixed
with water or oxygen
• Oxygen (O) - gaseous
element we breathe
• 21 % of the atmosphere
• Reacts with most other elements
Oxygen and iron react in burning thermite
GNU license photo
Iron in a river turns water and rocks red
See page 45
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Other Common Elements
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Sodium (Na) - soft metal that reacts with water
Chlorine (Cl) - yellow-green gas that is highly toxic
Mercury (Hg) - liquid at room temperature metal.
Silver (Ag) - precious metal mined in British Columbia
Silicon (Si) - brittle, grey, semiconductor that is second most
common element in Earth’s crust.
Na
Cl
Hg
Ag
Take the Section 2.1 Quiz
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Si
See pages 46 - 47
2.2 Periodic Table
• Origin of the periodic table
• Chemists in the 10th century wished to organize elements
• Attempts focused on grouping elements with similar properties
• In 1867, Dimitri Mendeleev found patterns in the elements and
organized them into table
• The resulting table had holes for elements
not yet discovered
See page 52
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Periodic Table
• The Periodic Table provides information
on the physical and chemical properties of elements
Atomic Mass - mass of average atom
Atomic Number - number of protons
Ion Charge - electric charge that forms
when an atom gains or loses electrons
See page 53
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Periodic Table
See page 54
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Metals, Non-metals, Metalloids
• Period table has interesting patterns
• Due to Mendeleev’s organization, interesting patterns are created,
such as the groups: metals, non-metals and metalloids.
See page 55
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Periods and Families
• Each horizontal row in the periodic table is a period
• Vertical columns form groups or chemical families
• Alkali metals - highly reactive group 1
• Alkaline earth metals - group 2, burn
in air if heated
• Halogens - group 17, highly reactive
non-metals
• Noble gases - group 18, stable and
unreactive non-metals
Take the Section 2.2 Quiz
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See pages 56 - 57
2.3 Periodic Table and Atomic
Theory
• Elements with similar properties have similar
electron arrangements
• Bohr models show electron arrangement in shells
See page 64 - 65
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Bohr model patterns
• Chemical families on the periodic table have the same
number of valence electrons
• Elements in the same period have the same number of shells
• Period number indicates the number of electron shells
See page 66
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Atom Stability
• Noble gases are very unreactive because their atoms
have filled valence shells. Filled shells make atoms stable.
Atoms with filled shells do not easily trade or share electrons.
• Other atoms gain or lose electrons in order to achieve the stability
of noble gases. Gaining or losing electrons makes atoms into ions.
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Metals lose electrons to form positive ions
Non-metals gain electrons to form negative ions
Ions have a similar electron arrangement to the nearest noble gas
Example: Sodium ion (Na+) has 11 protons (11+) and
10 electrons (10-) for a total charge of 1+
Take the Section 2.3 Quiz
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See pages 66 - 67
3.1 Compounds
• Compounds are pure substances made of more than
one kind of atom joined together. The atoms are held together with
chemical bonds.
• Compounds come in two basic types: covalent and ionic.
Covalent compounds share
electrons to form molecules.
Example: water
In ionic compounds, atoms
gain or lose electrons to form
ions. Example: NaCl
See pages 76 - 78
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Ionic Compounds
• Ionic solids exist as a solid in the form of an ionic lattice.
• The positive ions attract all of the negative ions, and vice versa. In
the example of table salt (NaCl) the one-to-one ratio of ions results
in a simple square-shaped ionic cyrstal:
See page 78
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Polyatomic Ions
• Covalent and ionic bonds can occur together
• A molecule can gain or lose electrons to become charged,
forming a polyatomic ion.
• Polyatomic ions form compounds like
other ions.
• Example: Ammonium ion (NH4+)
• There are many types of
polyatomic ions, but they occur
in a few basic shapes.
Take the Section 3.1 Quiz
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See pages 79 - 80
3.2 Names and Formulas of
Ionic Compounds
•
The chemical name indicates the elements present in
the compound. Chemical names for ionic compounds are given
according to rules.
•
The positive ion is always the first part of the name
•
The negative ion is always the second part of the name
•
The non-metal ion’s name ends with the suffix “-ide”
See page 85
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Ionic Chemical Formulas
• In an ionic compound, the positive charges balance
the negative charges. This balance of charge is used to determine
the smallest whole number ratio of positive to negative ions.
See page 87
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Multivalent Metal Compounds
• Many metals are multivalent, meaning the metals
form two or more different positive ions with different charges
• For example, the atom iron forms two ions Fe2+ and Fe3+
• Too distinguish different ions for the same metal, roman
numerals are added to their name. For example, Fe3+ would
be named “iron(III)”
See page 88
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Writing Multivalent Formulas
• Writing ionic compound formulas with multivalent
ions follows the same rules as regular ionic compounds
See page 89
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Multivalent Compound Names
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Steps to writing multivalent compound names are
as follows:
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Identify the metal and verify it forms more than one ion
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Determine the ratio of ions - for example, Fe2O3 means 2 iron
ions for every 3 oxygen ions
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Note the charge on the negative ion: Oxygen is O2•
The positive and negative charges must balance, so 2 iron
ions of 3+ charge (Fe3+) are needed to balance the 3 oxygen
ions
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Write the name of the compound: Iron(III) oxide
See page 90
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Polyatomic Ion Compounds
•
Steps to writing names for formulas involving
polyatomic ions are similar to other ionic compounds
Take the Section 3.2 Quiz
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See page 91
3.3 Physical and Chemical
Changes
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In physical changes, the appearance of a substance
changes, but the chemical bonds holding the substance together
do not change. Examples: melting, freezing, boiling
In chemical changes, new substances are produced in the process
of breaking chemical bonds and forming new ones.
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Evidence of chemical change:
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Colour change
Heat, light, sound produced or consumed
Bubbles of gas form
Formation of a precipitate
The change is difficult to reverse
See page 98
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Energy Changes
• In both physical and chemical changes, energy changes
take place. This energy change can mean releasing to or absorbing
energy from the environment.
• Exothermic reactions involve the overall release of energy in the
form of heat and light.
• Endothermic reactions
involve the overall absorption
of energy.
Instant Cold Pack: Endothermic
Campfire: exothermic
See page 99
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Applications of Chemical
Changes
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Some chemical changes present problems, while
others provide opportunities and advantages
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Corrosion is major problem for steel structures - by protecting
steel surfaces, the chemical reaction of iron with oxygen can
be prevented.
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First nations people of the Pacific Coast have used smoking
as a means of preserving food. Smoke causes chemical
changes in meat that kill bacteria.
Take the Section 3.3 Quiz
(c) McGraw Hill Ryerson 2007
See page 100