Chapter One
The big picture
– the Periodic Table
Handouts and Worksheets
‘The Chemical Investigator’ pages 5 – 11
Study On – worksheets 1 & 2
Matter
• Every material thin that you can see, smell and
touch, that occupies space and has mass, is a
form of matter.
• Matter is made up of very small particles and
may exist in solid, liquid or gaseous states.
• The behaviour of these particles is explained
by the particle model, or kinetic theory of
matter.
The Kinetic Theory of Matter
Key Points
• Matter is made up of tiny, invisible moving particles
• Particles of different substances have different sizes
• Lighter particles move faster than heavier ones at a
particular temperature
• As temperature rises, the particles move faster
• In a solid, the particles are very close and vibrate in fixed
positions
• Ina liquid, the particles are a little further apart. They have
more energy and they can move around each other
• In a gas, the particles are far apart. They move rapidly and
randomly in all the space that surrounds them.
Properties of Solids, Liquids and gases
• Solids have definite shapes and volumes
– Crystalline solids (salt, diamonds) have particles
arranged in regular, repetitive patterns
– Particles are able to vibrate but not move
– Amorphous solids do not have this regular
structure (eg rubber, putty)
Properties of Solids, Liquids and gases
• Liquids have a definite volume, taking the
shape of the container but their surfaces are
always horizontal
– Liquid particles move further apart than those in a
solid and are in constant motion, free to move
– Liquids can flow
• Gases take the same shape and volume as
their container, free to move in any direction.
Matter is usually defined as anything that has mass
and occupies space.
Total disorder
Lots of empty space
Gas
Disorder
Some space
Particles closer
together
Liquid
Order
Particles fixed
in position
Solid
• Solids, Liquids, and Gases
– Gases have no defined shape or defined volume
• Low density
– Liquids flow and can be poured from one container to
another
• Indefinite shape and takes on the shape of the container.
– Solids have a definite volume
• Have a definite shape.
Review
• Matter On the Move jeopardy revision
• Complete revision questions page 4 (1 –
4). Check and review your answers
Changes in States of Matter
• As temperature varies the particles change in
energy and distance apart
• Changing states of matter is about changing
densities, pressures, temperatures, and other
physical properties. The basic chemical
structure does not change.
• Summarise the terms:
– Melting point, Freezing point, Evaporate, Boiling
Point, Condensation Point, Volatile
Changes in States of Matter
Draw the image with the appropriate terms
Review
• Complete revision questions page 6 (5 –
7). Check and review your answers
Atomic theory
Sub atomic particles
Atomic theory
• The theory attempts to explain the
microscopic structure of materials.
• All Matter is made up of Atoms
• Summarise the timeline (page 7) with the
– date, scientist and major discovery.
Models of the Atom
a Historical Perspective
Early Greek Theories
• 400 B.C. - Democritus thought matter could
not be divided indefinitely.
Democritus
• This led to the idea of atoms in a void.
fire
earth
Aristotle
air
water
• 350 B.C - Aristotle modified an earlier
theory that matter was made of four
“elements”: earth, fire, water, air.
• Aristotle was wrong. However, his
theory persisted for 2000 years.
John Dalton
• 1800 -Dalton proposed a modern atomic
model
• Allexperimentation
matter is made ofnot
atoms.
based on
on pure
reason.• Atoms of an element are identical.
• Each element has different atoms.
• Atoms of different elements combine
in constant ratios to form compounds.
• Atoms are rearranged in reactions.
• His ideas account for the law of conservation of
mass (atoms are neither created nor destroyed)
and the law of constant composition (elements
combine in fixed ratios).
Dalton’s Postulates
1. Every element is composed of tiny particles called atoms
1. All atoms of a given element are identical
1. Atoms of different elements have different properties
2. Atoms of an element are NOT changed into atoms of
another element by chemical processes
1. Matter can neither be created nor destroyed
3. Compounds are formed when atoms of more than one
element combine
Dalton’s Laws
1. The Law of Constant Composition:
“Any given compound always consists of the same
atoms and the same ratio of atoms. For example, water
always consists of oxygen and hydrogen atoms, and
it is always 89 percent oxygen by mass and 11 percent
hydrogen by mass”
2. The Law of Conservation of Mass:
“The total mass of materials before and after a chemical
reaction must be the same. For example, if we combine
89 grams of oxygen with 11 grams of hydrogen under
the appropriate conditions, 100 grams of water will be
produced—no more and no less.”
Dalton’s Laws
3. The Law of Multiple Proportions:
“If two elements combine to form more than one compound,
the masses of one of the elements that can combine with a given
mass of the other element are related by factors of small whole
numbers”
For example, water has an oxygen-to-hydrogen mass ratio of 7.9:1.
Hydrogen peroxide, another compound consisting of oxygen and
hydrogen, has an oxygen-to-hydrogen mass ratio of 15.8:1.
The ratio of these two ratios gives a small whole number.
Adding Electrons to the Model
Materials, when rubbed, can develop a charge
difference. This electricity is called “cathode rays” when
passed through an evacuated tube (demos).
These rays have a small mass and are negative.
Thompson noted that these negative subatomic particles
were a fundamental part of all atoms.
1) Dalton’s “Billiard ball” model (1800-1900)
Atoms are solid and indivisible.
2) Thompson “Plum pudding” model (1900)
Negative electrons in a positive framework.
3) The Rutherford model (around 1910)
Atoms are mostly empty space.
Negative electrons orbit a positive nucleus.
Ernest Rutherford
• Rutherford shot alpha () particles at gold foil.
Zinc sulfide screen
Thin gold foil
Lead block
Radioactive
substance path of invisible
-particles
Most particles passed through.
So, atoms are mostly empty.
Some positive -particles
deflected or bounced back!
Thus, a “nucleus” is positive &
holds most of an atom’s mass.
• Rutherfords gold foil experiment
Bohr’s model
• Electrons orbit the nucleus in “shells”
• Electrons can be bumped up to a higher
shell if hit by an electron or a photon of light.
There are 2 types of spectra: continuous spectra & line
spectra. It’s when electrons fall back down that they
release a photon. These jumps down from “shell” to
“shell” account for the line spectra seen in gas discharge
tubes (through spectroscopes).
The Structure of Atoms
•
Summarise and/or define:
•
Nuclear model of the atom, Protons, Neutrons,
Electrons, Sub-atomic particles, Ions, Elements, Atom,
Atomic number, Mass number, Isotopic symbol,
Isotopes, Atomic emission spectrum.
Atomic numbers, Mass numbers
•
•
•
There are 3 types of subatomic particles. We already know
about electrons (e–) & protons (p+). Neutrons (n0) were also
shown to exist (1930s).
They have: no charge, a mass similar to protons
Elements are often symbolized with their mass number and
16
atomic number
E.g. Oxygen: 8 O
• These values are given on the periodic table.
• For now, round the mass # to a whole number.
• These numbers tell you a lot about atoms.
# of protons = # of electrons = atomic number
# of neutrons = mass number – atomic number
• Calculate # of e–, n0, p+ for Ca, Ar, and Br.
Atomic
Mass
p+
n0
e–
Ca
20
40
20
20
20
Ar
18
40
18
22
18
Br
35
80
35
45
35
Review
•
Complete the revision questions page 9 (8 – 11).
Check and review your answers
•
What is the symbol for the Atomic number?
Z
What is the symbol for the Mass number?
A
•
•
•
Isotopes and Radioisotopes
Atoms of the same element that have different numbers
of neutrons are called isotopes.
• Due to isotopes, mass #s are not round #s.
• Li (6.9) is made up of both 6Li and 7Li.relative atomic mass
• Often, at least one isotope is unstable.
• It breaks down, releasing radioactivity.
• These types of isotopes are called radioisotopes
Q- Sometimes an isotope is written without its atomic
number - e.g. 35S (or S-35).
Q- Draw B-R diagrams for the two Li isotopes.
A- The atomic # of an element doesn’t change Although
the number of neutrons can vary, atoms have definite
numbers of protons.
•
6Li
7Li
3 p+
3 n0
2e– 1e–
3 p+
4 n0
2e– 1e–
Isotopes
Isotopes of Lithium illustration
For more lessons, visit
www.chalkbored.com
Isotopes
•
•
•
•
All atoms of a particular element have the same
number of protons.
Atoms with the same number of protons but a
different number of neutrons are called isotopes
Isotopes have similar chemical properties because
their electron structure is the same. They have
different physical properties due to their different
masses.
List the three naturally occurring isotopes of oxygen.
Oxygen
•
•
•
Isotopes of Oxygen
There are three stable isotopes of oxygen that lead
to oxygen (O) having a standard atomic mass of
15.9994(3) amu
Naturally occurring oxygen is composed of three
stable isotopes, 16O, 17O, and 18O, with 16O being
the most abundant (99.762% natural abundance).
Known oxygen isotopes range in mass number
from 12 to 24
Review
•
•
Work through the sample problem on page 10
Complete the revision questions pages 10, 11 (12
– 17). Check and review your answers.
Atomic Emission Spectrum
•
•
•
Every element emits light if it is heated by passing
an electric discharge through its gas or vapour
This happens because the atoms of the element
absorb energy, then lose it and emit it as light.
Atomic emission spectrum consist of separate
lines of coloured light, each line of the spectrum
corresponding to one particular frequency of light
being given off by the atom: therefore each line
corresponds to one exact amount of energy being
emitted.
Bohr’s
model
• Electrons orbit the nucleus in “shells”
• Electrons can be bumped up to a higher
shell if hit by an electron or a photon of light.
There are 2 types of spectra: continuous spectra & line
spectra. It’s when electrons fall back down that they
release a photon. These jumps down from “shell” to
“shell” account for the line spectra seen in gas discharge
tubes (through spectroscopes).
Bohr - Rutherford diagrams
• Putting all this together, we get B-R diagrams
• To draw them you must know the # of protons,
neutrons, and electrons (2,8,8,2 filling order)
• Draw protons (p+), (n0) in circle (i.e. “nucleus”)
• Draw electrons around in shells
He
p+
2
2 n0
Li
Li shorthand
p+
3
4 n0
3 p+
4 n0
2e– 1e–
Draw Be, B, Al and shorthand diagrams for O, Na
Be
B
Al
4 p+
5 n°
O
5 p+
6 n°
13 p+
14 n°
Na
8 p+ 2e– 6e–
8 n°
11 p+ 2e– 8e– 1e–
12 n°
Bohr’s energy levels
• How does Bohr’s model explain the Atomic
Emission Spectrum.
• Define the terms: ground state, energy levels,
excited state, photon
• Why does Bohr’s model not explain atoms more
complex than Hydrogen?
• What is the Quantum mechanics model?
Electron Shells
• The regions of space surrounding the
nucleus.
• The electron shells are labelled K, L, M, N
and numbered 1, 2, 3, 4
• A definite energy level is associated with
each shell (K – closest the nucleus –
lowest energy). Therefore an electron
has to gain energy to move away from the
nucleus.
Electron Shells
• If an electron gains enough energy to
completely leave the atom, the particle
that is left is not longer neutral and
called a positive ion.
• Explain how K can become K+. What is
the difference in the protons and
electrons when K becomes a positive
ion?
Electron Configuration
• Electron configuration – arrangement of
electrons in shells
• The maximum number of electrons that
each shell can hold is 2n2 where n is the
number or energy level
Electron Configuration
• Electron shells are filled in order from the
nucleus (lowest energy level first – K)
• For the first 20 elements the outer shell
never has more than 8 electrons
• The outer shell electrons mainly
determine the chemical properties of an
element.
Electron Configuration
• Each electron has its own distinct energy, this
energy corresponds to the energy level it
occupies
• Electrons can gain or lose energy, but that
amount of energy gained or lost is a fixed
amount of energy.
• This fixed amount of energy gained allows an
electron to move to a higher energy level.
Electron Configuration
• Ground state – electrons occupy the
lowest available oribitals.
• Excited state – unstable condition,
electrons temporarily move to a higher
energy level.
• When electrons are subject to stimuli
such as heat, light, or electricity,
electrons may absorb energy and
temporarily move to a higher energy
level.
Electron Configuration
• Chemical properties are based on the
number of electrons in the outer energy
level.
• Valence electrons are these outer
electrons.
• Quantum theory – explains chemical
behaviour of atoms.
• Quantum numbers – electrons are
described as a set of four numbers
Electron Configuration
• First number describes the major energy
level of the electron and is called the principle
energy level.
• Principle energy levels have sublevels.
There are as many sublevels as the number
of that energy level.
• s is the first sublevel, p is the second, d is the
third and f is the fourth.
• ie) 3s means third energy level and first
sublevel
Electron Configuration
• Electron configuration – distribution of
electrons in an atom.
• Electron configuration for oxygen would
be 1s22s22p4
Review
• Atomic structure on line multiple choice
• Periodic Table quiz
• atomic structure quiz
• Complete the revision questions pages
14, 15 (18 – 22). Check and review
your answers.
Elements make compounds
Explosive experiment
• Goggles, bench
mat, tongs,
bunsen, lighter and
magnesium.
• CAREFULLY put
magnesium in
flame.
Making a compound
•What energy is being released?
-When magnesium is burned, energy is released as
light.
•Is there a new substance being formed?
-A new substance forms, it is white.
•Is this new substance lighter or heavier?
-This new substance is heavier as there are two
different elements in the substance, they are
magnesium and oxygen.
Elements make compounds
• When elements join to make a compound,
energy is released and a new substance is
formed.
• A word which means ‘energy released’ is
exothermic.
• What do you think endothermic means?
Elements and Compounds
• All matter is made up of combinations of
elements with atomic theory explaining the
behaviour of these elements and their
compounds
• An element is a substance that cannot be
broken down into simpler substances because it
is made up of only one type of atom.
• Elements can be grouped according to their
physical and chemical properties as either
metals or non-metals.
Elements, Compounds and Mixtures
• 2. Language of chemistry
• Define and give examples of:
– An element
– A compound
– A mixture
• Mixtures and Pure Substances
– A mixture has unlike parts and a composition
that varies from sample to sample
– A heterogeneous mixture has physically
distinct parts with different properties.
– A homogeneous mixture is the same
throughout the sample
– Pure substances are substances with a fixed
composition
A classification scheme for matter.
– A physical change is a change that does not alter
the identity of the matter. Physical changes are
about energy and states of matter
– A chemical change is a change that does alter the
identity of the matter. Chemical changes happen
on a molecular level. ie burning sugar.
– A compound is a pure substance that can be
decomposed by a chemical change into simpler
substances with a fixed mass ratio
– An element is a pure substance which cannot be
broken down into anything simpler by either
physical or chemical means.
• Sugar (A) is a compound that can be easily
decomposed to simpler substances by heating. (B)
One of the simpler substances is the black element
carbon, which cannot be further decomposed by
chemical or physical means.
EXAMPLE
Isopropyl alcohol is a
A. heterogeneous mixture
B. homogeneous mixture
C. pure substance
D. Compound
E. pure substance and compound
E
Naming Compounds
Elements and symbols that
you should know:
1) Hydrogen
H
8) Oxygen
O
2) Helium
He
9) Fluorine
F
3) Lithium
Li
10)Neon
Ne
4) Beryllium
Be
11)Sodium
Na
5) Boron
B
12)Magnesium
Mg
6) Carbon
C
13)Aluminium
Al
7) Nitrogen
N
14)Silicon
Si
Elements and symbols that
you should know:
15) Phosphorus P
16) Sulphur
S
17) Chlorine
Cl
18) Argon
Ar
19) Potassium
K
20) Calcium
Ca
If two identical elements combine then the
name doesn’t change.
.e.g. oxygen + oxygen  oxygen
This happens with the following elements:
1) H2
2) N2
3) O2
These elements always go
around in pairs. For
example, hydrogen looks like
this:
When two elements join the name ends with
____ide
e.g. Magnesium + oxygen
oxide
1) Sodium + chlorine
1) Magnesium + fluorine
magnesium
When three or more elements combine and one of them
is oxygen the ending is _____ite or ________ate
e.g. Copper + sulphur + oxygen
1) Calcium + carbon + oxygen
2) Potassium + carbon + oxygen
3) Sodium + sulphur + oxygen
Copper sulphate
Review
• Complete the revision questions
pages17,18 (23 – 27)
Elements
• Discovery of Modern Elements
– Antoine Lavoisier suggested that burning was
actually a chemical combination with oxygen.
– Lavoisier realized that there needed to be a new
concept of elements, compounds, and chemical
change.
– We now know that there are 89 naturally-occurring
elements and at least 23 short-lived and artificially
prepared.
• Priestley produced a gas (oxygen) by using sunlight to
heat mercuric oxide kept in a closed container. The
oxygen forced some of the mercury out of the jar as it
was produced, increasing the volume about five times.
• Lavoisier heated a measured amount of
mercury to form the red oxide of mercury. He
measured the amount of oxygen removed from
the jar and the amount of red oxide formed.
When the reaction was reversed, he found the
original amounts of mercury and oxygen.
Names of Elements
– The first 103 elements have internationally
accepted names, which are derived from:
• The compound or substance in which the element
was discovered
• An unusual or identifying property of the element
• Places, cities, and countries
• Famous scientists
• Greek mythology
• Astronomical objects.
• Here are some of the symbols Dalton used for atoms
of elements and molecules of compounds. He
probably used a circle for each because, like the
ancient Greeks, he thought of atoms as tiny, round
hard spheres.
• The elements of aluminum, Iron, Oxygen, and Silicon
make up about 88 percent of the earth's solid surface.
Water on the surface and in the air as clouds and fog is
made up of hydrogen and oxygen. The air is 99 percent
nitrogen and oxygen. Hydrogen, oxygen, and carbon
make up 97 percent of a person. Thus almost
everything you see in this picture us made up of just six
elements.
Atomic Mass
Charge
(if ion)
Symbol
Atomic Number
Hydrogen
1
1
H
Protons: 1
Neutrons: 0
Electrons: 1
Sodium
23
11
Protons: 11
Neutrons: 12
Electrons: 11
Na
Rhenium
186
75
Protons: 75
Neutrons: 111
Electrons: 75
Re
Rhenium isotope
187
75
Protons: 75
Neutrons: 112
Electrons: 75
Re
EXAMPLE
How many protons, neutrons and electrons are
found in an
atom of
133
55
Cs
Atomic number = protons and electrons
There are 55 protons and 55 electrons
Mass number = sum of protons and neutrons
133 – 55 = 78
There are 78 neutrons
The Periodic Law
• Dmitri Medeleev gave us a functional scheme
with which to classify elements.
– Mendeleev’s scheme was based on chemical
properties of the elements.
– It was noticed that the chemical properties of
elements increased in a periodic manner.
– The periodicity of the elements was demonstrated
by Medeleev when he used the table to predict to
occurrence and chemical properties of elements
which had not yet been discovered.
• Mendeleev left blank
spaces in his table when
the properties of the
elements above and below
did not seem to match. The
existence of unknown
elements was predicted by
Mendeleev on the basis of
the blank spaces. When
the unknown elements
were discovered, it was
found that Mendeleev had
closely predicted the
properties of the elements
as well as their discovery.
• The Periodic Law
– Similar physical and chemical properties
recur periodically when the elements are
listed in order of increasing atomic number.
The Modern Periodic Table
• Introduction
– The periodic table is made up of rows of elements
and columns.
– An element is identified by its chemical symbol.
– The number above the symbol is the atomic
number
– The number below the symbol is the rounded
atomic weight of the element.
– A row is called a period
– A column is called a group
Arrangement of the Periodic Table
• Elements with the same group number
have the same number of electrons in the
outer shell – valence electrons
• Q – which groups do the following
elements below to? K, B, He, Cl, Ca (ie
how many electrons in their outer shell?)
Arrangement of the Periodic Table
• The period number refers to the number of
the outermost shell containing electrons.
• Q – which periods do the following
elements belong to? Ca, Mg, Si, N, O, He
Arrangement of the Periodic Table Groups
• Valence electrons mainly determine
chemical reactivity. Elements in the same
group usually have similar chemical
properties. There is a progressive change
in their physical properties.
Arrangement of the Periodic Table Groups
• Group 1 – Alkali Metals
–
–
–
–
–
–
React with water to form alkaline solutions
Low melting points
Low boiling points
Densities so low they can float on water
So soft they can be cut with a knife
Very reactive and must be stored under oil to prevent
them reacting with oxygen and water vapour in the air
– Eg – Na and K
Arrangement of the Periodic Table Groups
• Group 2 – Alkaline Metals
– They were first extracted from oxides found in
the earth’s crust
– Less reactive than alkali metals
– Eg Be, Mg
Arrangement of the Periodic Table –
Groups
• Group 17 – Halogens
– So reactive that they never occur freely in
nature
– They occur combined with different metals to
form salts
– Eg Cl, I
Arrangement of the Periodic Table Groups
• Group 18 – Noble Gases
– They do not react readily with other
substances
– The Noble Gases
Arrangement of the Periodic Table Groups
• Groups 3 – 12 – Transition Metals
– Hard metals
– Usually have high melting point
– Usually have high boiling point
– Usually form coloured compounds
Arrangement of the Periodic Table Periods
• Period number represents the number of
occupied electron shells in the atoms of
the elements in that group.
• Elements in the same period share a
gradual change in their physical and
chemical properties.
Arrangement of the Periodic Table Periods
• Period 6 - Lanthanides or Rare Earth
– Rare occurrence in nature
• Period 7 – Actinides
– Radioactive elements
Arrangement of the Periodic Table
• Metal and Non-Metals
– Diagonal step line starting at B
– Separate metals from non-metals
– Change from metallic to non-metallic is
gradual and some elements have
characteristics of both metals and non-metals.
Arrangement of the Periodic
Table
• Metalloids
– Elements with combinations of metallic and
non-metallic properties
– Eg Si and Ge have high melting points and
high boiling points (like metals), but have low
densities and are brittle (like non-metals)
Trends in the Periodic Table
• Across a period
1. Metallic character
•
•
Decreases across the table, while non-metallic character
increases
Elements in groups 1 and 2 are metals, group 18 are gases
2. Atomic size
•
Generally decreases from metals to non-metals across a
period – electrons are being added to the same outer shell
while number of protons in the nucleus in increasing
(increases electrostatic attraction between electrons and
nucleus, pulls outer electrons closer to the nucleus,
reducing atomic radii
Trends in the Periodic Table
• Across a period
1. Reactivity
•
•
•
Generally lowest in the middle of a period and increases at
either end (not including noble gases)
Eg period 3 reactivity with Hydrochloric acid
High Na Mg  Al Low, Si non reactive, Low P  S  Cl
High
2. Electronegativity (electron attracting power of an
atom)
•
Increases from metals to non-metals across a period due
to the electrons being increasingly attracted to the nucleus.
Trends in the Periodic Table
• Down a Group
1. Metallic character
•
Increase down a group, while non-metallic
character decreases
2. Atomic size
•
Generally increases down a group as electrons
are added to successive main shells and are
therefore further away from the nucleus
Trends in the Periodic Table
• Down a Group
1. Reactivity
•
•
•
•
Metals - Generally increases down a group
Eg reactivity with water – High K, Na, Li Low
Non-metals – Generally decreases down a group
Eg. Reactivity with water – High F, Cl, Br, I Low
2. Electronegativity
•
Generally decreases down a group – electrons are further
away from the nucleus and are ‘shielded’ from the
attraction of the nucleus by the inner shell electrons
Trends in the Periodic Table
• The Position of Hydrogen
– Sometimes in group 1, sometimes in group 17
– Group 1 loses an electron
– Group 17 – many of its properties are similar
to the Halogens – gains an electron when it
reacts with some elements
Review
• Complete the revision questions pages
21,22 (28 – 34)
• (A) Periods of the periodic table, and (B) groups of the periodic table.
• Periodic Patterns
– The chemical behavior of elements is
determined by its electron configuration
– Energy levels are quantized so roughly
correspond to layers of electrons around the
nucleus.
– A shell is all the electrons with the same value
of n.
• n is a row in the periodic table.
– Each period begins with a new outer electron
shell
– Each period ends with a completely filled outer
shell that has the maximum number of electrons
for that shell.
– The number identifying the A families identifies
the number of electrons in the outer shell, except
helium
– The outer shell electrons are responsible for
chemical reactions.
– Group A elements are called representative
elements
– Group B elements are called transition
elements.
• Chemical “Families”
– IA are called alkali metals because the react with
water to from an alkaline solution
– Group IIA are called the alkali earth metals
because they are reactive, but not as reactive as
Group IA.
• They are also soft metals like Earth.
– Group VIIA are the halogens
• These need only one electron to fill their outer shell
• They are very reactive.
– Group VIIIA are the noble gases as they have
completely filled outer shells
• They are almost non reactive.
Metal: Elements that are usually solids at room
temperature. Most elements are metals.
Non-Metal: Elements in the upper right corner of
the periodic
Table. Their chemical and physical properties are
different from metals.
Metalloid: Elements that lie on a diagonal line
between the Metals and non-metals. Their
chemical and physical properties are intermediate
between the two.
– When an atom or molecule gain or loses an
electron it becomes an ion.
• A cation has lost an electron and therefore has a
positive charge
• An anion has gained an electron and therefore
has a negative charge.
– Elements with 1, 2, or 3 electrons in their outer
shell tend to lose electrons to fill their outer shell
and become cations.
• These are the metals which always tend to lose
electrons.
– Elements with 5 to 7 electrons in their outer shell
tend to gain electrons to fill their outer shell and
become anions.
• These are the nonmetals which always tend to gain
electrons.
– Semiconductors (metalloids) occur at the
dividing line between metals and nonmetals.
EXAMPLE
What would the charge be on a sodium ion?
Since sodium in in Group IA it is a metal and so
would LOSE an electron
You can tell how many would be lost by the group
number
Group 1A elements lose 1 electron
So the charge would be +1
Remember an electron is negatively charged.
When you lose them atom becomes positively
charged when you gain them it becomes negatively
charged
EXAMPLE
How would you right the symbol for the sodium
CATION?
Na
+1
How many outer electrons does sodium have
before it loses one?
It has 1…remember the group number!
Chapter One Review
• Solids, liquids, gases on line multiple choice
• Periodic Table
• Complete the multiple choice questions page
24
• Consider each of the review questions 1 – 29
• Periodic table on line questions