4.1 ATOMIC THEORY & BONDING

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Transcript 4.1 ATOMIC THEORY & BONDING

1. Demonstrate knowledge of the three subatomic particles, their
properties, and their location within the atom.
2. Define and give examples of ionic bonding (e.g., metal and
non‐metal) and covalent bonding (e.g., two non‐metals, diatomic
elements).
3. With reference to elements 1 to 20 on the periodic table, draw
and interpret Bohr models, including protons, neutrons, and
electrons, of:
•
atoms (neutral)
•
ions (charged)
•
molecules ‐ covalent bonding (e.g., O2, CH4)
•
ionic compounds (e.g., CaCl2)
4. Identify valence electrons using the periodic table.
5. Distinguish between paired and unpaired electrons for a
single atom.
6. Draw and interpret Lewis diagrams showing single bonds
for simple ionic compounds and covalent molecules (e.g.,
NaCl, MgO, BaBr2, H2O, CH4, NH3).
7. Distinguish between lone pairs and bonding pairs of
electrons in molecules.
Alkali earth metals
Alkali metals
Anions
Atomic #
Atomic number
Atomic Theory
Atoms
Bohr diagram
Cations
Chemical Change
Chemical reaction
Compound
Covalent bonding
Covalent Compound
Electrons
Element
Family/Group
Halogens
Ionic bonding
Ionic compounds
Ions
Lewis Diagram
Matter
Metal
Metalloids
Mixture
Molecule
Neutron
Noble gases
Non-Metal
Nucleus
Period
Proton
Pure Substance
Stable outer shell
Subatomic particle
Transition metals
Valence electrons
Solutions
Mechanical
Suspensions
Elements
Compounds
• An atom is the smallest particle of an
element that still has the properties of that
element
50 million atoms, lined
up end to end = 1 cm
An atom = proton(s) +
neutron(s) + electron(s)
See pages 168 - 169
(c) McGraw Hill Ryerson 2007
• Atoms join together to form compounds.
– A compound is a pure substance that is composed of
two or more atoms combined in a specific way.
– Oxygen and hydrogen are atoms/elements; H2O is a
compound.
See pages 168 - 169
(c) McGraw Hill Ryerson 2007
A chemical
change occurs
when the
arrangement
of atoms in
compounds
changes to
form new
compounds.
See pages 168 - 169
(c) McGraw Hill Ryerson 2007
• Atoms are made up of smaller particles
called subatomic particles.
See page 170
(c) McGraw Hill Ryerson 2007
If the proton &
neutron were
enlarged, and each
had the mass of a
hippopotamus, the
electron, enlarged
to the same scale,
would have less
mass than an owl.
• The nucleus is at the
centre of an atom.
–The nucleus is composed of
-positive protons
-neutral neutrons
–Electrons exist in the space
surrounding the nucleus.
See page 170
(c) McGraw Hill Ryerson 2007
– # of protons = # of electrons in every atom
– Nuclear charge = charge on the nucleus = # of
protons
– Nuclear charge = Atomic number
– Atomic number = # of protons = # of electrons
See page 170
(c) McGraw Hill Ryerson 2007
INCREASING REACTIVITY
The Periodic Table
Where are the
following?
INCREASING REACTIVITY
• Atomic
number
See page 172
(c) McGraw Hill Ryerson 2007
• In the periodic table elements are listed in
order by their atomic number.
– Metals are on the left
– The transition metals range from group 3 -12
– Non-metals are on the right
– Metalloids form a “staircase” toward the right
side.
See page 171
(c) McGraw Hill Ryerson 2007
Metals (left of zig zag line)
Physical Properties of Metals: Shiny, good conductors of
heat and electricity, ductile (make wires) and malleable (thin
sheets). Easily lose electrons. Like to join with non-metals.
Corrode (tarnish/rust).
Nonmetals (right of zig zag line)
Physical Properties of Nonmetals: dull appearance, poor
conductor, brittle (breaks easily), not ductile or malleable.
Easily gain electrons. Like to join with metals, but will bond
to other non-metals.
Metalloids (on both sides of zigzag line)
Physical Properties of Metalloids: have properties of both
metals and nonmetals. Solid, shiny or dull, ductile and
malleable, conduct heat and electricity, but not very well.
The Periodic Table
Where are the
following?
INCREASING REACTIVITY
• Metals
• Non-metals
• Transition
metals
• Metalloids
See page 172
(c) McGraw Hill Ryerson 2007
– Rows of elements (across) are called periods.
• All elements in a period have their electrons in the
same general area around their nucleus.
• Example: period 3 all have 3 electron shells
sodium
magnesium
aluminum
See page 171
(c) McGraw Hill Ryerson 2007
– Columns of elements are called groups, or families.
• All elements in a family have similar properties and bond
with other elements in similar ways.
• Group 1 = alkali metals
• Group 2 = alkaline earth metals
• Group 17 = the halogens
• Group 18 = noble gases
18
1 2
17
See page 171
(c) McGraw Hill Ryerson 2007
Group 1 = alkali metals
very reactive metals
want to give away 1 electron
ie: lithium, sodium, potassium...
1
2
18
17
See page 171
(c) McGraw Hill Ryerson 2007
Group 2 = alkaline earth metals
somewhat reactive metals
want to give away 2 electrons
ie: beryllium, magnesium, calcium...
1
2
18
17
See page 171
(c) McGraw Hill Ryerson 2007
Group 17 = halogens
very reactive non-metals
want to accept 1 electron
react with alkali metals
ie: fluorine, chlorine, bromine......
1 2
18
17
See page 171
(c) McGraw Hill Ryerson 2007
Group 18 = noble gases
STABLE. Very non reactive gaseous non-metals
ie: helium, neon, argon......
18
1 2
17
See page 171
(c) McGraw Hill Ryerson 2007
The Periodic Table
Where are the
following?
• Period
INCREASING REACTIVITY
• Group/Family
•Alkali metals
• Alkaline earth
metals
• Halogens
• Noble gases
See page 172
(c) McGraw Hill Ryerson 2007
• Atoms gain and lose electrons to form bonds.
– The atoms become electrically charged
particles called ions.
See page 173
(c) McGraw Hill Ryerson 2007
• Atoms gain and lose electrons to form bonds.
– Metals lose negative electrons & become
positive ions.
– Positive ions are called CATIONS.
See page 173
(c) McGraw Hill Ryerson 2007
Some metals are MULTIVALENT and can
lose a varying number of electrons.
For example, iron, Fe, loses either two (Fe2+)
or three (Fe3+) electrons
See page 173
(c) McGraw Hill Ryerson 2007
• Atoms gain and lose electrons to form bonds.
– Non-metals gain electrons and become
negative ions
– Negative ions are called ANIONS
See page 173
(c) McGraw Hill Ryerson 2007
Atoms gain and lose electrons in an attempt to be STABLE.
The noble gases are stable because they have FULL outer
shells of electrons. They don’t need to lose or gain any e-s.
Atoms in each period want to have the same number of
electrons in their outer shell (VALENCE ELECTRONS) as
the noble gases on the end of their period.
See page 173
(c) McGraw Hill Ryerson 2007
• Bohr diagrams show how many electrons appear in each
electron shell around an atom.
– The first electron shell holds 2 electrons
– The second electron shell holds 8 electrons
– The third electron shell holds 8 electrons
– The fourth electron shell holds 18 electrons
 The noble gas
elements have
full electron
shells and are
very stable.
See page 174
(c) McGraw Hill Ryerson 2007
• Electrons appear in shells in a very predictable
manner.
– The period number = the number of shells in the
atom.
– Except for the transition elements (family 3-12), the last
digit of the group number = the number of electrons
in the valence shell.
See page 175
(c) McGraw Hill Ryerson 2007
• It has 2 + 8 + 8 = 18
electrons, and
therefore, 18 protons.
What element is this?
• It has three electron
shells, so it is in
period 3.
• It has eight electrons
in the outer (valence)
shell.
18 p
22 n
argon
See page 174
(c) McGraw Hill Ryerson 2007
• When two atoms get close together, their
valence electrons interact.
– If the valence electrons can combine to form a
low-energy bond, a compound is formed.
– Each atom in the compound attempts to have a
‘full’ outer shell of valence electrons.
See pages 176 - 177
(c) McGraw Hill Ryerson 2007
There are 2 types of compounds:
• IONIC COMPOUND: metals lose electrons and nonmetals gain electrons.
• Ionic bonds form when electrons are transferred from
positive (+) ions to negative (-) ions.
• The negative and positive ions are ATTRACTED to each
other and form a BOND.
See pages 176 - 177
(c) McGraw Hill Ryerson 2007
• Example ionic bond:
• lithium and oxygen form an ionic bond in the
compound Li2O.
+
lithium
oxygen
Electrons are transferred from the positive
ions to negative ions
Li+
O2Li+
lithium oxide, Li2O
See pages 176 - 177
(c) McGraw Hill Ryerson 2007
There are 2 types of compounds:
• COVALENT COMPOUND: atoms share electrons.
• Covalent bonds form when electrons are shared
between two non-metals.
• Electrons stay with their atom but overlap with other
shells.
See pages 176 - 177
(c) McGraw Hill Ryerson 2007
• Example covalent bond
• Hydrogen and fluorine form a covalent bond in
the compound hydrogen fluoride.
+
hydrogen
Hydrogen fluoride
fluorine
electrons are shared
(c) McGraw Hill Ryerson 2007
See pages 176 - 177
• Lewis diagrams illustrate chemical bonding by
showing only an atom’s valence electrons and
the chemical symbol.
Dots representing electrons
are placed around the element
symbols at the points of the
compass (north, east, south,
and west).
See page 178
(c) McGraw Hill Ryerson 2007
–Electron dots are placed singly until the
fourth electron is reached then they are
paired.
See page 178
(c) McGraw Hill Ryerson 2007
To write IONS using lewis diagrams follow these steps:
Step 1: Write the lewis diagram as you normally would.
Step 2: If the element has a POSITIVE combining capacity it will give away an
electron and become a POSITIVE ION (cation). Rewrite the lewis diagram to show
the element symbol in square brackets (no electrons needed as it has given them
away and they now have an EMPTY outer electron shell!) then add the + charge on
the outside of the brackets.
Step 3: If the element has a NEGATIVE ION (anion). Rewrite the lewis diagram to
show the element symbol in square brackets with extra electrons. They will now
have a FULL OUTER electron shell. Then add the - charge on the outside of the
brackets.
• Lewis diagrams and IONIC BONDS:
– For positive ions, one electron dot is removed from the
valence shell for each positive charge.
– For negative ions, one electron dot is added to each
valence shell for each negative charge.
– Square brackets are placed around each ion to indicate
transfer of electrons.
••
•
•
Be
••
••
•
•
•
•
Cl
••
•
•
••
Each beryllium has two
electrons to transfer away,
and each chlorine can
receive one more electron.
(c) McGraw Hill Ryerson 2007
•
•
Cl
••
••
• •
• •
Be
••
••
• •
• •
Cl
••
•
•
••
Since Be2+ can donate two
electrons and each Cl– can
accept only one, two Cl– ions
are necessary.
•
•
–
••
2+
••
–
• • Be • • Cl •
• •
••
•
••
••
••
Cl
beryllium chloride
See page 179
• Lewis diagrams and COVALENT BONDS:
– Like Bohr diagrams, valence electrons are drawn
to show sharing of electrons.
– The shared pairs of electrons are usually drawn
as a straight line.
See page 179
(c) McGraw Hill Ryerson 2007
• DIATOMIC MOLECULES, like O2 and H2, are also easy to draw as
Lewis diagrams.
The elements Hydrogen, Nitrogen, Fluorine, Oxygen, Iodine, Chlorine,
and Bromine are always found as diatomic molecules.
MEMORY TRICK: I Have No Bright Or Clever Friends
See page 180
(c) McGraw Hill Ryerson 2007