Download 4.1 PPT- Atomic Theory and 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
No new substances
produced
Only a change in state or
appearance
New substances produced
Very hard to reverse
= CHEMICAL REACTION
• 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
• 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
A chemical change occurs
when the arrangement of
atoms in compounds
changes to form new
compounds.
See pages 168 - 169
• Atoms are made up of smaller particles called subatomic
particles.
See page 170
• 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
•
•
•
•
# 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
INCREASING REACTIVITY
THE PERIODIC TABLE
Where are the
following?
INCREASING REACTIVITY
• Atomic
number
See page 172
• 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
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
INCREASING REACTIVITY
Where are the
following?
• Metals
• Non-metals
• Transition
metals
• Metalloids
See page 172
• 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
– 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
Group 1 = alkali metals
very reactive metals
want to give away 1 electron
ie: lithium, sodium, potassium...
1
2
18
17
See page 171
Group 2 = alkali earth metals
somewhat reactive metals
want to give away 2 electrons
ie: beryllium, magnesium, calcium...
1
2
18
17
See page 171
Group 17 = halogens
very reactive non-metals
want to accept 1 electron
react with alkali metals
ie: fluorine, chlorine, bromine......
18
1 2
17
See page 171
Group 18 = noble gases
STABLE. Very non reactive gaseous non-metals
ie: helium, neon, argon......
18
1 2
17
See page 171
THE PERIODIC TABLE
Where are the
following?
• Period
INCREASING REACTIVITY
• Group/Family
•Alkali metals
• Alkaline earth
metals
• Halogens
• Noble gases
See page 172
• Atoms gain and lose electrons to form bonds.
• The atoms become electrically charged particles called ions.
See page 173
• Atoms gain and lose electrons to form bonds.
• Metals lose negative electrons & become positive ions.
• Positive ions are called CATIONS.
See page 173
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
• Atoms gain and lose electrons to form bonds.
• Non-metals gain electrons and become negative ions
• Negative ions are called ANIONS
See page 173
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
• 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
• 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
• 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
• 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
There are 2 types of compounds:
• IONIC COMPOUND: metals lose electrons and non-metals 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
• 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
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
• Example covalent bond
• Hydrogen and fluorine form a covalent bond in the compound
hydrogen fluoride.
+
hydrogen
Hydrogen fluoride
fluorine
electrons are shared
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
• Dots representing electrons are placed around the element symbols at
the points of the compass (north, east, south, and west).
• Electron dots are placed singly until the fifth electron is reached then
they are paired.
See page 178
• 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 •
• •
••
•Cl •
• •
••
•Be •
• •
••
•Cl •
• •
••
•Cl •
• •
••
Each beryllium has two
electrons to transfer away,
and each chlorine can
receive one more electron.
Since Be2+ can donate two
electrons and each Cl– can
accept only one, two Cl– ions
are necessary.
2+
••
•
Be
•
••
•
•
••
–
•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
• 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