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Chemical Bonds
Ch. 6 Atomic Structure and
Chemical Bonds
Atom Structure

The nucleus
– Has a positive charge
– Contains protons and neutrons

The electron cloud
– Surrounds nucleus
– Contains electrons with a negative charge
traveling in energy levels

Each element has a different structure
consisting of a specific # of protons,
neutrons and electrons
Electron Arrangement

Each level (shell) that contains electrons has
a specific amount of energy
– The farther from the nucleus, the more energy and
the more electrons the shell can hold

Electron shell capacities
– 1st shell : energy level 1; 2 electrons ; least
energy ; hardest to remove
– 2nd shell : energy level 2 ; 8 electrons
– 3rd shell : energy level 3 ; 18 electrons (stable with
8- octet rule)
– 4th shell : energy level 4; 32 electrons ; most
energy ; easiest to remove
How Periodic Table predicts
electron configuration

Group number (top of column) shows the
number of electrons in the outermost shell for
representative elements
– Ex. Group 1 has 1 outer shell electron. Group 2
has 2, Group 13 has 3, Group 14 has 4 and so oncalled valence electrons

Each period (row) ends in a stable element
(noble gas). Elements in a period have
increasing melting points and densities in a
predictable pattern.
 Periodic Table
Element Families

Noble gases – Group 18
– Stable
– Almost completely unreactive
– Can be forced to react

Halogens – Group 17
– Seven outer shell electrons
– Very reactive, tend to gain one electron

Alkali metals – Group 1
– One outer shell electron
– Very reactive tends to lose one electron
Trends http://www.gpb.org/chemistryphysics/chemistry/403

Atomic number
– Increases as you go down and left to right

# of valence electrons
– Increases from left to right (representative
elements)

# of energy levels
– Increase as you move down a column (family,
group)

Atomic size
– atomic size is determined by how much space the
electrons take up.
– As you go down a group, the size increases. As
you go across a period the size decreases.

Ionization energy – tendency to lose electrons
– A low ionization energy means that it is easy for an atom to
lose electrons. A high ionization energy means that it is
hard for an atom to lose electrons.
– it gets harder for atoms to lose electrons as you go across
the periodic table (increased ionization energy) and it gets
easier for atoms to lose electrons as you go down the
periodic table.

Electron affinity - Tendency to gain electrons
– Thus the tendency of atoms to gain electrons increases as
we go from left to right across the periodic table. At least it
increases until we get to the inert gases. There it drops off to
zero because there is no room for additional electrons in the
valence energy level.
Electron Dot diagrams

Symbol for the element surrounded by
the number of electrons in it’s outer
shell
N
Chemical Bonds

Electrons are shared to fill the outermost
energy level
 The positively charged nucleus of each atom
attracts the negatively charged electrons at
the same time
 Can be represented by an electron-dot
diagram
 If an atom shares more than one electron it is
called a double (2 electrons) or triple (3
electrons) bond.
Polar Covalent Bonds
When electrons are shared unevenly in
a covalent bond it is a polar bond.
 The uneven sharing causes one end of
the molecule to be more negative

– Ex. H20 and HCl

When electrons are shared evenly it is a
nonpolar covalent bond
Molecules
The combination of atoms formed by a
covalent bond
 has all properties of that substance
 represented by a chemical formula
 subscripts show the number of atoms of
each element in the molecule

– Ex. H2O – 2 hydrogen atoms, 1 oxygen
atom in a molecule of water
Ionic Bonds

Bonding that involves a transfer of electrons
 One gains electrons one loses electrons
(electron affinity, ionization energy)
–
–
–
–

Fluorine (F) 7 valence electrons gains one (F-)
Sodium (Na) 1 valence electrons loses one(Na+)
bonding will give both filled outer levels
NaF sodium fluoride
Placement of ions results in a repeating
crystal lattice
Metallic Bonds
Metals give up electrons easily
 Outer electrons form a common
electron cloud - “sea of electrons”
 High melting point - good conductors
 Nuclei of atoms of metals are
surrounded by freely moving electrons

Predicting Bond Types
Elements at the left and in the center of
periodic table tend to lose electrons
easily-metals
 Elements at the right tend to gain
electrons easily - nonmetals
 Metal to nonmetal compound - ionic
 Nonmetal to nonmetal - covalent
 Metal to metal - metallic

Oxidation Number
The # of electrons an atom gains, loses,
or shares when it bonds
 describes its combining capacity
 Ex. Na=+1, Mg =+2, Cl = -1, O = -2
 Can be used to predict how many
atoms will combine and the formula for
the compound
 The sum of the oxidation #’s of the
atoms in a compound must be 0.
