Periodic Trends Bellringer
• 1) Which has a larger size, a Mg atom or a Mg
ion?
• 2) smaller ionization energy, K or Br?
• 3) smaller size, F atom or an I atom?
• 4) smaller electronegativity, O or Se?
• 5) less shielding, Ca or Sr?
Answers
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1) Mg atom
2) K
3) F
4) Se
5) Ca
Ch. 7 Ionic and Metallic Bonding
Valence Electrons
• Valence electrons = electrons in the highest
occupied energy level of an element’s atoms
– This # largely determines the chem. properties of the
element
– To find the # of valence e-’s in an atom of a
representative element, look at its group #
• Group 1 has 1, group 2 has 2, group 13 or 3A has 3, …
• Electron dot structures (Lewis dot diagrams) =
diagrams that show valence electrons as dots
• *Draw Lewis dot structures*
Octet Rule
• Octet rule = atoms gain or
lose electrons to achieve a
stable level of usually 8
• atoms of metals tend to
lose their valence electrons
leaving a complete octet in
the next-lowest energy
level
• Atoms of some nonmetals
tend to gain electrons or to
share electrons w/ another
nonmetal to achieve a
complete octet
Cations
• Cations (+ charged atom) are formed when
they lose one or more valence electrons
(ionization) in order to become stable
– Usually lose from 1 to 3 valence electrons tops
– Can use electron configurations to illustrate the
point but let’s use electron dot diagrams for
ease
– Exceptions are due to having to lose or gain too
many valence electrons to achieve a noble gas
state, so:
• Some atoms attain a pseudo noble-gas electron
configurations – Cu - copper, Ag - silver, Au - gold,
Cd - cadmium
Two atoms are walking down the street.
Says one atom to the other, "Hey! I think I lost
an electron!"
The other says, "Are you sure??"
"Yes, I'm positive!"
Anions
• Anions (- charged atoms) are formed when
they gain 1 or more valence electrons
– Typically suffix is –ide
– Halide ions (F-, Cl-, Br-, I-) are halogens that gain 1
e– CP 7.1, PP 1-2 pg. 193, sect. assessment 7.1 pg.
193 3-11
Common ions
Cation Name
H 1+
Li 1+
Na 1+
K 1+
Cs 1+
Be 2+
Mg 2+
Al 3+
Ag 1+
Anion Name*
hydrogen
lithium
sodium
potassium
cesium
beryllium
magnesium
aluminum
silver
H 1- hydride
F 1- fluoride
Cl 1- chloride
Br 1- bromide
I 1- iodide
O 2- oxide
S 2- sulfide
*The root is given in color.
Ionic Compounds
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Anions and cations are held together by opposite charges.
Ionic compounds are called salts
Electrically neutral
Simplest ratio is called the formula unit
The bond is formed through the transfer of electrons (called
an ionic bond)
Electrons are transferred to achieve noble gas configuration
Most are crystalline solids at room temp.
High melting points – large attractive forces result in a very
stable structure
Good conductors of electricity when melted or dissolved in
water Overheard at the mall
Teen 1: Did you hear oxygen and magnesium got together?? Teen 2: OMg!
Your mama's so ugly
Your mama's so ugly...even Fluorine won't bind to her
Formulas
• Chemical formula = shows the kinds and #s of
atoms in the smallest representative unit of a
substance
– NaCl is the chem. formula for sodium chloride
• Formula unit = the lowest whole-number ratio
of ions in an ionic compound
– MgCl2, NaCl, AlBr3
– CP 7.2, PP 12-13 pg. 196,
– 7.2 sect. assessment pg. 199 14-22
Ch. 7 Bellringer
• Write the Lewis electron-dot symbol for each of
the following
– A) sodium
– B) fluorine
– C) magnesium ion (Mg2+)
• Write the chemical formula that results when the
following pairs of ions combine to form an ionic
bond
– D) Mn4+ and O2– E) Li1+ and Cl1-
Metallic Bonds and Properties
• Metallic bonds = consist of the attraction of the free-floating
valence electrons for the positively charged metal ions
• How atoms are held together in the solid
• Metals hold onto their valence electrons very weakly
• Think of them as positive ions floating in a “sea of electrons”
• Electrons are free to move through the solid
• Metals conduct electricity
• Malleable - hammered into shape (bend)
• Ductile - drawn into wires
• Electrons allow cations to slide by each other under pressure
• Metals are crystalline structures and atoms are arranged in
very compact and orderly patterns
*Cu vs. Cu compounds hammer
demo – pg. 202*
Alloys
• Alloys = mixtures composed of 2 or more
elements, at least one of which is metal
– Brass – copper and zinc
– Sterling silver – silver and copper
– Bronze – copper and tin
– Steel – Fe, Cr, and others
– Impt. b/c their properties are often superior to
those of their component elements (usually
cheaper as well)
*Making an alloy DEMO* - pg. 205
Ch. 8 Covalent Bonding
Molecules
• Covalent bond = formed by sharing electrons b/w 2 or
more atoms
• Molecule = a neutral group of atoms joined by covalent
bonds
• Diatomic molecule = a molecule consisting of 2 of the same
atoms
– H2, N2, O2, F2, Cl2, Br2, I2 are the diatomic molecules in nature
• Molecular compounds tend to have relatively lower melting
and boiling points than ionic compounds
• Most molecular compounds are composed of atoms to 2 or
more nonmetals
• *Make a table comparing covalent bonding and ionic
bonding*
Molecular Formula
• Molecular formula = the chemical formula for
a molecular compound
• Shows how many atoms of each element a
molecule contains
• H2O, CO2, C2H6, O2
• 8.1 Sect. assessment,
Pg. 216
1-6
Covalent Bonding
• Electron sharing usually occurs so that atoms attain the electron
configuration of noble gases
• Combos of atoms of the nonmetals and metalloids in 4A, 5A, 6A,
and 7A are likely to form covalent bonds
• Single covalent bond = 2 atoms sharing 1 pair of electrons
–
–
–
–
–
–
H2
2 dots in an electron dot diagram represents this bond
A dash in a structural formula represents this bond
A molecular formula does NOT show this bond only the # of atoms
Halogens form these bonds in their diatomic molecules
Unshared pair = pair of valence electrons not shared in an electron
dot diagram
– *CP 8.1, PP 7-8 pg. 220
Double and Triple Covalent Bonds
• Double covalent bond =
involves 2 shared pairs of
electrons
– oxygen
• Triple covalent bond = involves
3 shared pairs of electrons
– nitrogen
Coordinate covalent bonds
• The shared electron pair comes from one of
the bonding atoms
– CO – look at pg. 223
• Polyatomic ion = a tightly bound group of
atoms that has a positive or negative charge
and behaves as a unit
– NH4+
• *c.p. 8.2, p.p. 9-12 pg. 225
A sign outside the chemistry hotel reads "Great
Day Rates, Even Better NO3-'s"
Bond Dissociation Energies
• Bond dissociation energy = the energy
required to break the bond b/w 2 covalently
bonded atoms
– a large bond dissociation energy corresponds to a
strong covalent bond
– H2 = 435 kJ/mol, C-C single bond = 347 kJ/mol,
C=C double bonds = 657 kJ/mol, and triple bonds
= 908 kJ/mol
Resonance Structure
• Resonance structure = a structure that occurs
when it is possible to draw 2 or more valid
electron dot structures that have the same #
of electron pairs for a molecule or ion
– Double-headed arrows are used to connect
– Double bonds are usually shorter than single
bonds but they are the same lengths b/c it is an
avg. of the 2 structures
– resonance
Exceptions to the Octet Rule
• Cannot be satisfied in molecules whose total #
of valence electrons is an odd #.
– NO2, ClO2, NO
• Sometimes w/ an even # as well (Fewer or
more)
– BF3
– PCl5
– SF6
*8.2 sect. assessment 13-22 pg. 229*
*Exceptions to the octet rule: A resonance
hybrid teacher DEMO*
Molecular Orbitals
• Molecular orbitals = orbitals that apply to the
entire molecule
– just as an atomic orbital belongs to a particular
atom, a molecular orbital belongs to a molecule as
a whole
• Bonding orbital = a molecular orbital that can
be occupied by 2 electrons of a covalent bond
Sigma Bonds
• Sigma bonds = formed when 2 atomic orbitals combine
to form a molecular orbital that is symmetrical around
the axis connecting 2 atomic nuclei
– Atomic orbitals overlap end to end
– Two s orbitals can combine to form a molecular orbital
• H2
– Two p orbitals
• F2
• The attractions b/w electrons and nuclei of two atoms
overpower the repulsions b/w the 2 nuclei or b/w the 2
sets of electrons = covalent bond (stable molecule)
Pi Bonds
• Pi bond = a covalent bond in which the
bonding electrons are most likely to be found
in sausage-shaped regions above and below
the bond axis of the bonded atoms
– Orbitals overlap side by side
– Atomic orbitals in pi bonding overlap less than in
sigma bonding – weaker than sigma bonds
– A typical double bond consists of 1 sigma & 1 pi
bond, triple bond is 1 sigma & 2 pi bonds
– In special cases, they form w/o any sigma bonds
VSEPR Theory
• VSEPR theory (valence-shell electron-pair
repulsion theory) = the repulsion b/w
electron pairs causes molecular shapes to
adjust so that the valence-electron pairs stay
as far apart as possible
– Explains the actual 3-D shapes of molecules
– http://gold.chem.wwu.edu/sdgchem121/Docs/W
WUDocs/geometry.pdf
– http://www.youtube.com/watch?v=i3FCHVlSZc4
Hybrid Orbitals
• Hybridization = several atomic orbitals mix to
form the same total number of equivalent
hybrid orbitals
• Single bonds – one 2s and three 2p orbitals
mix to form four sp3 hybrid orbitals
• Double bonds – one 2s and two 2p orbitals
mix to form three sp2 hybrid orbitals
• Triple bonds – one 2s and 1 2p orbitals mix to
form two sp hybrid orbitals
Electronic Geometry
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2 electron densities – Linear
3 e- densities – Trigonal Planar
4 e- densities – Tetrahedral
5 e- densities – Trigonal Bipyramidal
6 e- densities – octahedral
Linear
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Atoms connected in a straight line
All molecules w/ 2 atoms and some w/ 3
180⁰ bond angle
HCl, CO2
Hybridization – sp
2 bonds/0 lone pairs
Trigonal Planar
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Triangular flat
120⁰ bond angle
BCl3
Hybridization – sp2
3 bonds/0 lone pairs
– Bent – 2 bonds/1 lone pair
• 118⁰
• SO2
Tetrahedral
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4 surfaces
109.5⁰ bond angle
CH4
Hybridization – sp3
4 bonds/0 lone pairs
– Pyramidal (trigonal pyramidal) – 3 bonds/1 lone pair
• 107⁰
• NH3
– Bent – 2 bonds/2 lone pairs
• 105⁰
• H2O
Trigonal Bipyramidal
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90, 120, and 180⁰ bond angles
PF5
Hybridization – sp4
5 bonds/0 lone pairs
– See-saw – 4 bonds/0 lone pairs
• 90, 120, and 180⁰
• SF4
– Tee-shaped – 3 bonds/2 pairs
• 90 and 180⁰
• ClF3
– Linear – 2 bonds/3 lone pairs
• 180⁰
• XeF2
Octahedral
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90 and 180⁰ bond angles
SF6
Hybridization – sp5
6 bonds/0 lone pairs
– Square pyramidal – 5 bonds/1 lone pair
• 90 and 180⁰
• BrF5
– Square Planar – 4 bonds/2 lone pairs
• 90 and 180⁰
• XeF4
• 8.3 sect. assessment pg. 236 23-29
The VSEPR Model
The Shapes of Some Simple ABn Molecules
SO2
O
C
..
O
N
S
O
O
O
Linear
O
Bent
F
S
F
O
Trigonal
planar
F
Trigonal
pyramidal
AB6
F
F
F
Cl
T-shaped
F
F
F
F
F
Square
planar
F
F
P
Xe
F
F
F
S
F
F
F
F
F
Trigonal
bipyramidal
Octahedral
Bond Polarity
• Nonpolar covalent bonds = equal sharing of
electrons
– H2, O2, N2, Cl2
• Polar covalent bond = unequal sharing of
electrons
– HCl, H2O
– The more electronegative atom attracts electrons
more strongly and gains a slightly negative charge,
the less electronegative atom has a slightly
positive charge
Table 8.3 pg. 238
• C.P. 8.3, P.P. 30-31 pg. 239
Polar Molecules
• Polar molecule = one end of the molecule is
slightly negative, the other is slightly positive
• Dipole = a molecule that has 2 poles w/opposite
charges
• The shape of a molecule and the polarity of its
bonds together determine whether the molecule
is polar or nonpolar
• Equal and opposite directions arrows cancel =
nonpolar
• Arrows same direction = polar
Attractions b/w Molecules
• Weaker than ionic or covalent bonds
• Van der Waals forces = weak intermolecular
attractions
– 1. Dipole interactions = intermolecular forces
resulting from the attraction of oppositely
charged regions of polar molecules
– 2. Dispersion forces = (weakest of all)
attractions b/w molecules caused by the
electron motion of one molecule affecting the
electron motion on the other through electrical
forces
• Hydrogen bonds = attractive forces in which a
hydrogen covalently bonded to a very
electronegative atom is also weakly bonded to
an unshared electron pair of another
electronegative atom
– The combo of this strongly polar bond and the
lack of shielding effect in a hydrogen atom is
responsible for the relative strength of
hydrogen bonds (strongest of all intermolecluar
attractions)
• Table 8.4 pg. 244
• *8.4 section assessment 32-38 pg. 244