Title: Unit 5: Bonding
1Unit 5 Bonding
2Chemical Bond
- a mutual electrical attraction between the nuclei
and valence electrons of different atoms that
binds the atoms together - 2 types of Bonding
- Ionic Bonding
- results from an electrical attraction b/w large
numbers of cations and anions - Metals completely give up electrons/Nonmetals
gain - Between a metal and a nonmetal
3- Covalent Bonding
- Results from the sharing of electron pairs
between two atoms - B/w nonmetal and a nonmetal
- Nonpolar Covalent a covalent bond in which the
bonding electrons are shared equally by the
bonded atoms, resulting in a balanced
distribution of electrical charge - Polar Covalent a covalent bond in which bonded
atoms have an unequal attraction for the shared
electrons
4Polarity and Electronegativity
- Polar an uneven distribution of charge
- Remember Electronegativity
- What are the trends on the periodic table?
- Which element is the most electronegative?
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7Some Examples
- Using electronegativity values, designate whether
the bonds would be nonpolar covalent, polar
covalent, or ionic. - C and H
- C and S
- O and H
- Na and Cl
- Cs and S
8Ionic Bonding
- Ionic Compound composed of positive and
negative ions that are combined so that the
number of positive and negative charges are equal - Does NOT form a molecule, instead a formula unit
- Formula unit simplest collection of atoms from
which an ionic compounds formula can be
established - Ionic compound must be electrically neutral
9Characteristics of Ionic Bonding
- Most compounds exist as crystalline solids
- Combine in an orderly arrangement known as a
crystal lattice - Lattice energy the amount of energy released
when 1 mole of an ionic crystalline compound is
formed from gaseous ions.
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11- Some Ionic Characteristics
- High melting points
- Hard, brittle substances (solids with lattice
structures) - Good insulators
- Conduct electricity in solution (electrolyte)
12Example Sodium Chloride, NaCl
- NaCl is a compound, but how?
- Form its most common cation and anion
- Na Cl-
- In order to be neutral you only need 1 Na 1 Cl
13Example CaF2
- Break into anions and cations
- Ca2 F-
-
In order to be neutral, there must be 2 F- ions.
??
- What happens?
- Ca2 donates two electrons to F-
Which will make CaF2
14Example Magnesium Chloride
- Break into cation and anion
- Mg2 Cl-
- Becomes MgCl2
- Notice that the charges are criss-crossing
- Mgs charge becomes a subscript for Cl
- Cls charge becomes a subscript for Mg
15Examples of Ionic Bonding
- Potassium Iodide
- Barium Chloride
- Lithium Bromide
16Nomenclature for Ionic Bonding
- Monoatomic Ions
- Naming Cations and Anions
- K F-
- potassium ion fluoride ion
17Nomenclature for Ionic Bonding
- Binary Ionic Compounds
- MgBr2
- - Aluminum Oxide Ionic Formula?
18Nomenclature for Stock System
- Roman numeral
- equal to the ion charge
- Does not equal the subscript
- Lower numeral ous ending
- Higher charge ic ending
- Example Ferrous ion, Iron (II) ion, Fe2
- Ferric ion, Iron (III) ion, Fe3
19Examples using Roman Numerals
- Iron (II) chloride
- 2. Manganese (III) nitride
- 3. Tin (IV) oxide
- Cr(OH)3
- PbCl2
- Co2S3
20Common Polyatomic Ions
- Must know the polyatomic ions that are circled
- Lets take a look at them
- This is very important for this unit
21Naming Polyatomic Ions
- Think of the -ate ending as base. (chlorate
will be our example, ClO3- - If chlorate loses an oxygen, the name becomes,
chlorite, ClO2- - If chlorate loses 2 oxygens, the name becomes,
hypochlorite, ClO- - If chlorate gains 1 oxygen, the name becomes,
perchlorate, ClO4- - This applies to all of polyatomic ions containing
oxygen.
22Polyatomic Ion Examples
- NaNO3
- Mg(OH)2
- Ca(NO2)2
- Potassium carbonate
- Ammonium phosphate
- Calcium Acetate
-
23Metallic Bonding
- Metals tend to have high melting points and
boiling points suggesting strong bonds b/w atoms. - Example Sodium
- All of the 3s orbitals on all of the atoms
overlap to give a vast number of molecular
orbitals which extend over the whole piece of
metal.
24- The electrons can move freely within these
molecular orbitals, and so each electron becomes
detached from its parent atom. - The electrons are said to be delocalized.
- The metal is held together by the strong forces
of attraction between the positive nuclei and the
delocalised electrons. (In a way, they act like
cement holding the positive metal ions in
relatively fixed positions)
described as "an array of positive ions in a sea
of electrons".
25Covalent Bonding
- Lewis reasoned that an atom might attain a noble
gas electron configuration by sharing electrons - A chemical bond formed by sharing a pair of
electrons is called a covalent bond
26- The diatomic hydrogen molecule (H2) is the
simplest model of a covalent bond, and is
represented in Lewis structures as
- The shared pair of electrons provides each
hydrogen atom with two electrons in its valence
shell (the 1s) orbital. - In a sense, it has the electron configuration of
the noble gas helium.
27- When two chlorine atoms covalently bond to form
Cl2, the following sharing of electrons occurs
Each chlorine atom shared the bonding pair of
electrons and achieves the electron configuration
of the noble gas argon.
- In Lewis structures the bonding pair of
electrons is usually displayed as a line, and the
unshared electrons as dots
28Method of Drawing Lewis Structures
- Sum the valence electrons of all atoms. Subtract
an electron for a () sign add an electron for
a (-) sign. - Put the atom wanting the most bonds in the
middle - An atom will form one bond for each electron it
wants. - Put the remaining atoms around the central atom,
giving them the number of bonds they want. - Fill in pairs of electrons until every atom has
eight electrons - Exceptions H, B, Be
29Multiple Bonds
- The sharing of a pair of electrons represents a
single covalent bond, usually just referred to as
a single bond - In many molecules atoms attain complete octets by
sharing more than one pair of electrons between
them. - Two electron pairs share a double bond
- Three electron pairs share a triple bond
Because each nitrogen contains 5 valence
electrons, they need to share 3 pairs to each
achieve a valence octet.
30Some examples
31Try some more
- CF4
- PF2Cl
- SiOBr2
- NCl3
- H2O
- C2H6 (hint no central atom)
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34- Some Covalent Characteristics
- Depends on polarity and whether the molecule is a
covalent network (these can drastically influence
characteristics) - Nonpolars generally are gases with low boiling
points, are not good conductors of electricity
(nonelectrolyte), and are volatile - Polars generally are liquids with higher boiling
points. Highly polar ones can conduct
electricity - Covalent networks can have very high melting
points and are usually nonconductors
35Molecular Geometries
- VSEPR Theory valence shell electron pair
repulsion - Repulsion between the sets of valence level
electrons surrounding an atom causing these to be
oriented as far apart as possible.
362 atoms bonded to central atom No lone pairs Type
of molecule AB2
3 atoms bonded to central atom No lone
pairs Type of molecule AB3
37Lets look at CCl4
However, the Lewis structure provides no
information about the shape of the molecule
Atoms Bonded to Central 4 Lone pairs 0 Type
of molecule AB4
Carbon tetrachloride is tetrahedral in structure
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39Bent 2 atoms bonded to the central atom 2 sets
of lone pairs
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45Molecular Dipole Moments
- Are the following Polar or nonpolar.? Draw the
structure in the correct geometry with the
correct bond angles labeled, with dipoles and
? ?- signs in order to explain your answer. - BF3
- CH2O
- CCl4
- CH3Cl
46Hybridization
- Hybridization is the process of combining two or
more atomic orbitals to create new orbitals,
called hybrids, that will fulfill the geometric
demands of the system. - Use the s,p,d f nomenclature
47Rules for hybridization
- Count each thing that is touching the atom you
want to figure the hybridization for. - Lone pairs count once
- Single bonds count once
- Double bonds count once
- Triple bonds count once
48Examples
sp sp2 sp3
Carbon dioxide Boron trifluoride Carbon
tetrachloride
49Examples for hybridization
- Are the following Polar or nonpolar? Draw the
structure in the correct geometry with the
correct bond angles labeled, with dipoles and
? ?- signs in order to explain your answer.
What is the hybridization of the central
molecule? - BF3
- CH2O (find hybridization for C and O)
- CCl4
- CH3Cl
50For next class
- Notes on intermolecular forces
- Bring completed homework
- Quiz next time I see you
- Lab on the 1st and 3rd
- Just wondering Have you been studying your
objectives (on your unit plan) for the test
coming up?
51Can nonpolar molecules contain polar bonds?
- First assign dipoles and ? and ?- signs
- Each C-Cl bond is polar
- Overall though, CCl4 is non polar, due to its
symmetrical shape. - The more polar a molecule is the greater
separation of charge - The more polar a molecule is the more attracted
it will be to another polar molecule
52Intermolecular Forces
- Intermolecular forces are the forces of
attractions that exist between molecules in a
compound. - These cause the compound to exist in a certain
state of matter solid, liquid, or gas and
affect the melting and boiling points of
compounds as well as the solubilities of one
substance in another.
53- The stronger the attractions between particles
(molecules or ions), the more difficult it will
be to separate the particles. - In a solid the kinetic energy of the molecules is
small compared to the strength of the
intermolecular forces, so each molecule can only
move short distances around a fixed position. - In a liquid the kinetic energy of the molecules
is comparable to the intermolecular forces
between them, so the molecules can move around,
but usually stay within a molecular diameter of
each other. - In a gas the kinetic energy of the molecules is
much greater than the intermolecular forces
between them and the molecules move freely,
colliding far less frequently than in a liquid.
54Gas, Liquid, vs. Solid
- For each phase, a molecule or compound would most
likely be polar or nonpolar and why? - Solids
- The stronger the intermolecular forces the closer
the molecules pull themselves. - More density greater likelihood of being a solid
at room temperature. - Ionic compounds have a very large separation of
charge, therefore they have a strong force of
attraction ? likely to be solid at room
temperature.
55- Liquids
- At room temperature the atoms within a liquid are
not held as closely as those within solid. - Mostly polar, especially water due to hydrogen
bonding. - Gases
- At room temperature gases are nonpolar
- There is no separation of charge between most gas
molecules - Ex. O2, CH4
56What types of intermolecular forces are there?
57What causes intermolecular forces?
- Molecules are made up of charged particles
nuclei and electrons. - When one molecule approaches another there is a
multitude of interactions between the particles
in the two molecules. - Each electron in one molecule is subject to
forces from all the electrons and the nuclei in
the other molecule.
58Ionic Bonding Ion-Dipole Forces
- When an ionic substance dissolves in a polar
solvent (that is, a solvent whose molecules have
a permanent dipole moment) the majority of the
solvent molecules orient themselves with the
oppositely charged end of the solvent molecule
near an ion. - This attraction between the ions and the solvent
molecules can win out over the attraction of the
ions to each other, allowing the substance to
stray in solution. - For an insoluble ionic substance this is not the
case. Ion - dipole forces are responsible for the
dissolution of ionic substances in water.
59Van der Waals Forces
- Include
- Hydrogen Bonding
- Dipole-dipole
- Dipole-Induced Dipole
- London Dispersion Forces
60Covalent Bonding Dipole-Dipole Forces
- If two neutral molecules, each having a permanent
dipole moment, come together such that their
oppositely charged ends align, they will be
attracted to each other. - In a liquid or solid these alignments are favored
over those where like-charged ends of the
molecules are close together and hence repel each
other.
61Hydrogen Bonding
- Hydrogen bonds form only between a limited number
of elements bonded in a specific sequence. - So a necessary condition for hydrogen bonding is
that one molecule must contain a H bonded to
either N, O, or F and the other molecule must
contain either N, O, or F. - If a hydrogen bond can form between a pair of
molecules it will be stronger than other
intermolecular forces between the molecules. - Hydrogen bonding is responsible for the
unexpectedly high boiling point of water
62Dipole-Induced Dipole Forces
- A polar molecule (lower left) carries with it an
electric field and this can induce a dipole
moment in a nearby non-polar molecule (lower
right). - This will cause an attraction between the
molecules. - This type of force is responsible for the
solubility of oxygen (a non-polar molecule) in
water (polar).
63London Dispersion Forces
- Arise from the temporary variations in electron
density around atoms and molecules. - Nonpolar molecules have a certain minimum
symmetry to their average shape and electron
distribution. Picture 1 at left depicts two
nonpolar molecules. - However at any instant the electron distribution
around an atom or molecule will likely produce a
dipole moment (figure 2 on left) which will
average out to zero over a period of time.
64London Dispersion Continued
- But even a temporary dipole moment can induce a
(temporary) dipole moment in any nearby molecules
(picture 3 on left) causing them to be attracted
to the first molecule. - Unlike forces between molecules with permanent
dipole moments, dispersion forces always act to
attract the molecules to each other regardless of
the relative orientation of the molecules - Molecules containing large atoms (e.g. bromine or
iodine) have large polarisabilities and so give
rise to large dispersion forces. This explains
the increasing melting and boiling points of the
halogens going down that group of the periodic
table.
65Review Examples
- For each of the molecules below, list the types
of intermolecular forces which act between pairs
of these molecules, hybridization of central
atom, name of molecule, shape of molecule,
polarity, dipoles, and partial negative/positive
charge. - (a) CH4 (b) PF3 (c) CO2 (d) HCN, (e) HCOOH
(methanoic acid)
66Answers to Examples
- (a) CH4 is a tetrahedral molecule it does not
have a permanent dipole moment it does not
contain N, O, or F. Therefore only dispersion
forces act between pairs of CH4 molecules. - (b) PF3 is a trigonal pyramidal molecule (like
ammonia, the P has a single lone pair) it does
have a permanent dipole moment it does contain
F, however the fluorine is not bonded to a
hydrogen. Therefore dispersion forces and
dipole-dipole forces act between pairs of PF3
molecules. - (c) CO2 is a linear molecule it does not have a
permanent dipole moment it does contain O,
however the oxygen is not bonded to a hydrogen.
Therefore only dispersion forces act between
pairs of CO2 molecules. - (d) HCN is a linear molecule it does have a
permanent dipole moment it does contain N,
however the nitrogen is not directly bonded to a
hydrogen. Therefore dispersion forces and
dipole-dipole forces act between pairs of HCN
molecules. - (e) HCOOH is a non-linear molecule it does have
a permanent dipole moment it does contain O, and
the oxygen is directly bonded to a hydrogen.
Therefore dispersion forces, dipole-dipole forces
and hydrogen bonds act between pairs of HCOOH
molecules.