Unit 5: Bonding

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Unit 5 Bonding
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Chemical 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

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• 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

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Polarity 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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Some 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

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Ionic 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

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Characteristics 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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• Some Ionic Characteristics
• High melting points
• Hard, brittle substances (solids with lattice
  structures)
• Good insulators
• Conduct electricity in solution (electrolyte)

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Example 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

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Example 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
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Example 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

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Examples of Ionic Bonding

• Potassium Iodide
• Barium Chloride
• Lithium Bromide

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Nomenclature for Ionic Bonding

• Monoatomic Ions
• Naming Cations and Anions
• K F-
• potassium ion fluoride ion

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Nomenclature for Ionic Bonding

• Binary Ionic Compounds
• MgBr2
• - Aluminum Oxide Ionic Formula?

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Nomenclature 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

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Examples using Roman Numerals

• Iron (II) chloride
• 2. Manganese (III) nitride
• 3. Tin (IV) oxide
• Cr(OH)3
• PbCl2
• Co2S3

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Common Polyatomic Ions

• Must know the polyatomic ions that are circled
• Lets take a look at them
• This is very important for this unit

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Naming 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.

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Polyatomic Ion Examples

• NaNO3
• Mg(OH)2
• Ca(NO2)2
• Potassium carbonate
• Ammonium phosphate
• Calcium Acetate

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Metallic 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.

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• 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".
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Covalent 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

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• 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.

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• 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

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Method 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

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Multiple 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.
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Some examples

• CH4
• CO2
• NH3
• PSCl
• BF3

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Try some more

• CF4
• PF2Cl
• SiOBr2
• NCl3
• H2O
• C2H6 (hint no central atom)

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• 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

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Molecular 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.

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2 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
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Lets look at CCl4

• Lewis Structure

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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Bent 2 atoms bonded to the central atom 2 sets
of lone pairs
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Molecular 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

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Hybridization

• 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

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Rules 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

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Examples
sp sp2 sp3
Carbon dioxide Boron trifluoride Carbon
tetrachloride
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Examples 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

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For 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?

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Can 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

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Intermolecular 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.

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• 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.

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Gas, 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.

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• 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

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What types of intermolecular forces are there?
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What 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.

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Ionic 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.

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Van der Waals Forces

• Include
• Hydrogen Bonding
• Dipole-dipole
• Dipole-Induced Dipole
• London Dispersion Forces

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Covalent 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.

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Hydrogen 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

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Dipole-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).

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London 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.

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London 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.

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Review 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)

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Answers 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.