Chemical Bonding

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Chemical Bonding
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Chemical Bonds

• Form when atoms or ions are strongly attached to
  one another
• Defined as forces of attraction that hold two
  atoms together and allows them to function as a
  unit
• Three main types
• Ionic
• Covalent
• Metallic

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The Octet Rule

• Atoms tend to gain, share, or lose electrons in
  order to obtain a full set of valence electrons
  (in most cases this equals 8)
• An octet of electrons consists of full s and p
  sublevels on an atom.
• Exceptions transition elements and rare earth
  elements

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Determining Types of Bonds

• Determined by electronegativity differences -
  difference determines the percentage ionic or
  percentage covalent character
• Type of bonds are determined by which percentage
  is prevalent

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Ionic Bonding

• Occurs when electrons are transferred from one
  atom to another, forming two ions
• Cation positive ion Anion negative ion
• The ions stay together because of electrostatic
  attractions ionic bond
• Ionic bonds NEVER form molecules
• Ionic bonds form easily between alkali metals and
  halogens
• All ionic compounds are electrically neutral

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Example
-

Na
Cl
Na

Cl

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Properties of Ionic Compounds

• Ionic compounds do not form molecules they form
  a crystal lattice
• Formula unit the simplest collection of atoms
  from which an ionic compounds formula can be
  established
• The ions lower their potential energy forming
  orderly, 3-D array in which the cations and
  anions are balanced
• Formula units arrange themselves in repeating
  patterns

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This is a crystal of CaCl2. Each ion is held
rigidly in place by strong electrostatic forces
that bond it to several oppositely charged ions
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Other Properties

• Normally form between metals and nonmetals
• Ionic compounds have ions that form very strong
  bonds, which makes them hard and brittle
• They have high melting points and high boiling
  points
• Most are solids at room temperature

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Properties continued

• When dissolved in water, the solution will
  conduct electricity
• Conduct electricity in the molten state, but will
  not conduct electricity in the solid state
• Tend to be soluble in water
• Crystallize as sharply defined particles
• One atom has a low electronegativity and a low
  electron affinity
• The other is vise versa

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Types of Ions

• There are two types of ions
• Monatomic cation or anion that consists of a
  single atom. Examples Na and Cl-
• Polyatomic two or more atoms that act as a
  single ion (or particle). Examples (CO3)2- and
  (OH)-

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Monatomic Ions
Group Atoms that commonly form ions Charge on ions
1A H, Li, Na, K, Rb, Cs 1
2A Be, Mg, Ca, Sr, Ba 2
3A B, Al 3
5A N, P, As 3-
6A O, S, Se, Te 2-
7A F, Cl, Br, I 1-
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Common Polyatomic Ions
Ion Name Ion Name
NH4 Ammonium NO2- Nitrite
NO3- Nitrate OH- Hydroxide
CO32- Carbonate SO42- Sulfate
O22- Peroxide C2H3O2- Acetate
SO32- Sulfite ClO3- Chlorate
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Metallic Bonds

• Metal members of the representative groups have
  some, if not all, vacant p orbitals
• Many of the transition metals contain vacant d
  orbitals - allow electrons to roam freely
  throughout the metal
• Electrons are delocalized - they do not belong to
  any one atom (sea of mobile electrons)
• Metallic bonding is the result of the attraction
  between metal atoms and the surrounding sea of
  electrons
• Responsible for metallic properties such as
  conductivity, malleability, ductility, and luster

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Forming Covalent Bonds

• covalent (co sharing valent outermost shell)
  - when electrons are shared between two nuclei
• AKA molecular bonds
• molecules a group of atoms held together by
  covalent bonds

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Electron Pairs in Covalent Bonds

• Unshared pairs pairs of electrons that do not
  participate in bonding and belong to only one
  atom - also called lone pairs
• Bonding pairs pairs of electrons being shared
  between two atoms thus creating a covalent bond
• Electrons are not always equally shared
• Unequal/equal sharing is determined by
  electronegativity differences

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Covalent Bonds

• Electronegativity - tendency to attract
  electrons in a chemical bond
• polar having opposite ends
• covalent bonds in which the bonding electrons are
  more strongly attracted by one of the bonding
  atoms
• nonpolar not having opposite ends
• covalent bonds in which the bonding electrons are
  shared equally between the 2 bonding atoms

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Properties of Covalent Compounds

• Typically low melting points
• Most are gases, liquids, or very soft solids at
  room temperature
• Do not conduct electricity
• Are brittle when solids
• Typically form between nonmetals

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Types of Covalent Bonds

• Covalent bonds, unlike ionic bonds, can form
  multiple bonds
• Single bonds two atoms share two electrons (one
  pair)
• Double bonds two atoms share four electrons
  (two pair)
• Triple bonds two atoms share six electrons
  (three pair)

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Types of Covalent Bonds

• Coordinate covalent bonds a covalent bond in
  which a single atom contributes both of the
  electrons to a shared pair
• Covalent bonds are separated into two types of
  bonds
• Sigma bonds (s)
• pi bonds (p)

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Sigma and Pi Bonds

• Sigma bonds
• Formed along the horizontal axis between two
  atoms
• The primary bonds
• Pi bonds
• formed above and below the horizontal axis
  between two atoms
• The secondary bonds

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Bond Composition

• Single bonds - consist of only one sigma bond
• Double bonds - consist of one sigma bond and one
  pi bond
• Triple bonds - consist of one sigma bond and two
  pi bonds
• http//Sigma and Pi Bonds

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Bond Guide
Percent Ratio Electronegativity Difference Occurrence
0 Ionic/100 Covalent 0.00 Rarely
50 Ionic/50 Covalent gt1.7
100 Ionic/0 Covalent gt 3.3 Never
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Polarity

• Remember Polar vs Nonpolar are also determined
  by electronegativity differences
• Polar Bonds 0.4 lt x lt 1.7 ( x represents
  electronegativity difference)
• Nonpolar Bonds x lt 0.4
• Areas of partial charge build up because a shift
  in electron charge density occurs
• Shift itself indicated by use of arrows along the
  bond
• Partial charge indicated by uses of

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Bond Length

• Bond length average distance between bonded
  atoms
• Measured from center of one nucleus to center of
  the neighboring nucleus
• Different pairs of atoms form bonds of different
  length
• Atomic radius of each atom participating in the
  bond therefore directly affects bond length
• Not fixed because atoms vibrate through the bond
  in a spring-like fashion
• Multiple bonds are shorter than single bonds

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Bond Angle

• Bond angle the angle between the two bond axes
• Bond angles are also not fixed because of the
  atom vibration

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Bond Energy

• Bond energy energy required to break a bond
• indicates bond strength
• usually reported in units of kilojoules/mole
• the closer the atoms the greater the bond energy
  required to separate them
• more energy is required to break multiple bonds
  than single bonds
• also called bond dissociation energy