Bonding An Introduction to Chemical Reactions

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BondingAn Introduction to Chemical Reactions

• Pg. 156-215

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Chemical Bonds

• Properties of many materials can be understood in
  terms of their microscopic properties
• connectivity between atoms,
• three dimensional shape of the molecule.
• When atoms are strongly attracted to one another
  chemical bond
• What causes this attraction between atoms?

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Electrostatics

• Electrostatics- attraction and repulsion
  determines bonding between atoms and forces of
  attraction that can exist between molecules.

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Coulombs Law

• q1 q2 are charges on particles 1 2
• D is the distance between the particles
• ke constant
• As distance between charges increases, the
  electrostatic force __________.
• As the charge on the particles increases, the
  electrostatic force __________.

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Questions to Consider

• Where do these charges exist in an atom?
• How does the organization of the atoms electrons
  affect this electrostatic force?
• Make a connection between reactivity of atoms in
  the periodic table and the organization of
  electrons using this concept of electrostatic
  force.

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Bond Types

• 1. Ionic Bonds- __________________________________
  ______ These oppositely charged ions are
  attracted to each other through electrostatic
  forces.
• 2. Covalent Bonds-
• 3. Metallic Bonds-

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Metallic Bonds

• Positive ions in a sea of mobile electrons.

Delocalized Valence Electrons
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Metallic Bonds

• Form between two or more metals
• Atoms of metals achieve stability by sharing
  their valence electrons. Delocalized valance
  electrons.
• Metallic bonds are the attractive forces between
  fixed positive ions and the moving valence
  electrons of the metal.

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Composition of Selected Alloys
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Ionic Bonds

• Static electricity and the clothes dryer
• Static electricity is the basis for ionic bonds.
• Octet Rule dictates that some substances gain
  electrons- __________, while others lose
  electrons- ___________.
• Positive and negative ions are attracted to one
  another.

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Ionic Bonds
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Characteristics of Substances with Ionic Bonds

• Composed of _______
• Have ________ melting points
• Solids at room temperature, many soluble in water
• ________________________________________________
• Tend to be ______________

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Covalent Bonds

• Formed by a shared pair of electrons between two
  atoms.
• Molecule

Glycine- AA
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Types of Formulas

• Molecular formula- indicates the number of atoms
  that are in a single molecule of a compound.
  C6H12O6
• Empirical formula- indicates the lowest whole
  number ratio of atoms in a molecule. CH2O
• Structural Formulas- specifies which atoms are
  bonded to each other in a molecule

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Structural Formulas- Lewis Structures

• Valence electrons are indicated around the symbol
  for the element

Oxygen has 6 valence electrons
Nitrogen has 5 valence electrons
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Drawing Lewis Structures

• Imagine each side (top, bottom, left, right) of
  the symbol of the element can hold 2 electrons
  for a total of 8 electrons.
• Each side will hold one electron first, then will
  double up.
• In covalent bonding the number of single electron
  sides (unpaired electrons) indicates the number
  of covalent bonds the atom must have to satisfy
  its octet.

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• Oxygen has 6 valence electrons.
• Two unpaired electrons means that oxygen must
  form two bonds to satisfy its octet.
• Draw the Lewis structure for the following
• Chlorine
• Phosphorus
• Carbon

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Lewis Structures

• Atoms share electrons to fill their octets.
• A solid line indicates a shared pair of
  electrons.
• Dots are used to indicate unshared pairs of
  electrons.

Formation of a single covalent bond
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Double and Triple Bonds

• A unique characteristic of covalent compounds is
  their ability to form multiple bonds between two
  atoms.
• Refer back to the Lewis Structures for nitrogen
  and oxygen.
• Nitrogen needs to share three electrons
• Oxygen needs to share two electrons.

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Technique for Drawing Lewis Structures

• Determine the number of valence electrons in each
  atom making up the molecule
• Add the valence electrons and divide by two
• Draw the skeleton. If carbon is present, place
  it at the center of the molecule.
• Distribute the pairs of electrons around the
  skeleton to satisfy each atoms octet. (Remember
  Hydrogen only needs two electrons to fill its
  octet.)

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Practice

• Draw Lewis Structures for the following
  compounds
• Ammonia
• Ethyne- C2H2
• Carbon Dioxide
• HCN

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Exceptions to the Octet Rule

• Atoms with more than an octet
• SF4
• Molecules with an odd number of electrons
• NO
• Generally short lived, unstable molecules

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Properties of Molecular Compounds

• Composed of 2 or more ___________
• ___________ electrons in bond formation
• Can be solids, liquids, or gases at room
  temperature.
• Some are soluble in water, others are not.
• Tend to be ___________________ conductive.
• Generally have _________ melting points.

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Questions to Consider for Lewis Structures

• What does it mean to share electrons in the
  formation of a bond.
• In your experience, is sharing always equal?
• Pick a bond in your Lewis structure and decide if
  the sharing of electrons is equal or unequal.
  Why is it so?
• How might this sharing affect the physical and
  chemical characteristics of the molecule?

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Covalent Bonds- Are the Atoms Really Sharing
Electrons?
Chlorine
Hydrogen
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Covalent Bond Types

• Polar Covalent Bonds- electrons in bond are
  ________________.
• Nonpolar Covalent Bonds- electrons in bond
  ___________________________________.

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Polar Covalent Bonds

•   When a bond is classified as polar covalent
  (H-O), the atom with the higher electronegativity
  has the greater attraction for the shared
  electrons
• As a result, a charge unbalance is produced in
  the molecule? by H and by O
•   Dipole charge unbalance
• d H O d-
•  
• The positive and negative ends of the dipole
  are not real charges (such as positive and
  negative ions) because no electrons have actually
  been transferred between the atoms. The dipole
  represents only an unbalanced charge distribution
  along the bond.

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Nonpolar Covalent MoleculesBrINCl HOF Elements

• Diatomics- elements that can combine with
  themselves in a nonpolar covalent molecule to
  form a stable compound.
• Memorize!

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Electronegativity
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Bond Type by Electronegativity
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Waters polarity allows it to pull at the ions in
an ionic crystal.
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Metallic Vs Ionic Bonding

• Much easier to deform materials with metallic
  than with ionic bonding. Why?

Ag (s)
NaCl (s)

• Sliding atom planes over each other (deformation)
  very unfavorable energetically in ionic solids!
• ? metals are ductile ceramics (ionic) are
  brittle

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Intermolecular Vs Intramolecular Forces
Intermolecular forces are forces ___________
molecules. Arises from interaction between
dipoles. Bond Polarity
Intramolecular forces ____________________________
_____

• Intermolecular vs Intramolecular
• 41 kJ to vaporize 1 mole of water (inter)
• 930 kJ to break all O-H bonds in 1 mole of water
  (intra)

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Types of Intermolecular Forces

• Dipole-Dipole Forces
• Hydrogen Bonding Forces
• London Dispersion Forces
• Tend to be less than 15 as strong as covalent or
  ionic bonds.

Measure of intermolecular force boiling
point DHvap melting point DHfus DHsub
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Intermolecular Forces
1. Dipole-Dipole Forces
solid
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Intermolecular Forces
2. Hydrogen Bond a special dipole-dipole
interaction between the hydrogen atom in a polar
N-H, O-H, or F-H bond and an electronegative O,
N, or F atom.
A B are N, O, or F
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Intermolecular Forces
3. London Dispersion Forces
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London Dispersion Forces among nonpolar molecules
instantaneous dipoles
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Chemical Reactions

• A process in which one or more substances are
  converted into new substances with different
  physical and chemical properties.
• Reactant- a substance that enters into a chemical
  reaction.
• Product- a substance that is produced by a
  chemical reaction.

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The Reason for Reactions

• During a chemical reaction, new substances are
  produced as existing bonds are broken, atoms are
  rearranged, and new bonds are formed.
• Substances undergo chemical reactions with other
  substances _____________________

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Chemical Equations

• Describes what happens in a chemical reaction-
  similar to mathematic equations.
• Word Equations- give the names of the reactants
  and the products.
• Calcium oxygen yields calcium oxide
• Formula Equations-chemical symbols replace the
  names of the reactants and products.
• Ca O2 ? CaO

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Law of Conservation of Mass and Balancing
Chemical Equations

• Matter is neither created nor destroyed during a
  chemical reaction. Therefore, all the atoms that
  were present at the start of the reaction must be
  present at the end of the reaction.

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Balanced?Ca O2 ? CaO

• Coefficients are used in chemical equations to
  balance an equation.
• Subscripts cannot be changed once the compound is
  written. Changing the subscript would change the
  compound!
• Ca O2 ? CaO
• A coefficient of 2 is placed in front of calcium
  and calcium oxide to balance the equation.
• 2Ca O2 ? 2CaO

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Steps to Balance Chemical Equations

• Write the formula equation with the correct
  symbols and formulas.
• Na Cl2 ? NaCl
• Count the number of atoms of each element on each
  side of the arrow.
• Balance atoms by using coefficients.
• 2Na Cl2 ? 2NaCl
• Check your work by counting atoms of each element.

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Edible Equations

• 1. Gather several thin pretzel sticks and a
  package of MMs.
• 2. Use the pretzels and MMs to make models of
  the following chemical reactions
• 2KClO3 ? 2KCl 3O2
• U 3F2 ? UF6
• Cd HCl ? CdCl2 H2
• Cs2 O2 ? CO2 SO2
• 3. How do your models illustrate the Law of
  Conservation of Matter?

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Practice

• Sodium phosphate is used to cut grease. Write a
  balanced equation for the reaction in which iron
  (II) chloride reacts with sodium phosphate to
  produce sodium chloride and iron (II) phosphate.
• Chlorine reacts with lithium bromide to produce
  lithium chloride and bromine.

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Classifying Chemical Reactions

• Types
• Direction Combination Reactions (Synthesis)- two
  or more reactants come together to form a single
  product
• A B ? AB
• 2 Na (s) Cl2 (g) ? 2 NaCl (s)
• 4Fe (s) 3O2 (g) ? 2 Fe2O3 (s)

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Direct Combination (Synthesis)
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• 2. Decomposition Reactions (Analysis)
• A reaction in which a single compound is broken
  down into two or more smaller compounds or
  elements.

AB ? A B 2H2O (l) ? 2H2 (g) O2 (g)
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Decomposition Reactions (Analysis)
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• 3. Single Replacement Reaction (REDOX)- an
  uncombined element displaces an element that is
  part of a compound.

A BX ? AX B BX and AX are generally
ionic compounds and A and B are elements.   Mg
(s) CuSO4 (aq) ? MgSO4 (aq) Cu (s)   Fe
(s) CuSO4 (aq) ? FeSO4 (aq) Cu (s)
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Single Replacement Reaction (REDOX)
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Single Replacement Reactions

• A more active element will replace a less active
  element.
• Table J- Activity Series
• Substances higher in the table will replace
  substances lower in the table
• Pb(s) CuSO4 (aq) ?

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Single Replacement Reactions

• Mg (s) 2HCl (aq) ?
•  2Al (s) 3ZnCl2 (aq) ?
• Al (s) NaCl (aq) ?
• NaCl (aq) H2(g) ?

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• 4. Double Replacement Reactions-
• Atoms or ions from two different compounds
  replace each other. An identifying
  characteristic of a double replacement reaction
  is the presence of two compounds as reactants and
  two compounds as products. Switch Partners
• AX BY ? AY BX
• CaCO3 2HCl ? CaCl2 H2CO3

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Double Replacement Reactions
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• Double replacement reactions do not occur unless
• The reactants are dissolved in water so that the
  compounds can separate into ions.
• And one of the following
• 1.
• 2.
• 3.
• Table F, which shows the solubilities of various
  ionic substances in water,can be used to help us
  to determine if a precipitate is formed.

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Predict if the following Double Replacement
reactions will occur and indicate why the
reaction does or does not occur.
  AgNO3 NaCl ? AgCl NaNO3 KOH
Al(NO3)3 ? KNO3 Al(OH)3   NaOH HCl ?
NaCl H2O   KBr NaNO3 ? KNO3
NaBr