Intramolecular and Intermolecular Forces

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Intramolecular and Intermolecular Forces

• Week 4 Tuesday

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Electronegativity

• A measure of an atoms ability to attract the pair
  of electrons that it shares with another atom
  with in a covalent bond.

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Electronegativity

• The atomic radius plays a big role in how
  electronegative an atom will be. The larger an
  atom is the weaker the attraction for shared
  electron pairs will be. This is due to electrons
  shielding each other and thus making the
  attraction of electron pairs less.

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Periodic trend

• As the number of shells increase
  electronegativity decreases (move down the
  periodic table)
• As more protons are added atomic radius decreases
  and thus increases electronegativity. (move left
  to right)

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Periodic trend
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Polar and Nonpolar Covalent Bonds (p41)

• Intramolecular Force - the attractive force
  between atoms and ions in a compound
• If the electron pair is shared in an equal manner
  (ie. Each atom has an equal attraction to the
  shared electron pair), then the bond is a
  nonpolar covalent bond.
• Ex H2, O2, Cl2.

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Polar and Nonpolar Covalent Bonds (p41)

• When electron pairs are not shared equally
  between atoms they will spend more time near one
  atom than another.
• This would mean that one of the atoms is slightly
  more negative than the other atom. To show these
  partial localized charges, we use the symbol d
  for a localized positive charge and d- for a
  localized negative charge. These bonds are called
  polar covalent bonds.
• Example HCl.

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• Whether or not a bond is polar covalent depends
  on the difference between the electronegativities
  of the bonded atoms.
• The larger the difference in electronegativities
  the more polar the bond is.
• A bond is considered ionic when the difference is
  1.7. Some ionic bonds will have less than this.

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Polar and Non-polar molecules (p42)

• Polar molecules - are molecules that have a
  positive charged end and a negative charged end.
• Nonpolar molecules - are molecules that do not
  have charged ends.

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Polar and nonpolar molecules

• The polarity of a molecule depends on two
  characteristics of the molecule.
• The presence of polar covalent bonds
• The three dimensional shape (geometry) of the
  molecule

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• Water is a polar molecule because it has polar
  covalent bonds and has a shape that allows it to
  have a positive end and a negative end. Methane
  on the other hand has polar covalent bonds but
  since the shape is pyramidal there is no real
  positive end or negative end.

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Intermolecular Forces

• Bonds between molecules forces of attraction
  that form between a molecule and its neighbouring
  molecules.

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There are 3 types

• 1. Van der waals forces (contains two of them)
• Dipole-dipole force- an attractive force between
  polar molecules. (polar molecules)
• London dispersion- an attractive force between
  all molecules, including non-polar molecules.
  (non-polar and polar molecules) Weaker than
  dipole-dipole

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• 2. Hydrogen bonding Similar to dipole-dipole
  but stronger. It occurs in highly polar molecules
  have either a H-N, H-O or H-F bond. This is why
  it requires large amounts of energy to break
  water molecules apart.(solid, liquid)
• -When water freezes it forms a lattice structure
  which makes it less dense than water.

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Homework

• Try this (p 44) if in class with model kit
• Page 45 1-6

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Molecular Nomenclature
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Molecular Nomenclature

• The names of molecular compounds often contain
  prefixes.
• The prefixes are used to count the number of
  atoms in the molecule.
• This is important as molecular compounds
  containing 2 elements can have different
  combinations which have different properties.

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Binary Molecular Compounds IUPAC

• 1. Write down the name of the first element. If
  there is more than one atom of this element
  attach a Greek prefix. (if there is only one atom
  do not attach the prefix)
• 2. Attach a Greek prefix (relating to the number
  of atoms) to the second elements name and add
  -ide.
• Example
• CO Carbon monoxide
• CO2 Carbon dioxide

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Prefixes used when naming binary molecular
compounds
of atoms 1 2 3 4 5 6 7 8 9 10
Prefix mono di tri tetra penta hexa hepta octa nona deca
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Practice
HF PBr CS2 NO H2S NO2 SCl2 N2O2 sulphur trioxide dihydrogen monoxide carbon tetrafluoride Silicon dibromide Trihydrogen mononitride Disulfur trioxide Carbon diphosphide Chlorine gas
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Word Equations
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Word Equations

• A word equation is a way of representing a
  chemical reaction it tells you what reacts and
  what is produced. Word equations are an efficient
  way to describe chemical changes, to help
  chemists recognise patterns, and to predict the
  products of a chemical reaction.

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Word Equations

• They are written in a particular order.
• Reactants are always on the left side of the
  arrow and products are always on the right side
  of the arrow.
• Multiple reactants or products are separated by a
  sign.

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Examples

• Silver nitrate copper ? silver copper(II)
  nitrate
• Hydrogen Oxygen ? water vapour
• Which are the products and which are reactants

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The Conservation of Mass
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The Conservation of Mass

• In a chemical reaction, the total mass of the
  reactants is always equal to the total mass of
  the products.
• This tells us a few things.

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Conservation of mass

• Atoms do not change in a reaction. The molecules
  that they form can be changed but the atoms
  themselves are not.
• Mass cannot be destroyed. If it could we could
  use E MC2

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Example

• Methane oxygen ? water carbon dioxide

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Skeleton equations

• A representation of a chemical reaction where the
  formulas of the reactants are connected to the
  formulas products by an arrow.
• CH4 O2 ? H2O CO2
• This however does not demonstrate the Law of
  Conservation of Mass states that the mass of the
  products will equal the mass of the reactants.

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Combustion of methane
Type of atom Reactants Products
C 1 1
H 4 2
O 2 2 1 3
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• We cant change the formulas of the products or
  reactants so the only thing we can do is change
  the number of molecules instead of their
  formulas.

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• CH4 O2 O2 ? H2O H2O CO2
• CH4 2O2 ? 2H2O CO2
• Now the chemical equation is balanced and the
  mass of the reactants will equal the mass of the
  products

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Type of atom Reactants Products
C 1 1
H 4 4
O 4 2 1 1 4
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Steps to balancing an equation

• Step 1
• Write the word equation of the reaction
• Aluminum bromine ? aluminum bromide
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• Step 2
• Write the skeleton equation by replacing each
  name with a correct formula.
• Al Br2 ? AlBr3

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• Step 3

Type of atom Reactants Products
Al 1 1
Br 2 3
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• Step 4
• Multiply each of the formulas by the appropriate
  coefficients to balance the number of atoms.
• Start out by picking the element with the most
  number of atoms and try to balance it first. We
  will start with Bromine. The 2 and 3 will be
  balanced if we multiply the reactant side by 3
  which would give it 6 Br, and multiply the
  product side by 2 to give us 6 Br. Now we have 2
  Al products which need to be balanced so we add a
  2 to the Al on the reactant side.
• 2Al 3Br2 ? 2AlBr3

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More examples

• Hydrogen gas Chlorine gas ? hydrogen chloride
• Sodium chlorine ? sodium chloride
• Nitrogen hydrogen ? ammonia (hydrogen nitride)
• N2 H2 ? NH3 N2 H2 ? NH3

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Types of chemical reactions (P48-52)
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Types of chemical reactions

• 1. Combustion Reaction
• The reaction of a substance with oxygen,
  producing oxides and energy
• Fuel oxygen ? oxides energy
• AB oxygen ? common oxides of A and B (ex AO,
  BO)

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Common oxides include
C CO2(g) H H2O(g) S SO2(g) N NO2(g)
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Example

• C2H5OH(g) 3O2(g) ? 2CO2(g) 3H2O(g)

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• Incomplete combustion occurs when there is not
  sufficient oxygen. When this occurs 4 products
  are produced instead of the usual H2O and CO2.
  Incomplete combustion also produces CO(g) and
  C(s). This is commonly seen when lighting an
  acetylene torch.
• 3C2H2(g) 3O2(g) ? 2CO2(g) 3H2O(g)
  2CO(g) 2C(s)

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2. Synthesis Reaction

• -A chemical reaction in which two or more
  substances combine to form a more complex
  substance.
• A B ? AB
• Example
• 2CO(g) O2(g) pt? 2CO2(g)

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3. Decomposition reaction

• -A chemical reaction in which a compound is
  broken down into two or more simpler substances.
• AB ? A B
• Example The decomposition of water.
• 2H2O(l) electricity ?2 H2(g) O2(g)

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Chemical reactions in solution

• A Solution is a homogenous mixture in which a
  pure substance, called the solute, is dissolved
  in another pure substance called the solvent.
•  
• The solution is often an aqueous solution which
  is a solution where water is the solvent.

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4. Single Displacement reaction

• -A reaction of an element with a compound to
  produce a new element and a new compound. The
  reaction will only occur if the element is higher
  on the reactivity series than the metal in the
  compound. (Reactivity series page 500)
• A BC ? AC B
•  
• Example
• Cu(s) 2AgNO3(aq) ? 2Ag(s) Cu(NO3)2(aq)

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5. Double displacement reaction

• - A reaction in which aqueous ionic compounds
  rearrange cations and anions, resulting in the
  formation of new compounds. If a new product does
  not happen then the reaction has not happened.
  (Solubility table page 54)
• AB CD ? AD CB
• Example
• CaCl2(aq) Na2CO3(aq) ? CaCO3(s) 2NaCl(aq)

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Double displacement is likely to occur if

• - a precipitate is produced
• - a gas is produced
• - a acid base neutralization occurs

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Examples

• 2As 3O2 ? 2As2O3
• Type of reaction?
• KClO4 ? KCl 2O2
• Type of reaction?
• Zn(s) CuCl2(aq) ? Will it occur? Look at
  the table.

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• Cu(s) ZnCl2(s) ?
• will it occur
• MgCl2(aq) K(OH)2(aq) ? will it occur?
• Name the precipitate

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Work sheet

• This will be handed in for marks.
• Do it, it is great practice.

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Review

• If you would like the review now, I can give you
  the paper copy with the practice questions.