Chemical Bonds Standard 2: chapter 7, 8, 9

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Chemical Bonds Standard 2 chapter 7, 8, 9

• Vocabulary
• leave enough space for definition AND an example
• Metallic bond
• Alloy
• Ionic bond
• Cation
• Anion
• Crystal
• Covalent bond
• Polar covalent bond
• Diatomic molecule
• Electron dot structure

• Siddall.
• Chemistry.

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Standard 2d Intermolecular forces

• Gases
• Particles have no attraction to each other
• Extremely low melting point
• low boiling point
• Particles move rapidly and randomly
• no fixed shape
• no fixed volume

• Solids
• Particles are strongly attracted to each other
• High melting point
• Particles vibrate in place
• fixed shape
• fixed volume

• Liquids
• Particles are weakly attracted to each other
• Low melting point
• Particles move around each other freely
• no fixed shape
• fixed volume

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study question 1

• If a substance has no fixed shape it could be
  _________ or _________
• If a substance has no fixed shape and no fixed
  volume it would be ___________

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• Physical state The state of a material depends
  on the balance between
• the kinetic energy of the particles.
• the attractions between particles
• Kinetic energy gt attractions gas
• Kinetic energy lt attractions liquid
• Kinetic energy ltlt attractions solid

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study question 2

• If the kinetic energy of particles in a substance
  is much greater than the forces between
  particles, the substance is a __________

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Phase changes.

• Melting a solid or evaporating a liquid requires
  energy to overcome the forces holding the
  particles together.

• Freezing a liquid or condensing a gas is caused
  by removing energy so attractive forces between
  particles dominate

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Physical state
gas
evaporating
Condensing
liquid
Energy added/absorbed intermolecular attractions
are overcome
Energy released/removed intermolecular
attractions take over
melting
freezing
solid
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study question 3

• Which processes occur when energy is removed from
  a substance?

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Honors Only Volatility degree of change.

• A substance with high volatility will change
  easily from solid to liquid or liquid to gas. For
  example
• carbon dioxide, oxygen are very volatile
  compounds (gas at room temperature)
• Water is less volatile (liquid at room
  temperature)
• Iron, salt are considered non-volatile compounds
  (solids at room temperature)

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Study question 4

• List the following compounds as extremely
  volatile, somewhat volatile or non-volatile.
  Explain each choice.
• Methane (natural gas)
• Alcohol
• Calcium carbonate (rocks)

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Standard 2a Types of Bonds
Type of atoms Electrons Solid/ liquid/gas
Metallic bond Metals Shared between atoms Solid
Ionic bond Metals non-metals Transferred from one atom to another Solid
Covalent bond Non-metals Shared between atoms Gas, liquid or solid
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study question 5

• What type of bonds are formed when non-metal
  atoms share electrons?

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Bonding in Metals

• Metal atoms share valence electrons.
• Atoms are very close together
• ? Metals are solid compounds
• Electrons move around (sea of electrons)
• ? Metals conduct electricity
• ? Metals are malleable.
• ex lead

Alloy mixture of different metals with specific
properties superior to individual metals. e.x.
steel frame construction.
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study question 6

• Explain why metals are solid and why they conduct
  electricity.

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Standard 2c Ionic bonds.

• An Ionic bond is formed between metal and
  non-metal atoms
• Each atom gains or loses electrons in order to
  form an octet
• An ion is a charged particle
• A cation positive ion a Metal.
• e.x. Na, Ca2, Al3
• An anion negative ion a Non-metal
• e.x. Cl-, S2-, P3-

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study question 7

• For the compound KF
• Which atom is the cation?
• Which atom is an anion?

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Crystal Lattice Structure

• All ionic compounds form a crystal lattice
  structure
• formed by a very large network of electrostatic
  attractions (positive and negative ions attracted
  to each other).
• Lattice energy The energy needed to break the
  electrostatic attractions holding the lattice
  structure.
• Ionic compounds are always solid because of the
  strong electrostatic attractions between ions

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Crystal Lattice Structure.





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study question 8

• Why do ionic compounds form crystal lattice
  structures?

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Properties of ionic compounds

• Electrostatic attractions are very strong
  therefore ionic compounds
• are solids at room temperature.
• have very high melting points.

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study question 9

• Which of the following are solid at room
  temperature?
• CaO
• CO
• NO2
• Na2O

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naming ionic compounds
e.x. Na2O

• Use cation name
• modify anion name (ide)
• Do NOT use prefixes
• Name sodium oxide

More examples CaF2 calcium fluoride K2O
potassium oxide
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study question 10

• Name the following
• MgCl2
• Al2O3
• NaBr

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Weird things

• Polyatomic ions act as one charged particle in
  an ionic bond
• Anion names are not modified for polyatomic ions
• Example NH4OH
• Ammonium hydroxide
• Example Al(NO3)3
• aluminum nitrate

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study question 11

• Name the following
• NaOH
• K2SO4
• Mg(NO3)2

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Writing ionic formulas from names

• example
• Sodium hydroxide (Na and OH-)
• Charges must cancel out NaOH
• example
• Magnesium hydroxide (Mg2 and OH-)
• For each Mg2 there must be 2 x OH- Mg(OH)2
• note use parenthesis only when showing more
  than one polyatomic ion

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study question 12

• Write formulas for the following compounds
• Aluminum hydroxide
• Potassium oxide
• Magnesium nitrate

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Standard 2b covalent bonds

• Covalent (molecular) compounds
• Formed when non-metal atoms bond.
• Bonds between atoms are strong
• But many covalent compounds are liquids or gases
  because molecules are not strongly attracted to
  each other
• ex H2O, CO2
• Properties
• many covalent molecules have very low melting
  points and high volatility
• Many covalent molecules are gases or liquids

O
C
O
O
H
H
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study question 13

• Identify the covalent compounds
• CO2
• CaO
• MgCl2
• CCl4

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Naming covalent compounds.

• e.x. CO2 1 carbon 2 oxygen
• Name carbon dioxide
• Modify name of second atom (ide).
• Add pre-fix to indicate number of atoms.

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Prefixes

1. mono
2. di
3. tri
4. tetra
5. penta

1. hexa
2. hepta
3. octa
4. nona
5. deca

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• Examples
• CCl4
• Carbon tetrachloride.
• N2O3
• Dinitrogen trioxide.
• Exception The mono prefix is usually omitted
  from the first atom
• NO nitrogen monoxide

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Diatomic molecules.

• molecules formed from 2 atoms
• H2
• N2
• O2
• F2
• Cl2
• Br2
• I2

• hydrogen
• nitrogen
• oxygen
• fluorine
• chlorine
• bromine
• iodine

NOTE Nitrogen N2 N2 is a molecule N a
nitrogen atom
News flash you must know these
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study question 14

• Name the following
• CO
• CO2
• Cl2
• NO
• N2O

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Standard 2e Lewis Dot Diagrams

• diagrams show
• Chemical symbol
• Valence electrons

• Each atom has 4 valence electron orbitals (one s
  orbital and 3 p orbitals)
• Each orbital can hold 2 electrons.
• Electrons like to be alone.
• electrons pair up if necessary.

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Examples of Lewis Dot diagrams
e.x. Nitrogen atom
e.x. Sulfur atom



electron
S

N







Electron orbitals
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study question 15

1. Draw a Lewis Dot Diagram for an oxygen atom
2. Draw the Lewis Dot Diagram for a chlorine atom

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Creating an Octet

• Non-metal atoms form covalent bonds in order to
  share electrons and create an octet.

• Example H2
• Each hydrogen has one electron.
• Each hydrogen needs two electrons (like He).

H
H
H
H

Covalent bond 2 shared electrons (show in
between atoms)
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study question 16

• Draw the Lewis dot diagram for a chlorine
  molecule (Cl2)

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e.x. Oxygen molecule (O2).

• Oxygen atom
• needs 2 more electrons (to have an octet)
• forms 2 bonds (using 2 unpaired electrons)

O
O



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Double bond

• O

O



One bond

• The oxygen molecule still has the same total
  number of electrons
• But each atom thinks it has an octet.

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study question 17

1. Draw the Lewis dot diagram for N2

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• Rules for Dot Diagrams
• Count total number of valence electrons for all
  atoms.
• Determine number of bonds needed for each atom.
• Allocate unpaired electrons to bonds.
• Allocate unshared (paired) electrons to orbitals
  so each atom has an octet.
• Re-count total number of electrons in diagram.

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e.x. CH4 (methane molecule)

• Hydrogen
• has 1 electron
• needs 1 electron
• forms 1 bond.

• Carbon
• has 4 electrons
• needs 4 electrons
• forms 4 bonds.

H
H


C
H

H
Total number of electrons 8
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H
octet


Single bond.


H
C
H




Helium electron configuration
H

• The total number of electrons did not change.
• Each atom thinks it has an octet.

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study question 18

• Draw the correct Lewis Dot Diagram for H2O

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Danger!

• HONORS STUDENTS ONLY BEYOND THIS POINT.

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VESPR TheoryValence Shell Electron Pair Repulsion

• An atom with no unshared electrons forms four
  bonds

4 single bonds tetrahedral
2 single bonds 1 double bond trigonal planar
2 double bonds linear
1 single bond 1 triple bond tetrahedral
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Study question 19

• Draw the lewis dot diagram for CF4 and determine
  the shape of the molecule

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• An atom with 1 unshared pairs of electrons forms
  3 bonds
• The electron pair takes up the space of an
  orbital
• An atom with 2 unshared pairs of electrons forms
  2 bonds

3 single bonds trigonal pyramidal NOT trigonal
planar
2 single bonds bent NOT linear
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Study question 20

• Draw the lewis dot diagram for water and
  determine the shape of the molecule

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2g electronegativity and bonds.

• Polar Covalent Bonds Formed when atoms share
  electrons unequally.
• Electrons have a negative charge
• Atoms with high electronegativity attract the
  electrons in a bond. This creates a dipole within
  a molecule.
• Dipole charge difference
• Polar molecules (molecules with a dipole) are
  attracted to each other like magnets.
• Molecules that are extremely polar are usually
  liquids because the molecules are close together
• Non-polar molecules are usually gases

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• e.x. Hydrogen fluoride

Slightly negative

d
d-

F

H

Slightly positive
Arrow points to negative part of molecule tail
shows for positive
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study question 21

• Draw the Lewis dot diagram for HCl and use
  symbols to indicate the dipole

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• Calculating polar or non-polar bonds.
• Non-polar bond difference in electronegativity lt
  0.5
• Polar bond difference in en 0.5
• Ionic bond difference in en gt 2
• Example H-F
• 4.0 2.1 1.9 polar bond
• Example H-C
• 2.5 2.1 0.4 non-polar bond

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study question 22

• Label the following molecules polar or non-polar
• CO
• CN
• O2

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• Hydrogen bonding (this is NOT a bond)
• A special type of intermolecular attraction
• Hydrogen bond intermolecular attraction between
  a hydrogen nucleus in one molecule and a very
  electronegative atom in another molecule.
• very electronegative atoms O, N, F, Cl

• Hydrogen bonding results in low volatility
• For water there are some unique properties
• Lower density of solid water
• High boiling point

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study question 23

• Which of the following molecules might hydrogen
  bond? CH4, CH2O, HF, H2O, H2S

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Van der Waals forces

• Intermolecular forces
• Exist between all types of atoms/molecules
• Electrons orbiting atoms become unevenly
  dispersed creating a temporary dipole
• Also called London Dispersion Forces

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Study question 24

• Which atoms would probably create stronger Van
  der Waals forces? Small atoms or large atoms.
  Why?
• Would Van der Waals forces make a liquid more
  volatile or less volatile?