Bonds, Reactions

1
Bonds, Reactions Amounts

• The noble gases are inert, this suggests that
  their electronic structures are stable. When
  other atoms acquire the same number of electrons
  in their outermost or valence shell they will
  become more stable.
• The octet rule
• Atoms gain, lose or share electrons to
  acquire the configuration of the nearest noble
  gas.

2
Ionic Bonding

• Formed when one atom loses one or more electrons
  and another gains one or more electrons. Loss of
  electrons from metals forms cations and gain of
  electrons by nonmetals forms anions. The ions are
  held together by a strong electrostatic
  attraction. This results in a lower potential
  energy.
• Metals form cations because it is easy for them
  to lose electrons but hard to gain them.
• Nonmetals form anions because it is difficult for
  them to lose electrons but easy to gain them.

3
Cations Anions

• Na 2e-8e-1e- ? Na 2e-8e- Same as Ne
• O 2e-6e- ? O2- 2e-8e- Same as Ne
• Ca 2e-8e-8e-2e- ? Ca2 2e-8e-8e- Same as
  Ar
• Cl 2e-8e-7e- ? Cl- 2e-8e-8e- Same as Ar
• Ions or atoms with the same configuration are
  said to be isoelectronic.
• Li 2e-1e- ? Li 2e- Same as He

4
Lewis Symbols Ions
Lewis symbols can be used to represent ionic
bonding only between representative non-metals
and metals. Instead of using complete electron
configurations to represent the loss and gain of
electrons, Lewis symbols can be used.

Na Cl ? Na1 Cl 1-


5
Formulas of Ionic Compounds

• Compounds are always neutral, so the proportion
  of cations to anions must reflect this in the
  formula.
• One Na combines with one Cl- to form neutral
  NaCl.
• If both ions have same charge use one of each,
  e.g. Ca2 O2- gives CaO.

6
Formulas III

• To find the charge on ions of representative
  elements-
• Metals charge is the same as the group number,
  e.g. Al is in group 3A, so the ion is Al3.
• Nonmetals charge is (group - 8), e.g. S is in
  group 6A, so charge is 6 - 8 -2 and the ion is
  S2-.
• Formula obtained from crossover rule
• Al3 S2-
• Al2S3

7
Properties of Ionic Compounds

• Ionic compounds are crystalline solids, with
  alternating cations and anions making up the
  crystal structure.
• The high electrostatic attraction causes them to
  be hard, with high melting points.
• They are not molecules and the smallest
  collection of ions with the correct proportions
  is called a formula unit.

8
Ion Names

• Metal ions are given the name of the metal except
  when more than one charge can exist.
• Nonmetals typically shorten the name and have an
  -ide ending - carbide, nitride, phosphide, oxide
  sulfide halogens replace the n with a d.
• In compounds the metal is placed before the anion
  name.

9
Multiple Charges 1

• Some metals form ions with more than one charge.
• These are named by the Stock system that shows
  the charge on the ion by a roman numeral after
  the name of the metal.
• Some also have names using the latin form of the
  name followed by the ending ic for the higher of
  two charges. The ending ous is used for the
  lower charge.

10
Multiple Charges 2

• Fe2 iron II or ferrous
• Fe3 iron III or ferric
• Cu copper I or cuprous
• Cu2 copper II or cupric
• Pb2 lead II or plumbous
• Pb4 lead IV or plumbic
• Sn2 tin II or stannous
• Sn4 tin IV or stannic

11
Polyatomic Ions

• Many atoms form groups that are like molecules
  except that they are charged, the majority
  involve a non-metal and varying numbers of oxygen
  atoms. You will definitely need to learn these
• H3O - hydronium NH4 - ammonium
• OH- - hydroxide SO42- - sulfate
• NO3- - nitrate CO32- - carbonate
• HCO3- - hydrogen carbonate or bicarbonate
• PO43- - phosphate.

12
Formulas with Polyatomic Ions

• In formulas one ion needs no subscript, but with
  more than one they are placed within parentheses
  and the subscript placed after that. Examples
• aluminum sulfate - Al2(SO4)3
• sodium sulfate - Na2SO4
• ammonium sulfate - (NH4)2SO4
• ammonium bicarbonate NH4HCO3

13
Covalent Bonds

• Because nonmetals cannot form cations, they form
  bonds by sharing one or more pairs of electrons
  this constitutes a covalent bond.
• Combinations of two or more atoms connected to
  each other by covalent bonds are called
  molecules.
• Polyatomic ions are constructed the same way but
  carry a charge.

14
Hydrogen

• Hydrogen When two atoms approach each other, as
  they get close both nuclei attract them so
  electrons move to the region between the nuclei.
  The nuclei also repel each other so the lowest
  energy level is when the attractions and
  repulsions balance. Can represent bonds in two
  ways

15
Multiple Bonds

• If one shared pair is not enough to give each
  atom an octet then more pairs must be used.
• O2 CO2 have double bonds and N2 has a triple
  bond.

16
Naming Covalent Compounds

• Other than established common names (e.g. water
  and ammonia), the least electronegative element
  is named first. Prefixes is placed before the
  element show how many of each type of atom is
  present
• mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa-
  6 hepta- 7 octa- 8 nona- 9 deca- 10
• Mono is frequently omitted and the a is left out
  when preceding a vowel. Examples
• NO - nitrogen monoxide N2O4 - dinitrogen
  tetroxide PCl5 - phosphorus pentachloride.

17
Bond Polarity

• When different atoms combine one of the pair will
  have a stronger attraction for electrons than the
  other one as a result this atom acquires a
  slight negative charge and the other an equal
  positive charge and this is a polar bond.

18
Polarity II

• The pull an atom has on electron pairs in a bond
  is called electronegativity. Values increase
  from left to right across a period and from
  bottom to top within a group noble gases have a
  0 value so F is the most electronegative.
• When the electronegativity difference is 1.8
  bonds are ionic.
• With a difference gt0 but lt 1.8 polar covalent
• With 0 difference nonpolar covalent.

19
Lewis Structures

• 1. Decide which atoms are bonded.
• 2. Count all valence electrons.
• 3. Add one electron for each unit of negative
  charge and subtract one for each unit of
  positive charge.
• 4. Place two electrons in each bond.
• 5. Complete the octets of the outer atoms by
  adding electron pairs.
• 6 Place any extra electrons on the central atom
  in pairs.
• 7. If the central atom still has not got an
  octet, make it do so by forming double or triple
  bonds.
• Pairs of electrons that are not involved in bonds
  are called lone pairs.

20
Carbonate Ion
Carbonate ion Formula CO32- Valence electrons
4 (3x6) 22 Charge 2- Total electrons 22 2
24 Using rules 1 5 above leads to
This uses all available electrons, compensate
according to rule 7 by forming one double
bond
21
Sulfur Tetrafluoride
Formula SF4. Total electrons 6 (4 x 7)
34 Rules 1-5 lead to
Short by 2 electrons, apply rule 6.
22
Molecular Shapes

• With three or more atoms the shape can be
  described. Two or more orbitals can combine to
  give hybrid orbitals with the same energy level,
  midway between that of the original ones.

23
VSEPR Notation
Table 10.1 in the text summarizes various
possibilities for molecular geometries in
relation to electron-group geometries. In the
VSEPR notation used to describe molecular
geometries, the central atom in a structure is
denoted as A, terminal atoms as X, and the lone
pairs of electrons as E. For structures with no
lone-pair electrons (AXn), the molecular geometry
and electron-group geometry are identical.
24
Polarity of Molecules

• Molecules that have the same outer atom and are
  AX2, AX3 or AX4 will be non-polar, no matter how
  polar the individual bonds.
• Those that have lone pairs such as AX2E, AX3E or
  AX2E2 will be polar.

25
Physical vs. Chemical Changes

• When a physical change takes place there is no
  change in composition, as in boiling water. But
  if an electrical current is passed through water
  it will form oxygen and hydrogen. This change in
  composition is what distinguishes a chemical
  change from physical one. Typical signs of a
  chemical change are change in color, production
  of a gas, heat or formation of a solid in a
  solution there are more.

26
Law of Conservation of Matter

• In a chemical reaction matter is neither created
  nor destroyed.
• This means that whatever elements are present
  before a reaction takes place must still be
  present after the reaction is over, though in
  different combinations.

27
Chemical Equations

• We show this by writing chemical equations. For
  example when water is formed from the reaction of
  hydrogen with oxygen there must be as many atoms
  of each kind present before and after the
  reaction. To show this we use an equation as
  below. The initial substances are called
  reactants and are written on the left, the
  materials formed are called products and are on
  the write.
• 2H2 O2 ? 2H2O

28
Equations 2

• The formulas cannot be changed so the numbers in
  front of each, called coefficients show how many
  molecules, formula units or atoms are needed to
  balance the equation to agree with the
  conservation law. In doing so polyatomic ions are
  kept as intact units, unless they undergo a
  change themselves.
• 3CaCl2(aq) 2Na3PO4(aq) ? Ca3(PO4)2s) 6NaCl(aq)

29
Types of Reaction

• In the following equations the following
  subscripts are used to show what physical state a
  substance has
• (g) gas (l) liquid (s) solid (aq)
  aqueous, i.e. a solution in water.
• Combination
• N2(g) 3H2(g) ? 2NH3(g)
• Decomposition
• 2NaHCO3(s) ? Na2CO3(s) H2O(g) CO2(g)
• Single replacement
• Zn (s) H2SO 4(aq) ? H 2(g) ZnSO4(aq)
• Combustion
• C3H 8(g) 5O2(g) ? 3CO2(g) 4H2O(g)
• Double replacement
• AgNO3(aq) NaCl(aq) ? AgCl(s) NaCl(aq)

30
Oxidation Reduction

• Oxidation
• 1. Gain of oxygen 4Fe 3O2 ? 2Fe2O3
• The iron has been oxidized to iron III oxide
  (rust).
• 2. Loss of hydrogen
• CH3CH2OH O CH3CHO H2O
• The ethyl alcohol has been oxidized to
  acetaldehyde.
• 3. Loss of electrons
• Fe 2HCl ? FeCl2 H2
• The iron has been oxidized to the Fe2 ion,
  losing two electrons (the same process takes
  place forming Fe3)

31
Reduction

• 1. Loss of oxygen FeO CO ? Fe CO2
• Here the iron II oxide loses oxygen to form iron.
• 2. Gain of hydrogen
• CH3COCH3 2H ? CH3CH(OH)CH3
• Here acetone is reduced to isopropyl alcohol.
• 3. Gain of electrons Mg Cl2 ? MgCl2
• Here the chlorine atoms are reduced to chloride
  ions.

32
Oxidizers Reducers 1

• 2Al 3S ? Al2S3
• Oxidation never takes place without reduction.
  Here the aluminum is oxidized to the aluminum ion
  - so sulfur is the oxidizing agent and is itself
  reduced to form the sulfide ion.
• Similarly the aluminum is the reducing agent when
  it becomes oxidized.

33
Oxidizers Reducers 2

• Oxidizing agents - Oxygen 21 of air, oxidizes
  metals and nonmetals to oxides, and hydrocarbon
  fuels to CO2 and H2O. Halogens. H2O2. Various
  ions - ClO- MnO4- Cr2O72-.
• Reducing agents - Hydrogen not found free,
  secondary fuel when burned with oxygen also
  reduces metal oxides to metals. Some metals and
  carbon reduce other metal ores to metals.

34
Biological Oxidation Reduction
Energy is obtained from carbohydrates- C6H12O6
6O2 ? 6CO2 6H2O Each carbon has on
average lost 2 hydrogens and gained 1 oxygen, so
oxidation has occurred. The reaction is reversed
in photosynthesis so this is reduction.
35
The Mole

• N2 3H2 ? 2NH3
• 1 molecule N2 3 molecules H2 ? 2
  molecules NH3
• 10 molecules N2 30 molecules H2 ? 20
  molecules NH3
• 1 x 106 mlcls N2 3 x 106 mlcls H2 ? 2
  x 106 mlcls N2
• 6.02 x 1023 mlcls N2 18.06 x 1023 mlcls H2 ?
  12.04 x 1023 molecules NH3
• 6.02 x 1023 particles 1 mole of particles
• (abbreviation "mol")
• 1 mol N2 3 mol H2 ? 2 mol NH3

36
Mole Examples

• 1 mol Atomic Mass in grams
• 1 mol Molecular Mass in grams
• 1 mol Formula Mass in grams (for ionic
  compounds)
• Example
• 1 mol O 6.02 x 1023 atoms O 16.00 grams O
• 1 mol N2 28.02 g N2
• 1 mol H2 2.016 g H2 1 mol NH3 17.03 g NH3

37
Mole Equation

• 1 mol N2 3 mol H2 2 mol NH3
• 1 x 28.02 g N2 3 x 2.016 g H2 2 x 17.03 g
  NH3
• 28.02 g N2 6.048 g H2 34.07 g NH3
• Note Demonstrates Law of Conservation of Mass

38
Moles and Grams

• 1. How many moles in 25.0 g of aluminum sulfate?
• Find molar mass of aluminum sulfate, Al2(SO4)3
• Al 26.98 x 2 53.96
• S 32.07 x 3 96.21
• O 16.00 x 12 192.00
• 342.17
• Therefore 1 mol Al2(SO4)3 342.17 g Al2(SO4)3
• 25.0 g Al2(SO4)3 x 1 mol Al2(SO4)3
  0.0731 g Al2(SO4)3
• 342.17 g Al2(SO4)3

39
Moles and Grams II

• 2. How many grams in 0.175 mol of sodium
  chloride?
• Find molar mass for sodium chloride, NaCl
• Na 22.99 x 1 22.99
• Cl 35.45 x 1 35.45
• 58.44
• Therefore 1 mol NaCl 58.44 g NaCl
• 0.175 mol NaCl x 58.44 g NaCl 10.2 g NaCl
• 1 mol NaCl

40
Stoichiometry

• Consider the combustion of 35.0 g of ethane
  according to the following reaction.
• 2C2H6 7O2 ? 4CO2 6H2O
• 1. How many moles of water are produced?
• 2. How many grams of oxygen are needed?
• 3. How many grams of carbon dioxide are produced?

41
Problem 1

• The following relationships exist and are needed
  to answer the first problem?
• 1 mol C2H6 30.07 g C2H6
• 1 mol C2H6 3 mol H2O
• 35.0 g C2H6 x 1 mol C2H6 x 3 mol H2O 3.49
  mol H2O
• 30.07 g C2H6 1 mol C2H6

42
Problem 2

• For the second problem we also need the
  following
• 2 mol C2H6 7 mol O2 1 mol O2 32.00 g O2
• 35.0 g C2H6 x 1 mol C2H6 x 7 mol O2 x
  32.00 g O2 130. g
• 30.07 g C2H6 2 mol C2H6 1
  mol O2

43
Problem 3

• For the third problem we also need
• 1 mol C2H6 2 mol CO2 and 1 mol CO2 44.01 g
  CO2
• 35.0 g C2H6 x 1 mol C2H6 x 2 mol CO2
  x 44.01 g CO2 103 g
• 30.07 g C2H6 1 mol
  C2H6 1 mol CO2

44
Reaction Rates
1. Collision. 2. Orientation. 3.
Temperature. 4. Concentration. 5. Catalysis.
45
Chemical Equilibrium
Many reactions can go in the reverse direction
(back reaction) and reform the reactants.
Reactant concentration decreases during a
reaction, therefore the forward rate
decreases. Products concentration increases
during a reaction, therefore the back rate
increases. When the rates of the forward and back
reactions are the same the system is in dynamic
equilibrium. For any one reaction the proportion
of reactants to products at a certain temperature
will always be the same.
46
Le Châteliers 1
If a stress is put on a system in equilibrium it
will respond to minimize the stress to maintain
the balance. Example N2 3H2 ? 2NH3
Heat Add N2 and/or H2 - produces more product, a
shift to the right. Remove N2 and/or H2 - removes
product and forms more reactants - shift to the
left. Add NH3 - shifts to the left. Remove NH3 -
shifts to the right. Raise temperature (add heat)
- shifts to left. Lower temperature - shifts to
right.