AP Chemistry Chapter 6 Notes

AP Chemistry Chapter 6 Notes (Student’s edition)

Chapter 6 problem set: 12, 15, 17, 21, 27, 34, 47, 48, 57, 58, 62, 64, 65, 67-71, 74, 76, 77, 86, 88

6.1 More about the Periodic Table: Origin of the periodic table:

A German scientist called Johann Dobereiner put forward his law of triads in 1817. Each of Dobereiner's triads was a group of three elements. The appearance and reactions of the elements in a triad were similar to each other. At this time, scientists had begun to find out the relative atomic masses of the elements. Dobereiner discovered that the relative atomic mass of the middle element in each triad was close to the average of the relative atomic masses of the other two elements. This gave other scientists a clue that relative atomic masses were important when arranging the elements.

John Newlands, an English chemist, wrote a paper in 1863 which classified the 56 established elements into 11 groups based on similar physical properties, noting that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. In 1864 Newlands published his version of the periodic table and proposed the Law of Octaves (by analogy with the seven intervals of the musical scale). This law stated that any given element will exhibit analogous behavior to the eighth element following it in the table.

So, Johann Dobereiner and John Newlands both noted atomic masses are related to chemical properties.

(1864) Lothar Meyer published a periodic table which described the placement of 28 elements. Meyer's periodic table ordered the elements into groups arranged in order of their atomic weights. His periodic table arranged the elements into 6 families according to their valence, which was the first attempt to classify the elements according to this property.

Dimitri Mendeleev , a Russian scientist, publish first real periodic table in 1869. He based on the table on the chemical and physical properties. He listed the elements in order of increasing atomic mass mass. He left spaces for undiscovered elements. Mendeleev formulated the original. Periodic Law -

In 1911, Mosely (English) discovers the proton so.... new Periodic Law - ______

Reading the periodic table:

Periods: left to right on the periodic table elements in the same period have the same # of energy levels elements in the same period do not have similar properties also known as rows, shells, and energy levels

Groups: up and down on the periodic table elements in the same group have the same # of outer shell electrons elements in the same group do have similar properties also known as columns, families, and valence

Why do elements in the same column have similar properties?

Periods of elements Row #1  2 elements  2 electrons in shell #1 Row #2  8 elements  8 electrons in shell #2 Row #3  18 elements  18 electrons in shell #3 Row #4  32 elements  32 electrons in shell #4

Groups of elements and their Properties:

Group 1 - Alkali Metals

“Alkali” comes from Arabic - means “ashes” - early chemists separated sodium and potassium compounds from ashes - the hydroxides of these compounds are strongly basic. These compounds are not found alone in nature - why? explosive with water - they are stored under kerosene - very reactive. They react with nonmetals to form salts. Many of the compounds they form are white in color. They are silvery, shiny (luster), have a low melting point, conduct electricity, and are soft (so soft, you can cut them with a knife). They are malleable (able to flattened into a sheet) and ductile (able to be drawn into a wire). Sodium and Potassium are particularly important in body chemistry.

Group 2 - Alkaline Earth Metals

“Earth” - chemists term for oxides of these elements - it was originally thought that the oxides of these elements were actually the elements themselves. Tend to form white colored compounds. Strongly basic - 2nd most reactive elements. Also not “lone state” elements. Harder, denser than group 1. Common in sea salts.

Transition Metals

Groups 3-12. Harder, more brittle, higher melting point than groups 1 and 2. Form colored compounds. Conduct heat and electricity well and are shiny. Pd, Pt, Au - very unreactive (Noble metals).

Metalloids

B, Si, As, Te, At, Ge, Sb. Stairs and 2 people under the stairs. Properties of metals and nonmetals. Brittle - used in semiconductors, computers.

Halogens

Group 17. Most reactive of the nonmetals. Not found free in nature. “Halogen” - Greek for salt former. Solids, liquids, and gases in this group. Widespread – found in sea salts, minerals, living tissue. Many applications - bleach, photography, plastics, insecticides.

Noble Gases

Group 18 . Used to be called inert - not so since Kr, Xe, Rn made compounds. Used to be called rare (He and Ar fairly abundant). Least reactive elements. Used in air conditioners, double pane windows, lights, balloons.

Lanthanides

f block. Also known as the rare earth elements - not really rare. They are shiny, silver, and reactive. Used to make TV’s glow and in creating metal alloys

Actinides

f block. Unstable and radioactive. All but 4 are artificially created. Uranium used as nuclear fuel and for coloring glass and ceramics (fiesta ware). Also have found use in deep sea diving suits and smoke alarms. f block elements are called inner transition elements - they were put into their current position by Glenn Seaborg - the only living person ever to have an element named after himself.

Note…some key points… Be able to label group/family, period, Groups IA – VIIIA and B, label 1-18 Be able to explain the relationship of electron configuration to the table Representative elements – “A” columns – last (differentiating) electron is in s or p subshell of the outer shell. Properties easy to predict. Transition Elements – “B” columns – Last electron goes into an inner d subshell. Properties are similar across this section of the periodic table. Inner transition elements – Last electron goes into f subshell

Periodicity in properties:

Coulombic Attraction - properties are related to the attraction of a positive nucleus for negative electrons - depends on .

2 properties that depend on coulombic attraction - ionization energy and electronegativity.

2 properties that are based more on # of electrons - atomic radius and ionic radius.

6.2 Atomic Radii:

basic idea is “how big an atom is” - atoms are not spheres with outer boundaries due to the wave mechanical model. Atomic radius is measured in nanometers or Angstroms.

covalent atomic radius - distance from the nucleus to the outer shell when it’s

involved in a covalent bond. + +

Van der Waals radius - half the distance between two atoms when they aren’t bonded together.

atomic radius of metals - half the distance between the nuclei of two metal atoms.

+ + atomic radius is measured in nanometers.

Van der Waals radius is generally greater than covalent.

Metallic radius is also generally greater than covalent. Why? - covalent bonds hold atoms tighter together.

Predictions (2 trends):

↓ p-table = size - natural, logical - add more

 p-table = size - not logical! why?

From left to right -

trend looks like … Page 6

With transition metals, electrons are added to inner shell so radius increases slightly moving across the table. There are “blips” due to other complicating factors (electron repulsion, etc.)

6-6 Electronegativity:

Electronegativity- basic idea - the ability of an atom to attract electrons (Linus Pauling)

Decreased distance from the nucleus = increase pull on an adjacent atom’s electron.

Electronegativity is related to atomic size:

↓ size = electronegativity

trend looks like..... Page 7

Real definition of electronegativity - the ability of an atom to attract electrons that are shared with another atom in a covalent bond.

Electronegativity values are based on Pauling’s work with bond energies. 6.3 Ionization Energy:

ionization energy - energy required to remove the most loosely held electron from the outer energy level of an atom in its gas phase.

+ - A(g) + energy  A (g) + e

As the distance between the protons and the outer shell electrons decreases, the protons’ hold on the electrons increases. Increase hold on the electrons means that more energy is required to remove the electrons.

trend looks like … Page 9

IE is related to atomic radius - 2 reasons why smaller going down the table 1. greater distance from the nucleus - attraction 2. kernel electrons “ ” outer electrons from the nucleus

There is also a 2nd and 3rd IE - always higher than the first.

IE of elements greatly increases when the outer shell has been emptied.

Which has a higher 2nd IE - Na or Mg?

Which has a higher 3rd IE - Al or Mg?

Position of Electrons:

IE and Electronegativity are related, but different.

IE involves the attraction of a nucleus for an electron. IE can be measured. Thus, increase attraction equals increase energy required to take the electron.

Electronegativity is not a measurement of energy - it can’t be directly studied. It is determined mathematically by equations based on bond energy values (Pauling)

IE is related to atomic radius - 2 reasons why smaller going down the table 1. ______2. ______Students should also recognize anomalies in the graph. There are downward “blips” at column IIIA and VIA. These are caused by natures tendency to find stability in empty and full shells as well as when 1 electron is in each orbit.

Valence Electrons:

Valence electrons - outer shell electrons involved with bonding

Column # =

1 8 2 3 4 5 6 7 Usually 2

6.4 Electron Affinity:

Electron Affinity - energy change when an electron is acquired by a neutral atom gaseous atom. (we say absorbed since for most atoms, this is an endothermic process – certainly not favored by nature)

A + e -  A- + energy (exothermic, negative delta H, high EA) small atom

some atoms must be forced to accept an electron:

A + e- + energy  A- (endothermic, positive delta H, low EA) large atom

Basic idea - some atoms want to take on electrons - they have a high electron affinity value - they release a lot of energy when accepting electrons

examples : F = -322 kJ/mole Na = -53 kJ/mole

F has a higher electron affinity = higher, more negative value

general trend looks like … Page 11

Explanation......

Column # s p Energy

1 A little negative release of a little energy wants to fill the s sublevel or lose it?

2 Positive no release – only energy absorption wants to lose electrons

3 Positive wants to lose electrons

4 A little negative Wants to have an electron in

the pz

5 Not as negative (as 6 and 7)

6 Negative wants to finish filling the p sublevel

7 Highly negative almost has a full p sublevel 8 Positive wants to maintain, so must give a lot of energy to absorb an electron

Also - 2nd EA values are always positive

Groups 6 and 5 become negative ions after the 1st EA. So, trying to place a 2nd electron into a negative ion encounters significant repulsion; therefore, +EA.

An example is Fluorine - fluorine releases energy when taking an electron, but it only wants one electron. If it takes 2 electrons it has more electrons than it needs and becomes like sodium. The 2nd EA is positive as you need to force it to take a 2nd electron. F-1 has a negative charge and thus repels the new electron.

Activity:

For metals - larger atoms are more active - why? -

For nonmetals - smaller more active - why? -

metal activity trend nonmetal activity trend

Most active metals + most active nonmetals = most stable compounds

ex: RbF - very stable LiBr - less stable

Metallic Character:

Metallic character - some metals are said to be more metallic than others - really it is just a statement about their activity. If they are more active, they are said to be more metallic.

trend looks like this... 6.5 Ionic Radii:

Ions are created by electrons. Cation - Anion - Metals tend to become Nonmetals tend to become

Cations are smaller than the neutral atom - why? –

Anions are larger - why? -

trend looks like this.... Na Na+1  examples:

Li or Li+1 -

O or O-2 -

Li+1 or Be+2 -

O-2 or N -3 -

Page 15 Isoelectronic Species: Kinds of atoms that have the same

examples - Ne 1s2 2s2 2p6

so is F -1 , O -2 , N -3 , and Na +1 , Mg +2 , Al +3

All of the atoms above are considered to be isoelectronic.

In general - an isoelectronic series decreases in radius as atomic number increases.

6-7 Hydrogen and the Hydrides:

1. Hydrogen + column I, II metals → Ionic hydrides (note: high temperature and pressure)

H2 (g) + Li(l)  LiH(s)

H2 (g) + Ca(l) 

2. Metal hydrides + water → base + hydrogen

CaH2 (s) + 2 H2O(l) 

3. Hydrogen + nonmetal → molecular hydride

H2 (g) + Cl2 (g)  2 HCl (g)

2 H2 (g) + O2 (g) 

3 H2 (g) + N2 (g)  (note: this is called the Haber process)

4. molecular hydride + water → acid

HCl (g) + H2O (l) 

6-8 Oxygen and the Oxides:

1. Metal + oxygen → Ionic oxide

The probability of a group 1 metal to produce a peroxide or a superoxide increases going down the group.

4 Li (s) + O2 (g)  2 Li2O (s)

Na (s) + O2 (g) 

(note: this is a peroxide …this is sodium peroxide… not sodium oxide)

K (s) + O2 (g) 

(note: this is a superoxide … this is potassium superoxide…not potassium oxide)

Ca (s) + O2 (g) 

Increase the pressure and the following could happen:

Ca (s) + O2 (g) 

(note: this is a peroxide …this is calcium peroxide… not calcium oxide)

Oxygen does not form oxides with noble gases or with the noble metals (Au, Pd, and Pt). Metals that exhibit variable oxidation states react with a limited amount of oxygen to give oxides with lower oxidation states. They react with an excess of oxygen to give oxides with higher oxidation states.

For example, Iron combines with oxygen in the following series of reactions:

2 Fe (s) + O2 (g)  2 FeO (s)

(note: this is a Iron II Oxide… Fe+2 )

6 FeO (s) + O2 (g)  2 Fe3O4 (s)

4 Fe3O4 (s) + O2 (g)  6 Fe2O3 (s)

(note: this is a Iron III Oxide… Fe+3 )

2. Metal oxide + water → base Na2O (s) + H2O(l)  NaOH (aq)

CaO (s) + H2O(l) 

BaO (s) + H2O(l) 

Note: many metal oxides are relatively insoluble in water.

3. Nonmetal + oxygen → molecular oxide

2 C (s) + O2 (g)  2 CO (s)

(note: this is with excess Carbon and limited Oxygen)

2 C (s) + O2 (g)  2 CO2 (g)

(note: this is with limited Carbon and excess Oxygen)

P4 (s) + 3 O2 (g)  P4O6 (s)

(note: this is with limited Oxygen)

P4 (s) + 5 O2 (g)  P4O10 (s)

(note: this is with excess Oxygen)

S8 (s) + 8 O2 (g)  8 SO2 (g)

But sometimes the following occurs:

S8 (s) + 12 O2 (g)  8 SO3 (l)

4. Molecular oxide + water → acid

CO2 (g) + H2O(l)  H2CO3 (aq)

SO3 (l) + H2O(l) 

N2O5 (s) + H2O(l)  2 HNO3 (aq)

P4O10 (s) + 6 H2O(l)  4 H3PO4 (aq)

5. Metal oxide + nonmetal oxide → salt

CaO (s) + SO3 (l)  CaSO4 (s) MgO (s) + CO2 (g) 

6 Na2O (s) + P4O10 (s)  4 Na3PO4 (s)

6. Combustion reactions – redox reaction in which oxygen combines rapidly in a highly exothermic reaction with a visible flame.

C8H18 (l) + O2 (g) 

(note: this is with excess oxygen)

2 C8H18 (l) + 17 O2 (g)  16 CO (g) + 18 H2O (l)

(note: this is with limited oxygen)

2 C8H18 (l) + 9 O2 (g)  16 C (s) + 18 H2O (l)

(note: this is with very limited oxygen)

The combustion of pollutants containing sulfur…like coal:

2 S8 (s) + O2 (g)  8 SO2 (g)

2 PbS (s) + 3 O2 (g)  2 PbO (s) + 2 SO2 (g)

(note: this is called roasting and eventually leads to the isolation of Lead.)

2 SO2 (g) + O2 (g)  2 SO3 (l)

SO3 (l) + H2O (g) 

(note: oxides of sulfur are the main cause of acid rain.)