Download Ch. 7 Sections 7.10 and 7.12 Note Organizer

AP Chemistry
Chapter 7 Sections 7.10, 7.12 Note Organizer
The History of the Periodic Table, Periodic Trends in Atomic Properties
The History of the Periodic Table
• Originally constructed to represent the ____________ observed in the chemical
_______________ of the elements.
• First ______________ to recognize patterns was Johann Dobereiner (1780-1849).
• Noticed several groups of _____________ elements had similar properties, for example,
chlorine, bromine, and iodine.
• Tried to expand his model of ______________ but it was severely ______________.
• Next notable attempt was made by John Newlands in 1864.
• Suggested that elements should be arranged in ___________________.
• This was based on the idea that certain properties seemed to repeat for every ____________
element.
• Model did attempt to group based on _________________ but not generally successful.
• Present form of ________________ ___________ conceived by Julius Lothar Meyer (18301985) and Dmitri Mendeleev (1834-1907).
• Mendeleev is given most of the ______________ because he emphasized the table could be
used to predict the _________________ and _______________ of unknown elements.
• He published his table in 1872.
• Mendeleev predicted the __________________ and __________________ of the elements
gallium, scandium, and germanium from gaps in his periodic table.
• Germanium was ____________________ in 1886 and his predicted values and those
observed are in excellent _____________________.
• Mendeleev was also able to predict _______________ _____________ of several
_______________, including indium, beryllium and uranium.
• Mendeleev’s table was almost universally ________________ and remains one of the most
valuable of a chemist’s _______________.
• The fundamental difference between Mendeleev’s table and the modern periodic table is the
modern table uses ________________ _______________ to order the elements rather than
atomic mass.
Valence Electrons
• Valence electrons are the electrons in the _____________________ principal quantum level
(outermost _____________ ___________) of an atom.
• Electron configuration for nitrogen: 1s22s22p3
• The valence electrons for nitrogen are the 2s and 2p electrons; therefore, nitrogen has five
_________________ electrons.
• Valence electrons are ___________________ because they are involved in _____________.
• Core electrons are the ________________ electrons.
• Elements with the same _______________ configuration show similar ________________
behavior.
• Groups 1, 2, 13-18 are often called the _____________________ or representative elements.
• Every _______________ of these groups has the ___________ valence electron
configuration.
• Predicting the valence electron configurations of the ______________ metals, the
___________________, and the ___________________ is somewhat more _____________
because of the many exceptions.
Periodic Trends
• There are observed ______________ in several important __________________
_________________: ionization energy, electron affinity, and atomic size.
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Effective Nuclear Charge
The effective ______________ charge is the __________ that an electron “feels” from the
nucleus.
Effective Nuclear Charge (Zeff) = # _________________ - # core __________________
The ______________ an electron is to the nucleus, the more __________ it feels.
As effective nuclear charge _________________, the electron cloud is pulled in
_______________.
Ionization Energy
Ionization energy is the _________________ required to ____________ an electron from a
gaseous atom or ion:
X (g) → X+ (g) + eConsider the _________________ required to ____________ several electrons from
aluminum in the gaseous state.
Al (g) → Al+ (g) + eI1 = 580 kJ/mol
Al+ (g) → Al2+ (g) + e- I2 = 1815 kJ/mol
Al2+ (g) → Al3+ (g) + e- I3 = 2740 kJ/mol
Al3+ (g) → Al4+ (g) + e- I4 = 11,600 kJ/mol
The _________________ energy electron (the one bound least tightly is removed ________.
I1 is the _______________ ionization energy and for aluminum, this electron comes from the
3p orbital ([Ne]3s23p1).
I2 is the ______________ ionization energy and this electron comes from the 3s orbital.
Why is I1 _________________ than I2?
The first electron is removed from a ______________ atom and the second is removed from
a 1+ ____________.
The ________________ in ______________ charge binds the electrons more firmly and it
takes more energy to remove an electron.
Why is I4 so high?
The fourth electron is “____________” electron (Al3+ = 1s22s22p6) and core electrons are
bound more ________________ than valence electrons.
In general as we go across a period from left to right, the ___________ ionization energy
________________.
Reason: __________________ in effective nuclear ______________ (more protons in
nucleus) felt by the valence electrons across a period.
Causes the valence electrons to be held more ____________, which makes it more
________________ to remove them.
Note: there are ____________________ in ionization energy trends in going across a period.
Due to shielding and electron repulsions.
First ionization energy _______________ in going down a ____________.
Reason: going down a group the _______________ being removed are, on average,
_________________ from the nucleus.
As n increases, the size of the orbital increases, and the electrons are _______________ from
the nucleus, and thus are _________________ to remove.
Electron Affinity
Electron affinity is the _____________ _______________ associated with the
______________ of an electron to a gaseous atom:
X (g) + e- → X- (g)
If the addition of the electron is ___________________, the corresponding value for electron
affinity will carry a ___________________ sign.
The incoming electron experiences an __________________ to the _______________,
which causes the potential energy to be lowered as the electron approaches the atom.
The trends in electron affinity are ________________ to those for ionization energy.
Electron affinity becomes more _______________________ from left to right across a
period. A valence shell that holds its electrons tightly will also tend to bind an additional
electron tightly.
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Electron affinity becomes _______ ______________ down a group. A valence shell that
loses electrons easily (low IE) will have little attraction for additional electrons (small EA).
Note: there are exceptions.
Atomic Radius
The ______________ of an atom (r) is defined as half the distance between the nuclei in a
molecule consisting of _______________ atoms.
For ___________________ atoms that do not form diatomic molecules, the atomic radii are
estimated from their various ___________________ compounds.
The radii for _______________ atoms (metallic radii) are obtained from half the distance
between metal atoms in solid metal ___________________.
Atomic radii _____________ in going from left to right _____________ a ____________.
Due to ____________________ effective _______________ charge in going from left to
right. Valence electrons are drawn closer to the nucleus, decreasing the size of the atom.
Atomic radius _________________ down a _______________, because of the increases in
the orbital sizes in successive principal quantum levels.
Trends in the Sizes of Ions
Negative ions are always ________________ than the atoms from which they are formed.
When electrons are _______________ to an atom, the mutual repulsions between them
increase.
The causes the electrons to _____________ apart and occupy a _____________ volume.
Positive ions are always _________________ than the atoms from which they are formed.
When electrons are removed from the valence shell, the electron-electron repulsions
__________________, which allows the remaining electrons to be pulled closed together
around the nucleus.
Electronegativity
Valence electrons hold atoms together in chemical compounds.
In many compounds, the negative charge of the valence electrons is concentrated closer to
one atom than to another.
This uneven concentration of charge has a significant effect on the chemical properties of a
compound.
_____________________ is a measure of the __________ of an atom in a chemical
compound to ___________ electrons (the most electronegative element is ___________).
Electronegativity ______________ across each period.
Electronegativity _______________ or stays the ________ down a group.