General Periodic Trends Atomic and ionic size Atomic and ionic size Ionization energy Ionization...
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General Periodic Trends • Atomic and ionic size • Ionization energy • Electron Affinity • Electronegativity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly.
General Periodic Trends Atomic and ionic size Atomic and ionic size Ionization energy Ionization energy Electron Affinity Electron Affinity Electronegativity
General Periodic Trends Atomic and ionic size Atomic and ionic
size Ionization energy Ionization energy Electron Affinity Electron
Affinity Electronegativity Electronegativity Higher effective
nuclear charge Electrons held more tightly Larger orbitals.
Electrons held less tightly.
Slide 3
Atomic Size Size goes UP on going down a group. Size goes UP on
going down a group. Because electrons are added further from the
nucleus, there is less attraction. This is due to additional energy
levels and the shielding effect. Each additional energy level
shields the electrons from being pulled in toward the nucleus.
Because electrons are added further from the nucleus, there is less
attraction. This is due to additional energy levels and the
shielding effect. Each additional energy level shields the
electrons from being pulled in toward the nucleus. Size goes DOWN
on going across a period. Size goes DOWN on going across a period.
Size goes UP on going down a group. Size goes UP on going down a
group. Because electrons are added further from the nucleus, there
is less attraction. This is due to additional energy levels and the
shielding effect. Each additional energy level shields the
electrons from being pulled in toward the nucleus. Because
electrons are added further from the nucleus, there is less
attraction. This is due to additional energy levels and the
shielding effect. Each additional energy level shields the
electrons from being pulled in toward the nucleus. Size goes DOWN
on going across a period. Size goes DOWN on going across a
period.
Slide 4
Atomic Size Size decreases across a period owing to increase in
the positive charge from the protons. Each added electron feels a
greater and greater + charge because the protons are pulling in the
same direction, where the electrons are scattered. Large Small
Slide 5
Slide 6
Which is Bigger? Na or K ? Na or K ? Na or Mg ? Na or Mg ? Al
or I ? Al or I ?
Slide 7
Ion Sizes Does the size go up or down when losing an electron
to form a cation? Does the size go up or down when losing an
electron to form a cation?
Slide 8
Ion Sizes CATIONS are SMALLER than the atoms from which they
come. CATIONS are SMALLER than the atoms from which they come. The
electron/proton attraction has gone UP and so size DECREASES. The
electron/proton attraction has gone UP and so size DECREASES.
Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Forming a cation.
Slide 9
Ion Sizes Does the size go up or down when gaining an electron
to form an anion?
Slide 10
Ion Sizes ANIONS are LARGER than the atoms from which they
come. ANIONS are LARGER than the atoms from which they come. The
electron/proton attraction has gone DOWN and so size INCREASES. The
electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes. Trends in ion sizes
are the same as atom sizes. Forming an anion. F, 71 pm 9e and 9p F
-, 133 pm 10 e and 9 p -
Slide 11
Trends in Ion Sizes Figure 8.13
Slide 12
Which is Bigger? Cl or Cl - ? Cl or Cl - ? K + or K ? K + or K
? Ca or Ca +2 ? Ca or Ca +2 ? I - or Br - ? I - or Br - ?
Slide 13
Mg (g) + 738 kJ ---> Mg + (g) + e- This is called the FIRST
ionization energy because we removed only the OUTERMOST electron Mg
+ (g) + 1451 kJ ---> Mg 2+ (g) + e- This is the SECOND IE. IE =
energy required to remove an electron from an atom (in the gas
phase). Ionization Energy
Slide 14
Trends in Ionization Energy IE increases across a period
because the positive charge increases. IE increases across a period
because the positive charge increases. Remember, low energy means
easier. Removing an electron from Li than F Remember, low energy
means easier. Removing an electron from Li than F Metals lose
electrons more easily than nonmetals. Metals lose electrons more
easily than nonmetals. Nonmetals lose electrons with difficulty
(they like to GAIN electrons). Nonmetals lose electrons with
difficulty (they like to GAIN electrons).
Slide 15
Trends in Ionization Energy IE increases UP a group IE
increases UP a group Harder to remove an electron closer to the
nucleus Harder to remove an electron closer to the nucleus Because
size increases (Shielding Effect) Because size increases (Shielding
Effect)
Slide 16
Which has a higher 1 st ionization energy? Mg or Ca ? Al or S ?
Cs or Ba ?
Slide 17
Electron Affinity Definition - the energy change associated
with the addition of an electron The sign of the electron affinity
can be confusing. When an atom becomes less stable upon gaining an
electron, its potential energy increases, which implies that the
atom gains energy as it acquires the electron. In such a case, the
atom's electron affinity is positive. An atom with a negative
electron affinity is far more likely to gain electrons.
Slide 18
Periodic Trend: Electron Affinity
Slide 19
ELECTRON AFFINITY: ACROSS A PERIOD Electron affinities becoming
increasingly negative from left to right. (Easier to add e- to Cl
than Na) Just as in ionization energy, this trend conforms to and
helps explain the octet rule. The octet rule states that atoms with
close to full valence shells will tend to gain electrons. Such
atoms are located on the right of the periodic table and have very
negative electron affinities, meaning they give off a great deal of
energy upon gaining an electron and become more stable. Be careful,
though: the noble gases, located in the extreme right hand column
of the periodic table do not conform to this trend. Noble gases
have full valence shells, are very stable, and do not want to add
more electrons: noble gas electron affinities are positive.
Similarly, atoms with full subshells also have more positive
electron affinities (are less attractive of electrons) than the
elements around them.
Slide 20
ELECTRON AFFINITY: DOWN A PERIOD DECREASES AS YOU GO DOWN
(easier to add to Na than to K) Electron affinities change little
moving down a group, though they do generally become slightly more
positive (less attractive toward electrons). This is because there
is less attractions between the electrons you are trying to add and
the nucleus because there is more distance between the electron
shell and nucleus The biggest exception to this rule are the third
period elements, which often have more negative electron affinities
than the corresponding elements in the second period. For this
reason, Chlorine, Cl, (group VIIa and period 3) has the most
negative electron affinity.
Slide 21
Electronegativity, is a measure of the ability of an atom in a
molecule to attract electrons to itself. Concept proposed by Linus
Pauling 1901-1994 Concept proposed by Linus Pauling 1901-1994
Slide 22
Periodic Trends: Electronegativity In a group: Atoms with fewer
energy levels can attract electrons better (less shielding). So,
electronegativity increases UP a group of elements. In a period:
More protons, while the energy levels are the same, means atoms can
better attract electrons. So, electronegativity increases RIGHT in
a period of elements.
Slide 23
Electronegativity
Slide 24
Which is more electronegative? F or Cl ? Na or K ? Sn or I
?
Slide 25
Summary of Periodic Trends
Slide 26
Periodic Trends Summary TrendAs Move to the right Reason?As
Move down Reason? Atomic RadiusGets Smaller-Effective Nuclear
Charge Increases with no change in shielding Gets Larger-More
energy levels filled, more shielding between nucleus and electrons
Ionization Energy Energy Increases (harder to remove) -Octet rule:
easier to remove unpaired e- -Coulombs Law (has greater Effective
Nuclear Charge) Energy Decreases (easier to remove) -Electrons are
farther and have more shielding from Nucleus and easier to remove
according to Coloumbs Law Electron Affinity Energy Decreases
(easier to add) -Octet rule: easier to add e- to almost full
orbitals -Greater Effective nuclear charge, more attraction to add
electron Energy Increases (harder to add) -Electrons are farther
and have more shielding from nucleus and will have smaller
attractive forces to nucleus Electro- negativity Decreases (Easier
to attract electrons) -Effective nuclear charge increases=more
attracting to bonding electrons Increases (Harder to attract
electrons) - Electrons are farther away, more shielding, less
attraction to nucleus by bonding electrons