AP Chemistry Chapter 7 Lecture Notes
7.1 Development of the Periodic Table
•The periodic table is the most significant tool that chemists use for organizing and recalling chemical facts.
•Elements in the same column contain the same number of outer-shell electrons or valence electrons.
•How do we organize the different elements in a meaningful way that will allow us to make predictions about undiscovered
elements?
•Arrange elements to reflect the trends in chemical and physical properties.
•The periodic table arises from the periodic patterns in the electronic configurations of the elements.
•Elements in the same column contain the same number of valence electrons.
•The trends within a row or column form patterns that help us make predictions about chemical properties and
reactivity.
•In the first attempt Mendeleev and Meyer arranged the elements in order of increasing atomic weight.
•Certain elements were missing from this scheme.
•For example, in 1871 Mendeleev noted that As properly belonged underneath P and not Si, which left a
missing element underneath Si. He predicted a number of properties for this element.
•In 1886 Ge was discovered; the properties of Ge match Mendeleev’s predictions well.
•In the modern periodic table, elements are arranged in order of increasing atomic number.
7.2 Effective Nuclear Charge
•Effective nuclear charge (Zeff) is the charge experienced by an electron on a many-electron atom.
•The effective nuclear charge is not the same as the charge on the nucleus because of the effect of inner electrons.
•An electron is attracted to the nucleus, but repelled by electrons that shield or screen it from the full nuclear charge.
•This shielding is called the screening effect. OR THE SHIELDING EFFECT!
•The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of
electrons that block it from the nucleus.
•As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases.
Zeff = Z – S
•As the distance from the nucleus increases, S increases and Zeff decreases.
•S is called the screening constant which represents the portion of the nuclear charge that is screened from the
valence electron by other electrons in the atom.
•The value of S is usually close to the number of core electrons in an atom.
7.3 Sizes of Atoms and Ions
•Consider a collection of argon atoms in the gas phase.
•When they undergo collisions, they ricochet apart because electron clouds cannot penetrate each other to a
significant extent.
•The apparent radius is determined by the closest distances separating the nuclei during such collisions.
•This radius is the nonbonding radius.
•Now consider a simple diatomic molecule.
•The distance between the two nuclei is called the bonding atomic radius.
•It is shorter than the nonbonding radius.
•If the two atoms that make up the molecule are the same, then half the bond distance is called the covalent
radius of the atom.
Periodic Trends in Atomic Radii
•Atomic size varies consistently through the periodic table.
•As we move down a group the atoms become larger.
•As we move across a period, atoms become smaller.
•There are two factors at work:
•the principal quantum number, n, and
•the effective nuclear charge, Zeff.
•As the principal quantum number increases (i.e., we move down a group), the distance of the outermost
electron from the nucleus becomes larger. Hence the atomic radius increases.
•As we move across the periodic table, the number of core electrons remains constant, however, the nuclear
charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons.
This attraction causes the atomic radius to decrease.
Periodic Trends in Ionic Radii
•Ionic size is important:
•in predicting lattice energy.
•in determining the way in which ions pack in a solid.
•Just as atomic size is periodic, ionic size is also periodic.
•In general:
•Cations are smaller than their parent atoms.
•Electrons have been removed from the most spatially extended orbital.
•The effective nuclear charge has increased.
•Therefore, the cation is smaller than the parent atom.
•Anions are larger than their parent atoms.
•Electrons have been added to the most spatially extended orbital.
•This means total electron-electron repulsion has increased.
•Therefore, anions are larger than their parent atoms.
•For ions with the same charge, ionic size increases down a group.
•All the members of an isoelectronic series have the same number of electrons.
•As nuclear charge increases in an isoelectronic series the ions become smaller:
O2– > F– > Na+ > Mg2+ > Al3+
7.4 Ionization Energy
•The ionization energy of an atom or ion is the minimum energy required to remove an electron from the ground
state of the isolated gaseous atom or ion.
•The first ionization energy, I1, is the amount of energy required to remove an electron from a gaseous atom:
Na(g)  Na+(g) + e–
•The second ionization energy, I2, is the energy required to remove the second electron from a gaseous ion:
Na+(g)  Na2+(g) + e–
•The larger the ionization energy, the more difficult it is to remove the electron.
•There is a sharp increase in ionization energy when a core electron is removed.
Variations in Successive Ionization Energies
•Ionization energies for an element increase in magnitude as successive electrons are removed.
•As each successive electron is removed, more energy is required to pull an electron away from an
increasingly more positive ion.
•A sharp increase in ionization energy occurs when an inner-shell electron is removed.
Periodic Trends in First Ionization Energies
•Ionization energy generally increases across a period.
•As we move across a period, Zeff increases, making it more difficult to remove an electron.
•Two exceptions are removing the first p electron and removing the fourth p electron.
•The s electrons are more effective at shielding than p electrons. So, forming the s2p0 configuration is
more favorable.
•When a second electron is placed in a p orbital, the electron-electron repulsion increases. When this
electron is removed, the resulting s2p3 configuration is more stable than the starting s2p4 configuration.
Therefore, there is a decrease in ionization energy.
•Ionization energy decreases down a group.
•This means that the outermost electron is more readily removed as we go down a group.
•As the atom gets bigger, it becomes easier to remove an electron from the larger orbital.
•Example: for the noble gases, the ionization energies follow the order He > Ne > Ar > Kr > Xe
•The representative elements exhibit a larger range of values for I1 than transition metals.
Electron Configurations of Ions
•These are derived from the electron configurations of elements with the required number of electrons added or removed from
the most accessible orbital.
•
Li:
[He]2s1
becomes
Li+: [He]
•
F:
[He]2s22p5
becomes
F- :
[He]2s22p6 = [Ar]
•Transition metals tend to lose the valence shell electrons first and then as many d electrons as are required to reach
the desired charge on the ion.
•Thus electrons are removed from 4s before the 3d, etc.
7.5 Electron Affinities
•Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion.
•Electron affinity and ionization energy measure the energy changes of opposite processes.
•Electron affinity: Cl(g) + e–  Cl–(g)
∆ E = –349 kJ/mol
•Ionization energy: Cl(g)  Cl+(g ) + e–
∆E = 1251 kJ/mol
•Electron affinity can either be exothermic (as the above example) or endothermic:
Ar(g) + e–  Ar– (g)
E > 0
•Look at electron configurations to determine whether electron affinity is positive or negative.
•The extra electron in Ar needs to be placed in the 4s orbital which is significantly higher in energy than the
3p orbital.
•The added electron in Cl is placed in the 3p orbital to form the stable 3p6 electron configuration.
•Electron affinities do not change greatly as we move down in a group.
7.6 Metals, Nonmetals and Metalloids
•Metallic character refers to how much an element exhibits the physical and chemical properties of metals.
•Metallic character increases down a group.
•Metallic character decreases from left to right across a period.
Metals
•Metals are shiny and lustrous, as well as malleable and ductile.
•Metals are solids at room temperature and have very high melting temperatures (exceptions: mercury is liquid at room
temperature; gallium and cesium melt just above room temperature).
•Metals tend to have low ionization energies and tend to form cations easily.
•Metals tend to be oxidized when they react.
•Compounds of metals with nonmetals tend to be ionic substances.
•Metal oxides form basic ionic solids.
Metal oxide + water  metal hydroxide
•Most metal oxides are basic:
Na2O(s) + H2O(l)  2NaOH(aq)
•Metal oxides react with acids to form salts and water:
Metal oxide + acid  salt + water
NiO(s) + 2HCl(aq)  NiCl2(aq) + H2O(l)
Nonmetals
•Nonmetals are more diverse in their behavior than metals.
• Nonmetals are nonlustrous, poor conductors of heat and electricity, with lower melting points than metals.
•Seven nonmetallic elements exist as diatomic molecules under ordinary conditions:
•H2(g), N2(g), O2(g), F2(g), Cl2(g), Br2(l), I2(s)
Metal + nonmetal  salt
•When nonmetals react with metals, nonmetals tend to gain electrons:
2Al(s) + 3Br2(l)  2AlBr3 (s)
•Compounds composed entirely of nonmetals are molecular substances.
•Most nonmetal oxides are acidic:
Nonmetal oxide + water  acid
CO2(g) + H2O(l)  H2CO3(aq)
P4O10(s) + 6H2O(l)  4H3PO4(aq)
•Nonmetal oxides react with bases to form salts and water:
Nonmetal oxide + base  salt + water
CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)
Metalloids
•Metalloids have properties that are intermediate between those of metals and nonmetals.
•Example: Si has a metallic luster but it is brittle.
•Metalloids have found fame in the semiconductor industry.
7.7 Group Trends for the Active Metals
•The alkali metals (group 1A) and the alkaline earth metals (group 2A) are often called the active metals.
Group 1A: The Alkali Metals
•Alkali metals are all soft.
•Their chemistry is dominated by the loss of their single s electron: M  M+ + e–
•Reactivity increases as we move down the group.
•Alkali metals react with hydrogen to form hydrides.
•In hydrides, the hydrogen is present as H–, called the hydride ion. 2M(s) + H2(g)  2MH(s)
•Alkali metals react with water to form MOH and hydrogen gas: 2M(s) + 2H2O(l)  2MOH(aq) + H2(g)
•Alkali metals produce different oxides when reacting with O2:
•
4Li(s) + O2(g)  2Li2O(s) (oxide)
•
2Na(s) + O2 (g)  Na2O2(s) (peroxide)
•
K(s) + O2 (g)  KO2(s)
(superoxide)
•Alkali metals emit characteristic colors when placed in a high-temperature flame.
•The s electron is excited by the flame and emits energy when it returns to the ground state.
•The Na line occurs at 589 nm (yellow), characteristic of the 3p  3s transition.
•The Li line is crimson red.
•The K line is lilac.
Group 2A: The Alkaline Earth Metals
•Alkaline earth metals are harder and more dense than the alkali metals.
•Their chemistry is dominated by the loss of two s electrons:
M  M2+ + 2e–
•Reactivity increases down the group.
Mg(s) + Cl2(g) MgCl2(s)
•Be does not react with water.
2Mg(s) + O2(g)  2MgO(s)
•Mg will only react with steam.
•Ca and the elements below it react with water at room temperature as follows:
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
7.8 Group Trends for Selected Nonmetals
Hydrogen
•Hydrogen is a unique element.
•It most often occurs as a colorless diatomic gas, H2.
•Reactions between hydrogen and nonmetals can be very exothermic:
2H2(g) + 2O2(g)  2H2O(l)
∆Hf º = –571.7 kJ
•It can either gain another electron to form the hydride ion, H–, or lose its electron to become H+:
2Na(s) + H2(g)  2NaH(s)
2H2(g) + O2(g)  2H2O(l)
+
•H is a proton.The aqueous chemistry of hydrogen is dominated by H+(aq).
Group 6A: The Oxygen Group
•As we move down the group the metallic character increases. O2 is a gas, Te is a metalloid, Po is a metal.
•There are two important forms of oxygen: O2 and ozone, O3.
•O2 and O3 are allotropes.
•Allotropes are different forms of the same element in the same state (in this case, gaseous).
•Ozone can be prepared from oxygen:
3O2(g)  2O3(g)
∆Hfo = +284.6 kJ
•Ozone is pungent and toxic.
•Oxygen (or dioxygen, O2) is a potent oxidizing agent since the O2– ion has a noble gas configuration.
•There are two oxidation states for oxygen: –2 (e.g., H2O) and –1 (e.g., H2O2).
•Sulfur is another important member of this group.
•The most common form of sulfur is yellow S8.
•Sulfur tends to form S2– in compounds (sulfides).
Group 7A: The Halogens
•Group 7A elements are known as the halogens ("salt formers").
•The chemistry of the halogens is dominated by gaining an electron to form an anion:
X2 + 2e–  2X–
•Fluorine is one of the most reactive substances known:
2F2(g) + 2H2O(l)  4HF(aq) + O2(g)
∆H = –758.9 kJ
•All halogens consist of diatomic molecules, X2.
•Chlorine is the most industrially useful halogen.
•The reaction between chorine and water produces hypochlorous acid (HOCl), which is used to disinfect
swimming pool water:
Cl2(g) + H2O(l)  HCl(aq) + HOCl(aq)
•Halogens react with hydrogen to form gaseous hydrogen halide compounds:
H2(g) + X2  2HX(g)
•Hydrogen compounds of the halogens are all strong acids with the exception of HF.
Group 8A: The Noble Gases
•The group 8A elements are known as the noble gases.
•These are all nonmetals and monoatomic.
•They are notoriously unreactive because they have completely filled s and p subshells.
•In 1962 the first compounds of the noble gases were prepared: XeF2, XeF4, and XeF6.
HOMEWORK ASSIGNMENTS
#1- pg 292 #9, 13, 14, 15, 16, 18, 21, 22, 25, 26, 30,31
#2- pg 292 #5, 37, 38, 39, 42, 44, 46, 48, 49
#3- pg292- #54, 55, 57, 59, 60, 63, 65, 69, 76
AP Chemistry Chapter Seven Practice Test— Periodic Properties of the Elements
1) In general, as you go across a period in the periodic table
from left to right:
(1) the atomic radius __________;
(2) the electron affinity becomes _________ negative; and
(3) the first ionization energy __________.
A) decreases, decreasingly, increases
B) increases, increasingly, decreases
C) increases, increasingly, increases
D) decreases, increasingly, increases
E) decreases, increasingly, decreases
2) Element M reacts with chlorine to form a compound with the
formula MCl2 . Element M is more reactive than magnesium
and has a smaller radius than barium. This element is
__________.
A) Sr
B) K
C) Na
D) Ra
E) Be
3) The oxide of which element below can react with
hydrochloric acid?
A) sulfur
B) selenium
C) nitrogen
D) sodium
E) carbon
4) Metals can be __________ at room temperature.
A) liquid only
B) solid only
C) solid or liquid
D) solid, liquid, or gas
E) liquid or gas
10) Oxides of most nonmetals combine with water to form ___.
A) an acid
B) a base
C) water and a salt
D) water
E) hydrogen gas
11) Oxides of most nonmetals combine with base to form ___.
A) hydrogen gas
B) an acid
C) a base
D) water
E) water and a salt
12) The reaction of a metal with a nonmetal produces a(n) ___.
A) base
B) salt
C) acid
D) oxide
E) hydroxide
13) Of the hydrogen halides, only __________ is a weak acid.
A) HCl (aq)
B) HBr (aq)
C) HF (aq)
D) HI (aq)
E) They are all weak acids.
14) The only noble gas that does not have the ns 2 np6 valence
electron configuration is __________.
A) radon
B) neon
C) helium
D) krypton
2
6
E) All noble gases have the ns np valence configuration.
15) 2 F2 (g) + 2 H 2 O (l) ! __________.
A) 2 HF (aq) + 2 HFO (aq)
B) 2 F- (aq) + 2 H + (aq) + H 2 O 2 (aq)
C) 4 HF (aq) + O 2 (g)
5) Na reacts with element X to form an ionic compound with
the formula Na 3 X . Ca will react with X to form _________.
A) CaX 2
B) CaX
D) Ca 3 X 2
E) Ca 3 X
C) Ca 2 X3
6) The substance, __________ is always produced when an
active metal reacts with water.
A) NaOH
B) H 2 O
C) CO 2
D) H 2
E) O 2
7) What is the coefficient of H 2 O when the following equation
is completed and balanced?
Ba (s) + H 2 O (l) → ________
A) 1
B) 2
C) 3
D) 5
E) Ba(s) does not react with H 2 O (l).
8) Oxides of the active metals combine with water to form
________.
A) metal hydroxides
B) metal hydrides
C) hydrogen gas
D) oxygen gas
E) water and a salt
9) Oxides of the active metals combine with acid to form _____.
A) hydrogen gas
B) metal hydrides
C) water and a salt
D) oxygen gas
E) metal hydroxides
D) 2 HF2 (aq) + 2 OH - (aq)
E) 4 HF (aq) + 2 O2- (aq)
16) In which set of elements would all members be expected to
have very similar chemical properties?
A) O, S, Se
B) N, O, F
C) Na, Mg, K
D) S, Se, Si
E) Ne, Na, Mg
17) Electrons in the 1s subshell are much closer to the nucleus
in Ar than in He due to the larger __________ in Ar.
A) nuclear charge
B) paramagnetism
C) diamagnetism
D) Hund's rule
E) azimuthal quantum number
18) Of the following, which gives the correct order for atomic
radius for Mg, Na, P, Si and Ar?
A) Mg > Na > P > Si > Ar
B) Ar > Si > P > Na > Mg
C) Si > P > Ar > Na > Mg
D) Na > Mg > Si > P > Ar
E) Ar > P > Si > Mg > Na
19) Which of the following correctly lists the five atoms in
order of increasing size (smallest to largest)?
A) O < F < S < Mg < Ba
B) F < O < S < Mg < Ba
C) F < O < S < Ba < Mg
D) O < F < S < Ba < Mg
E) F < S < O < Mg < Ba
20) Screening by the valence electrons in
atoms is ________.
A) less efficient than that by core electrons
B) more efficient than that by core electrons
C) essentially identical to that by core
electrons
D) responsible for a general increase in
atomic radius going across a period
E) both more efficient than that by core
electrons and responsible for a general
increase in atomic radius going across
a period
28) Which ion below has the largest radius?
A) ClB) K +
C) Br -
21) Which one of the following atoms
has the largest radius?
A) O
B) F
C) S
D) Cl E) Ne
B) Cl- < Ar < K + < Ca 2+
22) In which of the following atoms is
the 3s orbital closest to the nucleus?
A) Br B) Cl C) At D) I
E) The 3s orbitals are the same distance
from the nucleus in all of these atoms.
23) Of the following elements, which has
the largest first ionization energy?
A) Na
B) Al
C) Se
D) Cl
E) Br
24) __________ have the lowest first
ionization energies of the groups listed.
A) Alkali metals B) Transition elements
C) Halogens
D) Alkaline earth metals
E) Noble gases
25) Which of the following has the
largest second ionization energy?
A) Si
B) Mg
C) Al
D) Na
E) P
26) Which equation correctly represents the first ionization of
aluminum?
A) Al- (g) ! Al (g) + eB) Al (g) ! Al- (g) + eC) Al (g) + e- ! Al- (g)
+
D) Al (g) ! Al (g) + e
-
E) Al+ (g) + e- ! Al (g)
27) Which of the following correctly represents the second
ionization of aluminum?
A) Al+ (g) + e- ! Al (g)
B) Al (g) ! Al+ (g) + eC) Al- (g) + e- ! Al2- (g)
D) Al+ (g) + e- ! Al2+ (g)
E) Al+ (g) ! Al2+ (g) + e-
D) F-
E) Na +
29) Of the compounds below, __________ has the smallest
ionic separation.
A) KF
B) K 2S
C) RbCl
D) SrBr2
E) RbF
30) Which isoelectronic series is correctly arranged in order of
increasing radius?
A) K + < Ca2+ < Ar < ClC) Ca 2+ < Ar < K + < ClD) Ca 2+ < K + < Ar < ClE) Ca 2+ < K + < Cl- < Ar
31) __________ is isoelectronic with argon and __________ is
isoelectronic with neon.
A) Cl- , FB) Cl- , Cl+
C) F+ , FD) Ne- , Kr +
E) Ne- , Ar +
32) Of the following elements, __________ has the most
negative electron affinity.
A) Na
B) Li
C) Be
D) N
E) F
33) Of the following elements, __________ has the most
negative electron affinity.
A) P
B) Al
C) Si
D) Cl
E) B
34) Chlorine is much more apt to exist as an anion than is
sodium. This is because __________.
A) chlorine is bigger than sodium
B) chlorine has a greater ionization energy than sodium does
C) chlorine has a greater electron affinity than sodium does
D) chlorine is a gas and sodium is a solid
E) chlorine is more metallic than sodium
35) Of the elements below, __________ is the most metallic.
A) sodium
B) barium
C) magnesium
D) calcium
E) cesium
36) Of the following metals, __________ exhibits multiple
oxidation states.
A) Al
B) Cs
C) V
D) Ca
E) Na
37) Which of these oxides is most basic?
A) K 2 O
B) Al2 O3
C) CO 2
D) MgO
E) Na 2 O
38) Of the following oxides, __________ is the most acidic.
A) CaO
B) CO 2
C) Al2 O3
D) Li 2 O
E) Na 2 O
Consider the following electron configurations to answer the
questions that follow:
(i) 1s 2 2s 2 2p6 3s1
(ii) 1s 2 2s 2 2p6 3s 2
(iii) 1s 2 2s 2 2p6 3s 2 3p1
(iv) 1s 2 2s 2 2p6 3s 2 3p 4
45) Which one of the following is not true about the alkali
metals?
A) They are low density solids at room temperature.
B) They all readily form ions with a +1 charge.
C) They all have 2 electrons in their valence shells.
D) They are very reactive elements.
E) They have the lowest first ionization energies of the elements.
(v) 1s 2 2s 2 2p6 3s 2 3p5
39) The electron configuration belonging to the atom with the
highest second ionization energy is __________.
A) (i)
B) (ii)
C) (iii)
D) (iv)
E) (v)
46) Alkaline earth metals __________.
A) have the smallest atomic radius in a given period
B) form monoanions
C) form basic oxides
D) exist as triatomic molecules
E) form halides with the formula MX
40) The electron configuration that belongs to the atom with the
lowest second ionization energy is __________.
A) (i)
B) (ii)
C) (iii)
D) (iv)
E) (v)
47) Alkali metals tend to be more reactive than alkaline earth
metals because __________.
A) alkali metals have lower densities
B) alkali metals have lower melting points
C) alkali metals have greater electron affinities
D) alkali metals have lower ionization energies
E) alkali metals are not more reactive than alkaline earth metals
41) The electron configuration of the atom with the most
negative electron affinity is __________.
A) (i)
B) (ii)
C) (iii)
D) (iv)
E) (v)
42) The electron configuration of the atom that is expected to
have a positive electron affinity is __________.
A) (i)
B) (ii)
C) (iii)
D) (iv)
E) (v)
43) Which of the following is not a characteristic of metals?
A) acidic oxides
B) low ionization energies
C) malleability
D) ductility
E) These are all characteristics of metals.
44) When two elements form a compound, the greater the
difference in metallic character between the two elements, the
greater the likelihood that the compound will be __________.
A) a gas at room temperature
B) a solid at room temperature
C) metallic
D) nonmetallic
E) a liquid at room temperature
48) Consider the general valence electron configuration of
ns 2 np5 and the following statements:
(i) Elements with this electron configuration are expected
to form -1 anions.
(ii) Elements with this electron configuration are expected
to have large positive electron affinities.
(iii) Elements with this electron configuration are
nonmetals.
(iv) Elements with this electron configuration form acidic
oxides.
Which statements are true?
A) (i) and (ii)
B) (i), (ii), and (iii)
C) (ii) and (iii)
D) (i), (iii,) and (iv)
E) All statements are true.
49) All of the following are ionic compounds except
__________.
A) K 2 O
B) Na 2SO 4
C) SiO 2
D) Li3 N
E) NaCl
50) Which one of the following compounds produces a basic
solution when dissolved in water?
A) SO 2
B) Na 2 O
C) CO 2
60) Explain the periodic trends of atomic size, ionic size,
ionization energy, and electron affinity based on atomic
structure. Include the concepts of effective nuclear charge and
core electron screening/shielding.
Atomic sizedefinition-
D) OF2
E) O 2
51) __________ is a unique element and does not truly belong
to any family.
A) Nitrogen
B) Radium
C) Hydrogen
D) Uranium
E) Helium
52) This element reacts with hydrogen to produce a gas with the
formula HX. When dissolved in water, HX forms an acidic
solution. X is __________,
A) Na
B) H
C) C
D) Br
E) O
53) Of the following elements, which have been shown to form
compounds?
helium
xenon
neon
argon
krypton
A) xenon and argon
B) xenon only
C) xenon, krypton, and argon
D) xenon and krypton
E) None of the above can form compounds.
54) Xenon has been shown to form compounds only when it is
combined with __________.
A) something with a tremendous ability to remove electrons
from other substances
B) another noble gas
C) something with a tremendous ability to donate electrons to
other substances
D) an alkali metal
E) an alkaline earth metal
55) Which noble gas has the highest first ionization energy?
56) Electron affinity measures how easy an atom gains an
electron.
57) The effective nuclear charge acting on an electron is larger
than the actual nuclear charge.
58) The atomic radius of iodine is one-half the distance
separating the iodine nuclei.
59) A group of ions all containing the same number of electrons
constitute an isoelectronic series.
TREND down a group-
TREND across a period (left to right)-
Ionic sizecations-
anions-
TRENDSisoelectronic seriesIonization energydefinitionTREND down a groupTREND across a periodSuccessive ionization energiesElectron affinitydefinitionpositive valuenegative valueTREND down a group-
TREND across a period-
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D
A
D
C
D
D
B
A
C
A
E
B
C
C
C
A
A
D
B
A
C
C
D
A
D
D
E
C
A
D
A
E
D
C
E
C
A
B
A
B
E
B
A
B
C
C
D
D
C
B
C
D
C
A
helium
TRUE
FALSE
TRUE
TRUE
ATOMIC SIZE- Xe is smaller than Li!!
definition- Bonding radius is half the distance between the two nuclei in a diatomic element. Bonding
radius cannot always be measured because not all atoms will form diatomic molecules. Non-bonding radius
measures how close two atoms can get to each other before electron cloud repulsions prohibit getting closer.
Non-bonding radius is the also the distance between atoms in a solid metal.
TREND down a group- Atomic size increases down a group because valence electrons are in a larger
atomic orbital. As quantum number n increases, the electron is less likely to be found near the nucleus and the
probability cloud gets larger. NOTE- Effective nuclear charge does increase slightly down a group because smaller
energy shells are more effective at screening valence electrons. An increase in Zeff would cause atoms to be smaller,
but the addition of another energy shell has much more effect on atomic size.
TREND across a period (left to right)- Atomic size decreases within a period as atomic number
increases. Because core electrons are responsible for the majority of screening (which reduces the nuclear charge to
the Effective Nuclear Charge, Zeff), elements in the same period experience nearly equal screening. If Zeff is equal to
Nuclear Charge(Z) – Screening (S), then Zeff increases as atomic number increases within a given period because
nuclear charge is increasing but screening is not. As effective nuclear charge increases, the electrons in an atom
are held closer to the nucleus thus reducing atomic size.
Ionic size- cations- Cations are always smaller than their parent atoms because the largest orbitals (containing
valence
electrons) are no longer filled.
anions- Anions are always bigger than their parent atoms because the addition of another electron increases
repulsions and requires that valence electrons spread out to accommodate the new electron(s). NOTE- There is only an
electron added, so effective nuclear charge is not increased.
TRENDS- Size trends for ions are similar to the trends for atoms, and for the same reasons. As you move
across a period from metals to non-metals ionic size decreases, due to increased Zeff, until you reach the non-metals.
Non-metals form anions, which are much bigger than cations, so at the non-metals there is a large increase in ionic
size. After this increase anions decrease in size due to increased Zeff.
isoelectronic series- Isoelectronic means having the same electron configuration. Because elements seek to
obtain noble gas configurations, most ions have a configuration like a noble gas. This means that anions directly
before a noble gas and cations directly after a noble gas, have the same configuration. The size of ions in an
isoelectronic series depends only on the nuclear charge because the screening is identical. As atomic number
increases, the size of ions in an isoelectronic series decreases due to an increase in Zeff.
Ionization energy- Energy required to remove an electron from a gaseous atom. 2nd ionization energy = energy
required to remove an electron from a +1 ion.
TREND down a group- Ionization energy decreases down a group because valence electrons are farther
from the nucleus where they experience a lower Zeff and are therefore easier to remove. As principal quantum number
n increases, there are more core electrons, which are very effective at screening valence electrons from the nucleus.
TREND across a period- As atomic number increases within a given period, effective nuclear charge
increases and the distance of valence electrons from the nucleus decreases. Both of these factors contribute to the
increase in ionization energy across a period. The higher the Zeff the harder it is to remove an electron. The closer the
electron is to the nucleus, the harder it is to remove.
Anomalies-Sometimes ionization energy does not increase across a period because the loss of an
electron is favored by the slight energy differences between sublevels in an energy level. Removing a single
electron from group 3A elements is easier than trends predict because the electron is alone in a p sublevel. Removing a
single electron from a group 6A element can be easier than predicted by trends because that electron is doubled up in
the p orbital (s2p4) and removing it minimizes repulsions.
Successive ionization energies- There is always a HUGE jump in energy required to remove a core
electron. Successive ionization energies can be used to identify an atom. If the second ionization energy is huge it is
an alkali metal. If the fourth ionization energy is huge, it is a group 3A element.
Electron affinity- Energy required to add an electron to a gaseous atom.
positive value- Energy is required to add the electron. This means it wont happen (non-spontaneous)
negative value- Means energy is released when an electron is added and that the anion formed is more
stable than the atom. A large negative electron affinity is correlated with spontaneous electron grabbing.
TREND down a group- Electrons are not as attracted to larger atoms because the nucleus is far away and
shielded, but as the atom gets bigger there is more room for added electrons and electron-electron repulsions are not as
significant. These two phenomena counteract each other so that no particular trend in electron affinity is seen within a
group.
TREND across a period??- Halogens have the highest negative electron affinities and group 6A elements
have the second highest because they are more stable as anions with a full octet. Noble gases have
positive electron affinities because a noble gas anion is energetically unfavorable and unstable therefore it
doesn’t form. Group 5A elements can have positive electron affinities (or low negatives) because the electron added is
the fourth electron in the p sublevel and requires doubling up. Group 2A elements have positive electron affinities
because the electron would be added to the empty p sublevel, which is higher in energy.