9.2 Metals Nonmetals & their Ions 9-2 Metals & Nonmetals & Their ...
4s versus 3d Problem!
9.2 Metals Nonmetals & their Ions
9-2 Metals & Nonmetals & Their Ions
According to the Aufbau Principle and atomic spectroscopy, the 4s orbital fills before the 3d one. Once we start looking at the actual chemistry of the transition metals and their cations, however, we find that the 4s electrons are removed first. To help you remember this chemical aspect, we decided to write 3d before 4s in the energy sequence of the orbitals for the transition metals.
Platinum Sulfur
For example, the ground state electronic configuration of iron is [Ar]4s23d6, which we now write as [Ar]3d64s2. Then, when we go on to form Fe2+, the ground state electronic configuration of the Fe2+ cation will be [Ar]3d6.
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Make predictions based on e- configurations Metals
Non Metals
Metals tend to give up electrons Metals
Noble gases
9.2 Metals Nonmetals & their Ions Nonmetals tend to gain electrons
9.2 Metals Nonmetals & their Ions Noble gases neither gain nor lose e-
Non Metals
Noble gases
9.2 Metals Nonmetals & their Ions Recall noble gas configuration
9.2 Metals Nonmetals & their Ions
ns2np6
Noble gases: complete octet ns2np6 Eight valence electrons: complete shell
*
Noble gas configuration is very stable Kr (Z=36) [Ar]3d10 4s Noble gases
4p
[Ar]3d104s24p6 Metals lose electrons to achieve ns2np6 stability Nonmetals gain electrons to achieve ns2np6 stability
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Metals tend to lose electrons
Metals tend to lose electrons 19
K
39.098
[Ar]4s1
K [Ar]4s1
K+ + 1e[Ar]
Noble gas configuration for the cation
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Metals tend to lose electrons
Metals tend to lose electrons 56
Ba
137.33
[Xe]6s2
Ba [Xe]6s2
Ba2+ + 2e[Xe]
Noble gas configuration for the cation
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Nonmetals tend to gain electrons
Nonmetals tend to gain electrons 35
Br
79.904
Br + 1e[Ar]3d104s24p5
Br [Ar]3d104s24p6 = [Kr]
[Ar]3d104s24p5 Noble gas configuration for the anion
9.2 Metals Nonmetals & their Ions Nonmetals tend to gain electrons 16
S
32.06
Kryptonite
[Ne]3s23p4
9.2 Metals Nonmetals & their Ions Nonmetals tend to gain electrons
9.2 Metals Nonmetals & their Ions Transition metal ions also lose electrons 22
S + 2e2 [Ne]3s 3p4
S2[Ne]3s23p6 = [Ar]
Noble gas configuration for the anion
Ti
47.88
[Ar]3d24s2
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Transition metal ions can achieve noble gas
Most transition metal ions do not achieve
configuration e.g., Ti4+
noble gas configuration e.g. Fe2+, Fe3+
Ti [Ar]3d24s2 Ti [Ar]3d24s2
Ti2+ + 2e[Ar]3d2 Ti4+ + 4e[Ar]
Fe2+ + 2e-
Fe [Ar]3d64s2
[Ar]3d6 Fe3+ + 3e-
Fe [Ar]3d64s2
[Ar]3d5
Noble gas configuration
9.2 Metals Nonmetals & their Ions Which would you expect to be more stable Fe2+ or Fe3+?
Fe2O3 Fe3+
Answer: Fe3+ Why?
Extra stability associated with a half-filled 3d shell.
9.2 Metals Nonmetals & their Ions Most transition metal ions do not achieve noble gas configuration Cu [Ar]3d104s1
Cu+ + 1e[Ar]3d10 = [Ne]3s23p63d10 Pseudo noble gas configuration 18e- in outer shell
9.3 The Sizes of Atoms & Ions
9-3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions There are size trends for atoms and ions
9.3 The Sizes of Atoms & Ions How do we define atomic, ionic radius? Electron probability cloud…border uncertain We can determine internuclear distances We define radii based on distance between two nuclei
9.3 The Sizes of Atoms & Ions Covalent Radius (for covalent bonded atoms) r
r
O2 (oxygen molecule)
Internuclear distance is 2 x atomic radius (same species)
2r = 143 pm So r = 71.5 pm
9.3 The Sizes of Atoms & Ions Ionic Radius (for ionic bonded species) Na+
Cl-
1893 Windsor, Ontario
9.3 The Sizes of Atoms & Ions Metallic Radius (for metal atoms in crystal) Crystalline Solid Metal
Na+Cl- (ionic compound) Ag
Ag r
rNa+ rCl99pm 181pm
r
rNa+ + rCl- = 280 pm Internuclear dist.
Internuclear distance is sum of the two radii If you know rNa+…you can calculate rCl-
2r
2r = 288 pm So r = 144 pm Internuclear distance is 2 x metallic radius (same species)
9.3 The Sizes of Atoms & Ions The ions in crystalline LiI are arranged as
9.3 The Sizes of Atoms & Ions What are the radii of Li+ and of I-? a
shown below. If a = 600pm, what is the radius of Li+? What is the radius of I-?
r-
Ia=600pm Li+
9.3 The Sizes of Atoms & Ions The ions in crystalline LiI are arranged as shown below. If a = 600pm, what is the radius of Li+? What is the radius of I-? IrI- = 212pm
a=600pm
Li+ + rLi = 88pm
-
r c r-
a r 2r+ a=600pm rrrLi+ = 88 pm
c
a = 2rI- + 2rLi+ 600pm=2*212pm+2rLi+ 2 2 2 2 c = a + a = 2a = 2(600pm)2 (Pythagorean) c = 849 pm = 4rIrI- = 212 pm
8.10 Multielectron Atoms Electron Screening Outer eexperience a lower effective nuclear charge
Inner shell electrons screen outer efrom full attraction of the nucleus
Zeff < Z Zeff = Z - S
where S is charge that is screened
8.10 Multielectron Atoms
9.3 The Sizes of Atoms & Ions
s orbitals are better at screening than p & d Screening s orbitals > p orbitals > d orbitals strength s Zeff > p Zeff > d Zeff (for same n) Orbital energy depends on n and Zeff 2 - Zeff hcR H En = n2
Valence electrons will shield each other to a Much Smaller degree than inner shell electrons
Atomic Radius (pm)
Can we explain the size trends? Atomic Number
9.3 The Sizes of Atoms & Ions
8.12 e- Configurations & The Periodic Table s block
Atomic Radius (pm)
r
r p block
Putting in valence electrons So, S ∼ constant while Z increases Zeff increases and radius decreases
Atomic radius usually decreases from left to right across a period in s & p blocks Atomic Number
9.3 The Sizes of Atoms & Ions
Filling ns orbitals
Filling np orbitals
9.3 The Sizes of Atoms & Ions There are size trends for atoms and ions
Atomic Radius (pm)
Increasing n down a group…orbitals get larger so radius increases r Atomic radius usually increases from top to bottom down a group Atomic Number
9.3 The Sizes of Atoms & Ions
8.12 e- Configurations & The Periodic Table Putting electrons in an inner shell So, S increases as Z increases Zeff ∼ constant and radius ∼ constant d block Transition metals
Atomic Radius (pm)
Atomic radius often stays relatively constant across a period for transition metals Atomic Number
Filling (n-1)d r ∼ const. orbitals
9.3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions
Refer to the periodic table and arrange N, O, and P in order of increasing atomic radius
For main group atoms radius decreases across a period
Predict O < N < P Atomic radius increases down a group
9.3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions
Refer to the periodic table and arrange N, O,
Cations are smaller than the parent atoms
and P in order of increasing atomic radius
Cations…lose electron(s)…Z stays constant
Predicted: O < N < P Actual: 73pm 75pm 110pm
Na Na+
r = 186 pm
9.3 The Sizes of Atoms & Ions
+ 1e-
r = 99 pm
9.3 The Sizes of Atoms & Ions
Anions are larger than the parent atoms
For isoelectronic cations…the more positive
Anion…gain electron(s)…Z stays constant
charge…the smaller the ionic radius
ClCl
+ 1e-
r = 99 pm
r = 181 pm
9.3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions
For isoelectronic cations…the more positive
For isoelectronic anions…the more negative
charge…the smaller the ionic radius
charge…the larger the ionic radius Cl- [Ne]3s23p6 P3- [Ne]3s23p6 isoelectronic Z=17 Z=15
Na+ 1s22s22p6 Mg2+ 1s22s22p6 isoelectronic Z=11 Z=12
P3-
ClMg2+
Na+ r = 99 pm
r = 72 pm r = 181 pm
9.3 The Sizes of Atoms & Ions Refer to the periodic table and arrange these species in order of increasing size: Na+,
r = 212 pm
9.3 The Sizes of Atoms & Ions Determine e- configurations from table All species have [Ne] configuration
Mg2+, O2-, F-, Ne Na+Mg2+
O2-F-Ne
For isoelectric configurations radius will increase with decreasing Z
9.3 The Sizes of Atoms & Ions Refer to the periodic table and arrange these species in order of increasing size: Na+, Mg2+, O2-, F-, Ne Answer: Mg2+ < Na+ < Ne < F- < O2Z=12 11 10 9 8 Increasing radius Actual: 72pm 99pm 71pm 133pm 140pm ?
9.3 The Sizes of Atoms & Ions Tutorial will cover other typical questions
9.4 Ionization Energy
9.4 Ionization Energy
9-4 Ionization Energy
Ionization energy for H-like species
Z 2 hcR H ΔE IE = 12
ΔEIE = ? e-
Energy required to strip an electron from a gaseous state atom (or ion)
9.4 Ionization Energy
n=2 H-like species Only one electron Atom absorbs hν photon ΔEIE=hν From ground state n=1
E
9.4 Ionization Energy
Definition Ionization Energy = I
In = nth Ionization energy...remove nth e-
Energy required to strip an electron from a gaseous state atom (or ion) e-
9.4 Ionization Energy
Al+(g) + 1e-
Al(g)
I1=577.6 kJ/mol
Al+(g)
Al2+(g) + 1e- I2=1,817 kJ/mol
Al2+(g)
Al3+(g) + 1e- I3=2,745 kJ/mol
Al3+(g)
Al4+(g) + 1e- I4=11,580 kJ/mol
Not spontaneous…requires energy input
[Ne] noble gas configuration…extra stable
9.4 Ionization Energy
I depends on Zeff and n
Periodic Trends for I I generally increases across a period
2 Zeff hcR H n2
Larger Zeff…higher I…e- held more tightly
I1 kJ/mol
I=
n=∞ n=5 n=4 n=3
I increases as Zeff increases across a period
Higher n…lower I…outer e- held less tightly 2 I ∝ Zeff and I ∝
1 n2
Atomic No. (Z)
8.12 e- Configurations & The Periodic Table s block
9.4 Ionization Energy Periodic Trends for I I generally decreases down a group
I I
p block
Filling ns orbitals
I1 kJ/mol
Putting in valence electrons So, S ∼ constant while Z increases Zeff increases
I decreases as n increases down a group
Filling np orbitals Atomic No. (Z)
9.4 Ionization Energy There are IE trends n I Zeff increases only slightly down a group
9.4 Ionization Energy
There are some exceptions
9.3 The Sizes of Atoms & Ions We would predict I1 Mg < I1 Al
9.4 Ionization Energy
9.4 Ionization Energy Remember that electrons are ionized from
There are some exceptions
the orbital with the highest n I1 = 737.7 kJ/mol Mg
This is not necessarily the last orbital to be Al is lower???
I1 = 577.6 kJ/mol Al Why lower? Mg
If there are different types of orbitals with the
[Ne]3s2
Easier to pull the electron from a higher energy p orbital than a lower energy s orbital
Al [Ne]3s23p1
9.4 Ionization Energy Example Sc Sc
Sc2+
First
[Ar]3d14s2 Sc+
[Ar]3d14s1 [Ar]3d1
same n, the e- is ionized from the highest energy orbital Lower energy s < p < d Higher energy
9.4 Ionization Energy Typical problems…refer to the periodic table
Last e- added
[Ar]3d14s2
Sc+
filled
and arrange elements in expected order of removed
[Ar]3d14s1
Sc2+ Sc3+
e-
[Ar]3d1
[Ar] +
+
+
1e-
increasing first ionization energy I1 Attend tutorials for examples
1e-
1e-
9.5 Electron Affinity
9.5 Electron Affinity
9-5 Electron Affinity
Definition Electron Affinity = EA
ΔEEA = ? e-
Energy change when an electron is added to a gaseous state atom (or ion)
e-
Energy change when an electron is added to a gaseous state atom (or ion)
9.5 Electron Affinity
9.5 Electron Affinity
EA is often negative This means energy is given off
F(g) + 1e1s22s22p5
(exothermic)
F-(g)
EA=-328 kJ/mol
1s22s22p6
Note that ionization energy was always positive…it required input of energy
9.5 Electron Affinity
9.5 Electron Affinity
EAn = nth Electron Affinity...add nth e-
EAn = nth Electron Affinity...add nth e-
O(g) + 1e-
O-(g)
EA1=-141.0 kJ/mol
O(g) + 1e-
O-(g)
EA1=-141.0 kJ/mol
O-(g) + 1e-
O2-(g)
EA2=+744 kJ/mol
O-(g) + 1e-
O2-(g)
EA2=+744 kJ/mol
Positive because it is hard to add a second electron due to repulsion
9.5 Electron Affinity
Electron affinity trends are harder to discern It is easy to get confused! A high affinity for electrons means a large negative electron affinity
9.5 Electron Affinity EA tends to become less negative (lower affinity) down a Group
9.5 Electron Affinity Highest EA values are found in group 17…ns2np5…readily gain one e-
9.7 Periodic Properties of the Elements
9.7 Periodic Properties of the Elements
9-7 Periodic Properties of the Elements
Summary of Atomic Properties Fig.9.11 EA becomes more negative
Arrows indicate increasing
We can understand the trends in terms of Zeff (screening); n (orbital size); e- repulsion
9.7 Periodic Properties of the Elements Periodicity…Predict Properties
9.7 Periodic Properties of the Elements
Physical properties often change uniformly down a group e.g., melting & boiling points
Periodicity…Predict Properties
9.7 Periodic Properties of the Elements Periodicity…Predict Properties
Halogens diatomic e.g., Cl2 Br2 I2
9.7 Periodic Properties of the Elements Periodicity…Predict Properties
Can we predict the melting point of Br2? Estimated mp = (172K+387K)/2 = 280K Actual mp for Br2 = 266K
9.7 Periodic Properties of the Elements Periodicity…Predict Properties Table 9.5
Can we predict the melting point of Br2? Estimated mp = (172K+387K)/2 = 280K Not Too bad! Actual mp for Br2 = 266K
9.7 Periodic Properties of the Elements As you learn reactions you will see trends
9.7 Periodic Properties of the Elements Periodicity…Predict Properties Table 9.5
Try predicting the boiling point of Br2 by averaging Example 9-5 pg. 360
LEO the lion says GER
LEO = loss of electrons Æ Oxidation
9.7 Periodic Properties of the Elements As you learn reactions you will see trends Groups 1 & 2 Metals…Lose ee- loss is Oxidation Group 1 & 2 metals are generally good “reducing agents”
9.7 Periodic Properties of the Elements As you learn reactions you will see trends Group 17 Nonmetals…gain eThey undergo reduction (gain of electrons) They act as oxidizing agents (they are reduced but drive oxidation in another species)
GER = gain of electrons Æ Reduction
9.7 Periodic Properties of the Elements As you learn reactions you will see trends Groups 1 & 2 Metals…lose ee- loss is Oxidation Group 1 & 2 Metals are generally good “reducing agents” A reducing agent undergoes oxidation…but drives reduction (gain of e- by another species)
9.7 Periodic Properties of the Elements In Science, we use models to predict & we test the predictions
Periodic Table is central for understanding Chemistry
The End!
THAT’ THAT’S ALL FOLKS!
GOOD LUCK IN THE REST OF THE CHEM 110 COURSE
9.2 Metals Nonmetals & their Ions
9-2 Metals & Nonmetals & Their Ions
According to the Aufbau Principle and atomic spectroscopy, the 4s orbital fills before the 3d one. Once we start looking at the actual chemistry of the transition metals and their cations, however, we find that the 4s electrons are removed first. To help you remember this chemical aspect, we decided to write 3d before 4s in the energy sequence of the orbitals for the transition metals.
Platinum Sulfur
For example, the ground state electronic configuration of iron is [Ar]4s23d6, which we now write as [Ar]3d64s2. Then, when we go on to form Fe2+, the ground state electronic configuration of the Fe2+ cation will be [Ar]3d6.
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Make predictions based on e- configurations Metals
Non Metals
Metals tend to give up electrons Metals
Noble gases
9.2 Metals Nonmetals & their Ions Nonmetals tend to gain electrons
9.2 Metals Nonmetals & their Ions Noble gases neither gain nor lose e-
Non Metals
Noble gases
9.2 Metals Nonmetals & their Ions Recall noble gas configuration
9.2 Metals Nonmetals & their Ions
ns2np6
Noble gases: complete octet ns2np6 Eight valence electrons: complete shell
*
Noble gas configuration is very stable Kr (Z=36) [Ar]3d10 4s Noble gases
4p
[Ar]3d104s24p6 Metals lose electrons to achieve ns2np6 stability Nonmetals gain electrons to achieve ns2np6 stability
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Metals tend to lose electrons
Metals tend to lose electrons 19
K
39.098
[Ar]4s1
K [Ar]4s1
K+ + 1e[Ar]
Noble gas configuration for the cation
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Metals tend to lose electrons
Metals tend to lose electrons 56
Ba
137.33
[Xe]6s2
Ba [Xe]6s2
Ba2+ + 2e[Xe]
Noble gas configuration for the cation
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Nonmetals tend to gain electrons
Nonmetals tend to gain electrons 35
Br
79.904
Br + 1e[Ar]3d104s24p5
Br [Ar]3d104s24p6 = [Kr]
[Ar]3d104s24p5 Noble gas configuration for the anion
9.2 Metals Nonmetals & their Ions Nonmetals tend to gain electrons 16
S
32.06
Kryptonite
[Ne]3s23p4
9.2 Metals Nonmetals & their Ions Nonmetals tend to gain electrons
9.2 Metals Nonmetals & their Ions Transition metal ions also lose electrons 22
S + 2e2 [Ne]3s 3p4
S2[Ne]3s23p6 = [Ar]
Noble gas configuration for the anion
Ti
47.88
[Ar]3d24s2
9.2 Metals Nonmetals & their Ions
9.2 Metals Nonmetals & their Ions
Transition metal ions can achieve noble gas
Most transition metal ions do not achieve
configuration e.g., Ti4+
noble gas configuration e.g. Fe2+, Fe3+
Ti [Ar]3d24s2 Ti [Ar]3d24s2
Ti2+ + 2e[Ar]3d2 Ti4+ + 4e[Ar]
Fe2+ + 2e-
Fe [Ar]3d64s2
[Ar]3d6 Fe3+ + 3e-
Fe [Ar]3d64s2
[Ar]3d5
Noble gas configuration
9.2 Metals Nonmetals & their Ions Which would you expect to be more stable Fe2+ or Fe3+?
Fe2O3 Fe3+
Answer: Fe3+ Why?
Extra stability associated with a half-filled 3d shell.
9.2 Metals Nonmetals & their Ions Most transition metal ions do not achieve noble gas configuration Cu [Ar]3d104s1
Cu+ + 1e[Ar]3d10 = [Ne]3s23p63d10 Pseudo noble gas configuration 18e- in outer shell
9.3 The Sizes of Atoms & Ions
9-3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions There are size trends for atoms and ions
9.3 The Sizes of Atoms & Ions How do we define atomic, ionic radius? Electron probability cloud…border uncertain We can determine internuclear distances We define radii based on distance between two nuclei
9.3 The Sizes of Atoms & Ions Covalent Radius (for covalent bonded atoms) r
r
O2 (oxygen molecule)
Internuclear distance is 2 x atomic radius (same species)
2r = 143 pm So r = 71.5 pm
9.3 The Sizes of Atoms & Ions Ionic Radius (for ionic bonded species) Na+
Cl-
1893 Windsor, Ontario
9.3 The Sizes of Atoms & Ions Metallic Radius (for metal atoms in crystal) Crystalline Solid Metal
Na+Cl- (ionic compound) Ag
Ag r
rNa+ rCl99pm 181pm
r
rNa+ + rCl- = 280 pm Internuclear dist.
Internuclear distance is sum of the two radii If you know rNa+…you can calculate rCl-
2r
2r = 288 pm So r = 144 pm Internuclear distance is 2 x metallic radius (same species)
9.3 The Sizes of Atoms & Ions The ions in crystalline LiI are arranged as
9.3 The Sizes of Atoms & Ions What are the radii of Li+ and of I-? a
shown below. If a = 600pm, what is the radius of Li+? What is the radius of I-?
r-
Ia=600pm Li+
9.3 The Sizes of Atoms & Ions The ions in crystalline LiI are arranged as shown below. If a = 600pm, what is the radius of Li+? What is the radius of I-? IrI- = 212pm
a=600pm
Li+ + rLi = 88pm
-
r c r-
a r 2r+ a=600pm rrrLi+ = 88 pm
c
a = 2rI- + 2rLi+ 600pm=2*212pm+2rLi+ 2 2 2 2 c = a + a = 2a = 2(600pm)2 (Pythagorean) c = 849 pm = 4rIrI- = 212 pm
8.10 Multielectron Atoms Electron Screening Outer eexperience a lower effective nuclear charge
Inner shell electrons screen outer efrom full attraction of the nucleus
Zeff < Z Zeff = Z - S
where S is charge that is screened
8.10 Multielectron Atoms
9.3 The Sizes of Atoms & Ions
s orbitals are better at screening than p & d Screening s orbitals > p orbitals > d orbitals strength s Zeff > p Zeff > d Zeff (for same n) Orbital energy depends on n and Zeff 2 - Zeff hcR H En = n2
Valence electrons will shield each other to a Much Smaller degree than inner shell electrons
Atomic Radius (pm)
Can we explain the size trends? Atomic Number
9.3 The Sizes of Atoms & Ions
8.12 e- Configurations & The Periodic Table s block
Atomic Radius (pm)
r
r p block
Putting in valence electrons So, S ∼ constant while Z increases Zeff increases and radius decreases
Atomic radius usually decreases from left to right across a period in s & p blocks Atomic Number
9.3 The Sizes of Atoms & Ions
Filling ns orbitals
Filling np orbitals
9.3 The Sizes of Atoms & Ions There are size trends for atoms and ions
Atomic Radius (pm)
Increasing n down a group…orbitals get larger so radius increases r Atomic radius usually increases from top to bottom down a group Atomic Number
9.3 The Sizes of Atoms & Ions
8.12 e- Configurations & The Periodic Table Putting electrons in an inner shell So, S increases as Z increases Zeff ∼ constant and radius ∼ constant d block Transition metals
Atomic Radius (pm)
Atomic radius often stays relatively constant across a period for transition metals Atomic Number
Filling (n-1)d r ∼ const. orbitals
9.3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions
Refer to the periodic table and arrange N, O, and P in order of increasing atomic radius
For main group atoms radius decreases across a period
Predict O < N < P Atomic radius increases down a group
9.3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions
Refer to the periodic table and arrange N, O,
Cations are smaller than the parent atoms
and P in order of increasing atomic radius
Cations…lose electron(s)…Z stays constant
Predicted: O < N < P Actual: 73pm 75pm 110pm
Na Na+
r = 186 pm
9.3 The Sizes of Atoms & Ions
+ 1e-
r = 99 pm
9.3 The Sizes of Atoms & Ions
Anions are larger than the parent atoms
For isoelectronic cations…the more positive
Anion…gain electron(s)…Z stays constant
charge…the smaller the ionic radius
ClCl
+ 1e-
r = 99 pm
r = 181 pm
9.3 The Sizes of Atoms & Ions
9.3 The Sizes of Atoms & Ions
For isoelectronic cations…the more positive
For isoelectronic anions…the more negative
charge…the smaller the ionic radius
charge…the larger the ionic radius Cl- [Ne]3s23p6 P3- [Ne]3s23p6 isoelectronic Z=17 Z=15
Na+ 1s22s22p6 Mg2+ 1s22s22p6 isoelectronic Z=11 Z=12
P3-
ClMg2+
Na+ r = 99 pm
r = 72 pm r = 181 pm
9.3 The Sizes of Atoms & Ions Refer to the periodic table and arrange these species in order of increasing size: Na+,
r = 212 pm
9.3 The Sizes of Atoms & Ions Determine e- configurations from table All species have [Ne] configuration
Mg2+, O2-, F-, Ne Na+Mg2+
O2-F-Ne
For isoelectric configurations radius will increase with decreasing Z
9.3 The Sizes of Atoms & Ions Refer to the periodic table and arrange these species in order of increasing size: Na+, Mg2+, O2-, F-, Ne Answer: Mg2+ < Na+ < Ne < F- < O2Z=12 11 10 9 8 Increasing radius Actual: 72pm 99pm 71pm 133pm 140pm ?
9.3 The Sizes of Atoms & Ions Tutorial will cover other typical questions
9.4 Ionization Energy
9.4 Ionization Energy
9-4 Ionization Energy
Ionization energy for H-like species
Z 2 hcR H ΔE IE = 12
ΔEIE = ? e-
Energy required to strip an electron from a gaseous state atom (or ion)
9.4 Ionization Energy
n=2 H-like species Only one electron Atom absorbs hν photon ΔEIE=hν From ground state n=1
E
9.4 Ionization Energy
Definition Ionization Energy = I
In = nth Ionization energy...remove nth e-
Energy required to strip an electron from a gaseous state atom (or ion) e-
9.4 Ionization Energy
Al+(g) + 1e-
Al(g)
I1=577.6 kJ/mol
Al+(g)
Al2+(g) + 1e- I2=1,817 kJ/mol
Al2+(g)
Al3+(g) + 1e- I3=2,745 kJ/mol
Al3+(g)
Al4+(g) + 1e- I4=11,580 kJ/mol
Not spontaneous…requires energy input
[Ne] noble gas configuration…extra stable
9.4 Ionization Energy
I depends on Zeff and n
Periodic Trends for I I generally increases across a period
2 Zeff hcR H n2
Larger Zeff…higher I…e- held more tightly
I1 kJ/mol
I=
n=∞ n=5 n=4 n=3
I increases as Zeff increases across a period
Higher n…lower I…outer e- held less tightly 2 I ∝ Zeff and I ∝
1 n2
Atomic No. (Z)
8.12 e- Configurations & The Periodic Table s block
9.4 Ionization Energy Periodic Trends for I I generally decreases down a group
I I
p block
Filling ns orbitals
I1 kJ/mol
Putting in valence electrons So, S ∼ constant while Z increases Zeff increases
I decreases as n increases down a group
Filling np orbitals Atomic No. (Z)
9.4 Ionization Energy There are IE trends n I Zeff increases only slightly down a group
9.4 Ionization Energy
There are some exceptions
9.3 The Sizes of Atoms & Ions We would predict I1 Mg < I1 Al
9.4 Ionization Energy
9.4 Ionization Energy Remember that electrons are ionized from
There are some exceptions
the orbital with the highest n I1 = 737.7 kJ/mol Mg
This is not necessarily the last orbital to be Al is lower???
I1 = 577.6 kJ/mol Al Why lower? Mg
If there are different types of orbitals with the
[Ne]3s2
Easier to pull the electron from a higher energy p orbital than a lower energy s orbital
Al [Ne]3s23p1
9.4 Ionization Energy Example Sc Sc
Sc2+
First
[Ar]3d14s2 Sc+
[Ar]3d14s1 [Ar]3d1
same n, the e- is ionized from the highest energy orbital Lower energy s < p < d Higher energy
9.4 Ionization Energy Typical problems…refer to the periodic table
Last e- added
[Ar]3d14s2
Sc+
filled
and arrange elements in expected order of removed
[Ar]3d14s1
Sc2+ Sc3+
e-
[Ar]3d1
[Ar] +
+
+
1e-
increasing first ionization energy I1 Attend tutorials for examples
1e-
1e-
9.5 Electron Affinity
9.5 Electron Affinity
9-5 Electron Affinity
Definition Electron Affinity = EA
ΔEEA = ? e-
Energy change when an electron is added to a gaseous state atom (or ion)
e-
Energy change when an electron is added to a gaseous state atom (or ion)
9.5 Electron Affinity
9.5 Electron Affinity
EA is often negative This means energy is given off
F(g) + 1e1s22s22p5
(exothermic)
F-(g)
EA=-328 kJ/mol
1s22s22p6
Note that ionization energy was always positive…it required input of energy
9.5 Electron Affinity
9.5 Electron Affinity
EAn = nth Electron Affinity...add nth e-
EAn = nth Electron Affinity...add nth e-
O(g) + 1e-
O-(g)
EA1=-141.0 kJ/mol
O(g) + 1e-
O-(g)
EA1=-141.0 kJ/mol
O-(g) + 1e-
O2-(g)
EA2=+744 kJ/mol
O-(g) + 1e-
O2-(g)
EA2=+744 kJ/mol
Positive because it is hard to add a second electron due to repulsion
9.5 Electron Affinity
Electron affinity trends are harder to discern It is easy to get confused! A high affinity for electrons means a large negative electron affinity
9.5 Electron Affinity EA tends to become less negative (lower affinity) down a Group
9.5 Electron Affinity Highest EA values are found in group 17…ns2np5…readily gain one e-
9.7 Periodic Properties of the Elements
9.7 Periodic Properties of the Elements
9-7 Periodic Properties of the Elements
Summary of Atomic Properties Fig.9.11 EA becomes more negative
Arrows indicate increasing
We can understand the trends in terms of Zeff (screening); n (orbital size); e- repulsion
9.7 Periodic Properties of the Elements Periodicity…Predict Properties
9.7 Periodic Properties of the Elements
Physical properties often change uniformly down a group e.g., melting & boiling points
Periodicity…Predict Properties
9.7 Periodic Properties of the Elements Periodicity…Predict Properties
Halogens diatomic e.g., Cl2 Br2 I2
9.7 Periodic Properties of the Elements Periodicity…Predict Properties
Can we predict the melting point of Br2? Estimated mp = (172K+387K)/2 = 280K Actual mp for Br2 = 266K
9.7 Periodic Properties of the Elements Periodicity…Predict Properties Table 9.5
Can we predict the melting point of Br2? Estimated mp = (172K+387K)/2 = 280K Not Too bad! Actual mp for Br2 = 266K
9.7 Periodic Properties of the Elements As you learn reactions you will see trends
9.7 Periodic Properties of the Elements Periodicity…Predict Properties Table 9.5
Try predicting the boiling point of Br2 by averaging Example 9-5 pg. 360
LEO the lion says GER
LEO = loss of electrons Æ Oxidation
9.7 Periodic Properties of the Elements As you learn reactions you will see trends Groups 1 & 2 Metals…Lose ee- loss is Oxidation Group 1 & 2 metals are generally good “reducing agents”
9.7 Periodic Properties of the Elements As you learn reactions you will see trends Group 17 Nonmetals…gain eThey undergo reduction (gain of electrons) They act as oxidizing agents (they are reduced but drive oxidation in another species)
GER = gain of electrons Æ Reduction
9.7 Periodic Properties of the Elements As you learn reactions you will see trends Groups 1 & 2 Metals…lose ee- loss is Oxidation Group 1 & 2 Metals are generally good “reducing agents” A reducing agent undergoes oxidation…but drives reduction (gain of e- by another species)
9.7 Periodic Properties of the Elements In Science, we use models to predict & we test the predictions
Periodic Table is central for understanding Chemistry
The End!
THAT’ THAT’S ALL FOLKS!
GOOD LUCK IN THE REST OF THE CHEM 110 COURSE
