Core electrons
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Core Electrons. These are the inner electrons of an atom . They are not exposed very much to the electrons of other atoms when chemical bonds are formed. Core Electrons.

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Core Electrons

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Core electrons

Core Electrons

These are the inner electrons of an atom. They are not exposed very much to the electrons of other atoms when chemical bonds are formed.


Core electrons1

Core Electrons

These are the inner electrons of an atom. They are not exposed very much to the electrons of other atoms when chemical bonds are formed.

Examples: Sc has the configuration

[Ar]4s23d1


Core electrons2

Core Electrons

These are the inner electrons of an atom. They are not exposed very much to the electrons of other atoms when chemical bonds are formed.

Examples: Sc has the configuration

[Ar]4s23d1

The core is1s22s22p63s23p6 which is represented by [Ar]


Core electrons

The electron configuration of N is

1s22s22px2py2pz


Core electrons

The electron configuration of N is

1s22s22px2py2pz

This can be represented as

[He]2s22px2py2pz


Core electrons

The electron configuration of N is

1s22s22px2py2pz

This can be represented as

[He]2s22px2py2pz

The core is 1s2 in this example.


Valence electrons

Valence Electrons


Valence electrons1

Valence Electrons

The outer shell is called the valence shell, and the electrons in it are called valence electrons.


Valence electrons2

Valence Electrons

The outer shell is called the valence shell, and the electrons in it are called valence electrons.

Examples: K has the electronic configuration [Ar]4s1. The valence electron configuration is 4s1.


Valence electrons3

Valence Electrons

The outer shell is called the valence shell, and the electrons in it are called valence electrons.

Examples: K has the electronic configuration [Ar]4s1. The valence electron configuration is 4s1.

Al has the configuration [Ne]3s23p1. The valence electron configuration is 3s23p1.


Core electrons

The valence electrons are the most important. They are the electrons that most strongly influence the nature and formation of chemical bonds.


Atomic radii and ionic radii

Atomic Radii and Ionic Radii


Atomic radii and ionic radii1

Atomic Radii and Ionic Radii

One way to estimate the size of an atom is to measure the atomic radius, defined to be half the distance between two identical atoms in a molecule.


Atomic radii and ionic radii2

Atomic Radii and Ionic Radii

One way to estimate the size of an atom is to measure the atomic radius, defined to be half the distance between two identical atoms in a molecule.

For example, the distance between two identical atoms (measured from the nuclei) in an I2 molecule is 2.66 Å. The radius of each iodine atom is taken to be 1.33 Å.


Core electrons

The atomic radius decreases as we move across a period(there are some minor exceptions to this). The increase in the nuclear charge as we move across a period results in a shrinkage of the atomic radius – that is, the charge cloud is more strongly attracted to the nucleus.


Core electrons

As we go down the periodic table, the atomic radius increases with increasing atomic number. The orbital size increases with increasing principal quantum number n.


Core electrons

As we go down the periodic table, the atomic radius increases with increasing atomic number. The orbital size increases with increasing principal quantum number n. The corresponding increase in the nuclear charge does not decide the issue of size as we go down the periodic table.


Core electrons

As we go down the periodic table, the atomic radius increases with increasing atomic number. The orbital size increases with increasing principal quantum number n. The corresponding increase in the nuclear charge does not decide the issue of size as we go down the periodic table.

The size of atoms plays an important role in the nature of chemical bonds.


Ionic radius

Ionic Radius


Ionic radius1

Ionic Radius

Ionic radius: The radius of a cation or an anion as measured in an ionic compound.


Ionic radius2

Ionic Radius

Ionic radius: The radius of a cation or an anion as measured in an ionic compound.

Anions (single atom ones) have larger radii than cations in the same period.


Core electrons

Consider the isoelectronic species:

Na+, Mg2+, Al3+


Core electrons

Consider the isoelectronic species:

Na+, Mg2+, Al3+

The smallest radii occurs for Al3+.

The next smallest is Mg2+, followed by Na+.


Core electrons

Consider the isoelectronic species:

Na+, Mg2+, Al3+

The smallest radii occurs for Al3+.

The next smallest is Mg2+, followed by Na+.

In the trivalent cation, the electron density is pulled inward towards the nucleus most strongly by the +3 charge on the nucleus.


Core electrons

For the anions O2- and F-, the oxide ion is the larger, because the extra electrostatic repulsion between the electrons in O2- will spread out the electron density to a greater extent than in F-.


Core electrons

For the anions O2- and F-, the oxide ion is the larger, because the extra electrostatic repulsion between the electrons in O2- will spread out the electron density to a greater extent than in F-. The N3- ion is larger than the oxide ion.


The periodic table

The Periodic Table


The periodic table1

The Periodic Table

Some trends


Core electrons

Definition: The ionization energy is the energy required to remove an electron from one mole of a substance in its ground state in the gas phase. For example, for substance X,

X(g) X(g)+ + e-

This defines the first ionization energy.


Core electrons

The smaller the ionization energy, the easier a cation may be formed.


Core electrons

The smaller the ionization energy, theeasier a cation may be formed.

The ionization energy gives information about the chemical reactivity of an element.


Core electrons

Element First ionization energy (kJ/mol)

Li 520

increasing reactivity


Core electrons

Element First ionization energy (kJ/mol)

Li 520

Na 496

increasing reactivity


Core electrons

Element First ionization energy (kJ/mol)

Li 520

Na 496

K 419

increasing reactivity


Core electrons

Element First ionization energy (kJ/mol)

Li 520

Na 496

K 419

Rb 403

increasing reactivity


Core electrons

Element First ionization energy (kJ/mol)

Li 520

Na 496

K 419

Rb 403

Cs 376

increasing reactivity


Core electrons

Element 1st IE

Li 520

Na 496

K 419

(units are kJ/mol)


Core electrons

Element 1st IE 2nd IE

Li 520 7300

Na 496 4560

K 419 3052

(units are kJ/mol)


Core electrons

Element 1st IE 2nd IE 3rd IE

Li 520 7300 11808

Na 496 4560 6900

K 419 3052 4410

(units are kJ/mol)


Core electrons

Element 1st IE 2nd IE 3rd IE

Li 520 7300 11808

Na 496 4560 6900

K 419 3052 4410

(units are kJ/mol)

From this table it is clear why we do not have Li2+, Na2+, K2+, or Li3+, Na3+, K3+ as common cations.


Core electrons

Element 1st IE(kJ/mol)

He 2370

Ne 2080

Ar 1520

Kr 1350

Xe 1170


Core electrons

Element 1st IE

Be 899

Mg 738

Ca 590

(units are kJ/mol)


Core electrons

Element 1st IE 2nd IE

Be 899 1757

Mg 738 1450

Ca 590 1145

(units are kJ/mol)


Core electrons

Element 1st IE 2nd IE 3rd IE

Be 899 1757 14850

Mg 738 1450 7730

Ca 590 1145 4900

(units are kJ/mol)


Core electrons

Element 1st IE 2nd IE 3rd IE

Be 899 1757 14850

Mg 738 1450 7730

Ca 590 1145 4900

(units are kJ/mol)

From this table it is clear why we do not have Be3+, Mg3+, Ca3+ as common cations.


Electron affinity

Electron Affinity


Electron affinity1

Electron Affinity

Electron affinity: The energy released (usually) when one mole of electrons are added to one mole of a species in its ground state and in the gas phase.


Electron affinity2

Electron Affinity

Electron affinity: The energy released (usually) when one mole of electrons are added to one mole of a species in its ground state and in the gas phase.

For species Y, we have

Y(g) + e- Y-(g)


Core electrons

The higher the electron affinity, the more likely the formation of an anion will occur.


Core electrons

Element Electron Affinity (kJ/mol)*

Cl 350

F 338

Br 330

I 300

S 164

O 145

H 77

Li 58

* Usually reported as positive values. Using proper convention, they are negative if energy is released.


Some general trends

Some general trends


Ionic and covalent compounds

Ionic and Covalent compounds

Review naming of inorganic compounds.


Chemical bonding

Chemical Bonding


Chemical bonding1

Chemical Bonding

Three main types:


Chemical bonding2

Chemical Bonding

Three main types:

1. Covalent Bonding


Chemical bonding3

Chemical Bonding

Three main types:

1. Covalent Bonding

2. Ionic Bonding


Chemical bonding4

Chemical Bonding

Three main types:

1. Covalent Bonding

2. Ionic Bonding

3. Metallic Bonding


Chemical bonding5

Chemical Bonding

Three main types:

1. Covalent Bonding

2. Ionic Bonding

3. Metallic Bonding

The focus will be on the first two.


Covalent bonding

Covalent Bonding


Covalent bonding1

Covalent Bonding

An early view of the covalent bond.


Covalent bonding2

Covalent Bonding

An early view of the covalent bond.

Covalent bond: A bond in which two electrons are shared by two atoms.


Covalent bonding3

Covalent Bonding

An early view of the covalent bond.

Covalent bond: A bond in which two electrons are shared by two atoms.

The key breakthroughs were made by Gilbert Lewis.


Lewis structures

Lewis Structures


Lewis structures1

Lewis Structures

Lewis structure: A diagram showing how electron pairs are shared between atoms in a molecule.


Lewis structures2

Lewis Structures

Lewis structure: A diagram showing how electron pairs are shared between atoms in a molecule.

Consider the dihydrogen molecule.

H. + .H


Lewis structures3

Lewis Structures

Lewis structure: A diagram showing how electron pairs are shared between atoms in a molecule.

Consider the dihydrogen molecule.

H. + .H

The electron pair will be replaced with a line to designate the bond, so will be replaced by

H H


Lewis structures4

Lewis Structures

Lewis structure: A diagram showing how electron pairs are shared between atoms in a molecule.

Consider the dihydrogen molecule.

H. + .H

The electron pair will be replaced with a line to designate the bond, so will be replaced by

H H

This is the Lewis structure for the dihydrogen molecule.


Covalent bond formation for h 2

Covalent bond formation for H2


Distribution of electron density in h 2

Distribution of electron density in H2


Core electrons

Only the valence electrons are represented in Lewis structures.


Core electrons

Only the valence electrons are represented in Lewis structures.

The difluorine molecule would be represented as


Core electrons

Only the valence electrons are represented in Lewis structures.

The difluorine molecule would be represented as

and with the covalent bond shown, as


Core electrons

The Lewis structure for the water molecule is

H HHH


Core electrons

The Lewis structure for the water molecule is

H HHH

The electrons that are not involved in bond formation are called nonbonding electrons or lone pairs.


Core electrons

The Lewis structure for the water molecule is

H HHH

The electrons that are not involved in bond formation are called nonbonding electrons or lone pairs.

In F2 each F atom has three lone pairs and the O atom in water has two lone pairs.


Core electrons

There are many compounds in which the same two atoms share two or even three pairs of bonding electrons.


Core electrons

There are many compounds in which the same two atoms share two or even three pairs of bonding electrons. In these cases we have multiple bonding.


Core electrons

There are many compounds in which the same two atoms share two or even three pairs of bonding electrons. In these cases we have multiple bonding.

The bond between two atoms sharing twopairs of electrons is called a double bond.


Core electrons

There are many compounds in which the same two atoms share two or even three pairs of bonding electrons. In these cases we have multiple bonding.

The bond between two atoms sharing twopairs of electrons is called a double bond. Examples are O2 and CO2.


Core electrons

There are many compounds in which the same two atoms share two or even three pairs of bonding electrons. In these cases we have multiple bonding.

The bond between two atoms sharing twopairs of electrons is called a double bond. Examples are O2 and CO2.


Core electrons

There are many compounds in which the same two atoms share two or even three pairs of bonding electrons. In these cases we have multiple bonding.

The bond between two atoms sharing twopairs of electrons is called a double bond. Examples are O2 and CO2.

C


Core electrons

Note that only the 2s and 2p electrons (valence electrons) are used in forming these bonds in O2 and CO2.


Core electrons

Note that only the 2s and 2p electrons (valence electrons) are used in forming these bonds in O2 and CO2.

A triple bond arises when two atoms share three pairs of electrons.


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