Section 8.4 Ions: Electron Configurations and Sizes Return to TOC Electron Configurations in Stable Compounds When two nonmetals react to form a covalent

Section 8.4

Ions: Electron Configurations and Sizes

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Electron Configurations in Stable Compounds

• When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence electron configurations of both atoms.

• When a nonmetal and a representative-group metal react to form a binary ionic compound, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom. The valence orbitals of the metal are emptied.

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Ions

An ion is an atom with a chargeCation – positively charged atomAnion – negatively charged atom

The question becomes….why do atoms form ions?

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Ions

Atoms will gain or lose e- in an attempt to form the same electron configuration as the closest noble gas.

***Move the LEAST number of e- possible.

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Ions

• Atoms in groups 1 and 2 will lose the outer valence e- first. These are the “s” e-

• Atoms in groups 13-15, below the “stairs” will lose their outer “p” e- first, then their outer “s” e-

• Exceptions: B, Al tend to lose both the “p” and “s” e- at the same time

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Transition Metals – lose their valence “s” e- first, then they may lose another e- from the “d” sublevel.

Exceptions: Ag, Zn, Cd – they do NOT lose e- from the “d” sublevel

Why?After they lose their “s” electrons, it takes too

much energy to take from the full “d” sublevel

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Rules for Filling Orbitals– Any orbital may contain 0, 1, or, at most, 2

electrons.– In filling the p, d, and f subsets, each orbital gets a

single electron with the same spin as the others before any pairing takes place.

– This is because more energy would be required to fill them in any other way.

ELECTRON CONFIGURATIONS

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Elements with atomic numbers 14 have only s electronsElements with atomic numbers 510 also have electrons in p orbitals

Elements 2130 have d electronsElements 5871 have electrons in f orbitals along with all their other electrons.

Figure 3.18, pg. 78

Investigating Chemistry, 2nd Edition

© 2009 W.H. Freeman & Company


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Beryllium has an atomic number of 4, with two 1s electrons and with two electrons in the 2s orbital. Adding the superscripts gives the total number of

electrons.


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Note the exceptions in red. Copper, Cu, also has an unexpected configuration.

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• It was once suspected that the deposed Emperor Napoleon was poisoned with arsenic. What is the electron configuration of arsenic, As, element number 33?

• Following the periodic table from H, to He, to Li, Be, B, C, N, O, F, etc., – We get 1s2, 2s2, 2p6, 3s2, 3p6…– So far we have 2 + 2 + 6 + 2 + 6 = 18 e’s.

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• 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p3 • Let’s check our math.• 18 + 2 + 10 + 3 = 33, the right number of electrons in

a neutral arsenic atom, As.• Since we followed the periodic table, we did not have

to memorize the fact that the 4s orbital is filled before the 3d orbitals.

• The set of three 4p orbitals is only half-filled.


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• Because the elements N and P are directly above arsenic, As, in the periodic table, they also have half-filled p subshells.

• As a result, these three elements have many chemical similarities.

• Now we can begin to see why Mendeleev was able to predict the properties of elements and compounds that had not yet been discovered in 1869.


Section 8.4


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Electron Orbital Configurations

The configuration may be written by using boxes to represent each orbital

All orbitals MUST be in increasing energy and MUST contain a label

1s 2s 2p

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1s 2s 2p

Arrows are used to represent each electron

Before we begin…..

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Three Rules

Aufbau’s Principle – lower energy orbitals fill before proceeding to higher energy orbitals

Hund’s Rule – When there are multiple orbitals available in a sublevel, one electron is placed in each orbital before doubling up the electrons

Pauli’s Exclusion Principle – Within each orbital, e- must spin in opposite directions; each orbital in a sublevel must spin in the same direction.

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Aufbau’s Principle

To get the orbitals in increasing energy, just follow the periodic table like you would read a book.

1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p

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Hund’s Rule

Never double up electrons in an orbital until each orbital in that sublevel has one electron.

Once each orbital in a sublevel has one electron, then begin to double up the electrons.

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Pauli’s Exclusions Principle

Electrons will take the lowest energy configuration possible. This means:

1. All unpaired e- must spin in the same direction.

2. All paired e- must spin in opposite directions

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1s 2s 2p

Hydrogen Atomic # =1 , 1e-

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1s 2s 2p

Helium – Atomic Number = 2

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1s 2s 2p

Boron – Atomic Number = 5

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Stable Compounds

• Atoms in stable compounds usually have a noble gas electron configuration.

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Noble Gas Configuration

What is a noble gas?

Noble gases are located in group 8A, 18 on the periodic table.

Noble gases are extremely unreactive, because their outer energy level is filled

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Noble Gas Configurations

Noble gases include:

He Ne Ar Kr Xe Rn

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Aufbau tells us that all lower sublevels MUST be filled before filling sublevels of higher energy.

This results in us writing the same information repeatedly when making short hand configurations:

Mn 1s22s22p63s23p64s23d5

Cl 1s22s22p63s23p5

Ca 1s22s22p63s23p64s2

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Rules: Choose the largest noble gas that has an atomic number LESS than the element you are working with.

For Mn, the largest noble gas is Ar

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Because we know that lower sublevels are already filled, we can substitute part of the configuration with a noble gas:

Mn 1s22s22p63s23p64s23d5

Ar 1s22s22p63s23p6

Therefore we write: [Ar] 4s23d5

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Now try it for Cl

Cl 1s22s22p63s23p5

The largest noble gas is Ne 1s22s22p6

[Ne] 3s23p5

Valence electrons – these ARE used in bonding

Core electrons – these are NOT used when bonding

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Write the noble gas configurations for:

As

I

Pb

Au

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Write the noble gas configurations for:

As [Ar] 4s23d104p3

I [Kr] 5s24d105p5

Pb [Xe] 6s24f145d106p2

Au [Xe] 6s24f145d9

W [Xe] 6s24f145d4

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Exceptional Configurations

….and ions

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Exceptions to Aufbau

There is a general stability associated with electron configurations

Filled sublevels are MOST stable ½ Filled sublevels are stable All other configurations for sublevels are

LEAST stable

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Sometimes by moving electrons between sublevels that are close in energy, atoms can achieve a more stable configuration.

Examples include:

s2d4

Because d5 is ½ filled and more stable, the atom takes on the configuration of

s1d5

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Cr [Ar] 4s13d5

Mo [Kr] 5s14d5

W [Xe] 6s14f145d5

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Another exception occurs with the configuration:

s2d9

Again, by moving 1e- from the “s” sublevel to the “d” sublevel, the “d” sublevel becomes filled.

s1d10

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Cu [Ar]4s13d10

Ag [Kr]5s14d10

Au [Xe]6s14f145d10

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WARNING

Exceptional configurations only happen between “s” and “d” sublevels….NEVER between “s” and “p” sublevels.

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Documents

Section 8.4 Ions: Electron Configurations and Sizes Return to TOC Electron Configurations in Stable Compounds When two nonmetals react to form a covalent