Chemistry: Lecture 9

Multi-electron atom: energy of orbitals

Energy level also depends on L, not just n: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p- energy of orbitals determines which orbitals get filled up first

Ionization energy

Energy required to remove one single electron in the gas state - always positive - Na(g)= Na+(g) + 1e - Na+(g)= Na2+(g) +1e (second IE) trends: - as the n gets larger, electrons experience less nuclear charge, meaning less ionization energy is required to remove electron, so down a group, there is a decrease in ionization energy - across a period, the ionization energy increases because electrons on the outside experience greater Zeff as you move across

How to do an orbital diagram?

First find number of electrons, then draw the boxes, the number of boxes for each subshell depends on the Ml number. Then draw in the arrows, make sure you follow hunds rule and do not fill in arrows in one box if another box is free

Penetration

If an electron can get closer to the nucleus than the 1s electrons, then it will experience full nuclear charge - as the outer electron undergoes penetration into the inner electron region, it experiences greater and greater nuclear charge

Exceptions to ionization energy

- boron less than Be- jump from 2s to 2p means that more shielding, and less energy required to take away - oxygen is lower than nitrogen because there is repulsion between the filled boxes of oxygen, making it lower in ionization energy

Ionic radius

- cations are smaller - anions are larger because there is an extra electron that is increasing the shielding to the other electrons, which means the nuclear charge decreases which means the electron is larger

Coulombs law

- describes the relationship between electron and the nucleus - an atom wants low potential energy, so if two particles are the same charge they will repel because the bigger the r, the lower the E: E is positive - for two different charges, the lower the r, the lower the potential energy which is why the two particles attract each other: E is negative

Metallic character

- how likely are electrons to be lost/gained: remember that metals are more likely to lose and nonmetals are more likely to gain decreases across a period and increases down a group

Why do second/third/fourth ionization energies differ so much?

- there is a huge jump when removing an electron from the atom's core

Two types of shielding

1. Core electrons shielding from nuclear charge 2. Valence electrons shielding each other - the core electrons are the ONLY ones relevant to the effective nuclear charge

Electron configuration of ion

1. Start with normal electron configuration, then add or remove electrons - anion: add extra electrons - cation: lose electrons - for the same n values: take away electrons from p subshell before s subshell - for different n values, take away electrons from the higher n value first

Degenerate orbitals: multi-electron species

2l+1= orbital states with same energy this is because there are effects of penetration and shielding in multi-electron orbitals

2p and 2s orbitals

2p and 2s are degenerate if there4 is no 1s, but if there is, 2s experiences a stronger effective nuclear charge and therefore, has a lower energy than 2p

2s^1

2s is sub shell, 2=n, l=0, 1=number of electrons in the sub shell

Metalloids

7 metalloids - semiconductors, semi metals

Shielding

Electrons in higher n values do not feel the full effects of the nucleus because of other electrons getting in their way - as orbitals increase, shielding increases: the less and less nuclear energy felt by electrons

How to calculate effective nuclear charge?

Actual Nuclear charge- amount of core electrons - so for lithium with 2 core electrons and a nuclear charge of 3, the Zeff is +1

Group 1A

Alkali metals- very reactive, violently lose one electron to form ions

Group 2A

Alkaline earth metals, tend to lose two electrons, not as reactive as alkali metals 2+

Dimagnetic

All electrons are paired, so they are not attracted to an external magnetic field

Paramagnetic

An unpaired electron has a spin, this spin makes a tiny magnetic field, so any electron with unpaired electrons is attracted to an external magnetic field

Nuclear charge and potential energy relationship

As you experience greater nuclear charge, you also experience a lower energy:

Why is it more energy to take away an electron in He+ than H?

Because the electron in He+ is experiencing a 2+ pull from the nucleus, while the H electron is only experiencing a 1+ charge from nucleus Because of this, He+ has a smaller radius than H

Pauli exclusion principle

Electrons in the same orbital have to have different spins (Ms quantum number)

1s subshell

Can only hold two electrons maximum - n=1, l=0 remember: each subshell holds two electrons, the amount of boxes for the orbital diagram depends on Ml- the individual orbitals s has 1 p has 3 D has 5 f has 7

Why do orbital energies look different in multi electron atom vs single?

Coulomb's law gives potential energy between electrons -positive E means there is repulsion, which increases energy of a system and causes destability -negative E means there is attraction, which decreases the energy of a system, making it more stable- this is more energetically favorable - electrons in multi electron species do not experience the full nuclear charge: electron feels a nuclear charge less than the nucleus, this is called the EFFECTIVE nuclear charge is less than 2 - as orbitals increase, the energy of the orbitals also goes up

Atomic radius

Distance between non bonding atoms that are in direct contact

Electron affinity trend

Electron affinities become more positive as we go down the 1A groupmet, and become more negative from left to right - higher Zeff means more negative electron affinity

Why does the grouping of atoms in metals, non metals, and metalloids show their properties

Electron configurations - 1A, 2A, 3A, etc. : show valence electrons

Orbital penetration

For a given n, the s gets closest to nucleus, p second closest, d third closest, and f the furthest. - in other words, s penetrates further: the more the electrons penetrate, the closer to the nucleus you get and the lower the energy of atom - less orbitals, more penetration= electrons on lower orbitals feel a larger nuclear charge

Valence electrons

For main group elements: the electrons in the outermost n level - for example: for Ge there are 4 valence electrons For transition elements: it is higher n electrons and outermost d shell too - 1A, 2A: group number gives charge of cations - 5A, 6A, 7A, ????

Covalent/bonding radius

Half the distance between two bonded atoms - nonmetals: half the distance between two bonded atoms - metals: half distance between two atoms next to each other - by the way, the approximate bond length between two covalently bonded atoms is the sum of their atomic radius

Group 7A

Halogens: gain one electron, 1-

Core elections

Inner electrons- all the electrons that are not valence- they do not bond, but are responsible for shielding effects

Noble-gas-core abbreviation

Instead of writing 1s^2, 2s^1 for lithium, you write [He] 2s^1 -use the noble gas you find after the element you want - this means that the core electrons are isoelectronic with a noble gas

How to do electron configurations?

Look at atomic number (the top left corner), that is the number of electrons that you need to incorporate. Then, write out the numbers, be aware that the lowest energy ones fill out first: fill them out until there are no electrons left

Electron affinity

Measure of how easily an atom accepts an additional electron - usually negative because atom releases energy when it gains electron Cl(g) + 1 e =Cl-(g)

Metal vs nonmetal

Metals: not H— tends to lose electrons, very reactive nonmetals- H and everything on the right, gain electrons easily

Valence electrons

Most important for bonding: - for transition elements: count both greatest n AND outermost d electrons - for main group: electrons in highest n level - electrons in a group have the same number of valence electrons

Noble gases (8A)

Most unreactive elements: inert, these atoms are very stable= the potential energy is very low - remember, systems with low potential energies try to stay that way and things with higher potential energies try to become lower - full octets

Shielding

Multi electron atoms: every electron experiences both the positive charge of the nucleus and the negative charge of the electrons: it is attracted to nucleus and repulsed by the other electrons. The repulsion of the electrons is known as screening or shielding: the electron will not feel the full effects of the nuclear charge - so, the effective nuclear charge is the nucleus minus the charge of the electrons

Van dear waals radius

Radius of the two atoms next to each other, multiplied by two

Transition elements: atomic radius across a period

Radius stays the same because the number of valence electrons stays roughly the same

Excited state

Some electrons are in higher energy orbitals

Atomic radius: transition elements

Stays roughly the same, this is because the amount of valence electrons stays roughly the same

Coulomb's law- general idea

The closer the electrons to the nucleus, the lower energy they will have, and lead to a more stable system

Density pattern

The density of an object tends to increase as it goes down a group, this is because there is a greater increase in the mass than the volume

Hydrogen atom: energy of orbitals

The energy of an orbital only depends on the value of n- for example, 2s and 2p orbitals are the same energy and 2p has a greater energy than 1s

Aufbau principle

The idea that the lowest energy levels get filled first before the higher number subshells

Degenerate

Two or more electron orbitals with the same value of n and have the same energy: for single electron atoms, since energy levels of orbitals only depends on n, all orbitals with the same value of n are degenerate - the orbitals within multi-electron atoms are not degenerate, they are splitting

Cations and anions

We know that cations is supposed to be a loss of electrons and anions is supposed to be a gain of electrons - for main group cations, you have to remove the electrons from the reverse order of filling: the first spot available you remove it from there - for transition elements, you have to remove the electrons from the highest n orbitals first

Ground state atom

When all electrons are in the lowest energy orbitals possible

Hund's rule

When orbitals of identical energy are available, electrons will occupy them first before going into the same orbital - this is because if it is alone in an orbital, the electron can maximize its spin- the best configuration is one that maximizes spin

Zeff across a period

Zeff keeps getting stronger as the core electrons stay the same, but the valence electrons grow: so, as effective nuclear charge grows, the radius gets smaller and smaller

Copper (Cu) electron configuration

[Ar] 4s^1 3d^10

Chromium (Cr) electron configuration

[Ar] 4s^1 3d^5

Silver (Ag) electron configuration

[Kr] 5s^1 4d^10

Molybdenum (Mo) electron configuration

[Kr] 5s^1 4d^5

Exceptions to the Aufbau Principle

chromium, molybdenum, silver, copper

Degenerate orbitals: single electron species (H)

orbitals that have the same energy n^2 is the orbitals with the same energy - for example: in n=2 shell, there are 4 same energy orbitals (one 2s orbital, 3 2p orbitals) - hydrogen has a degeneracy of 4 - depends on the idea that energy level is only dependent on the n number