Chemistry-Ch04_Atomic Energy Levels

8/9/2019 Chemistry-Ch04_Atomic Energy Levels

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44AtomicAtomic

energy levelsenergy levels



4.1 Characteristics of atoms4.1 Characteristics of atoms

Atoms:

± Possess mass

± Contain positive nuclei ± Contain electrons

± Occupy volume

± Have various properties

± Attract one another

± Can combine with one another



4.2 Characteristics of light4.2 Characteristics of light

The most useful tool for studying the

structure of atoms is electromagnetic

radiation

Light is one form of this radiation




Wave properties of light ± Light has wave-like properties

± A wave is a regular oscillation in some

particular property ± Wavelength (P) - the distance between two

successive crests (units: meters or nanometers)

± Frequency (R) - the number of waves passing a

certain point over a unit of time (units: s-1 = Hz) ± Amplitude: height of the wave measured from

the axis of propagation, a measure of intensity



4.24.2

CharacteristicsCharacteristicsof lightof light



± Phase refers to the starting position of a

wave with respect to one wavelength




± The wavelengths and frequencies of

electromagnetic radiation cover an immense

range

± What we perceive as white light actually

consists of a range of wavelengths







Particle properties of light

± Light carries energy

± The photoelectric effect shows how theenergy of light depends on its frequency and

intensity





± Several phenomena can be observed:

Below a certain frequency, no electrons were

observed, no matter what the intensity

The energy of the ejected electronsincreased linearly with the frequency of light

(E w R)

The number of emitted electrons increased

with light intensity All metals show the same pattern, but each

metal has a different threshold frequency




± Albert Einstein postulated that light

comes in µpackets¶ or µbundles¶

± These are called photons

± Each photon has an energy that is

directly proportional to its frequency

± Ephoton = hRphoton

± The proportionality constant betweenenergy and frequency is known as

Planck¶s constant = 6.626 x 10-34 J s




± Electrons are only ejected if the

frequency of light is high enough

If the photon does not have enough energy,

the electrons are not ejected.

± It is not dependent on the intensity of the

light

The greater the intensity of light, the more

photons, but if no individual photon hasenough energy (the correct frequency) to

remove an electron, then none are ejected.






Absorption and emission spectra

± Light and atoms

Energy must be supplied

to remove a bound

electron from an atom





Ground state: lowest energy state of an

atom.

Excited state: when an atom absorbs a

photon Energy level diagram: depicts the

changes in energy of an atom

When an atom emits a photon (or

radiates heat), it returns to the groundstate.

¨E=± hphoton



Atomic spectra

± Atoms absorb specific and characteristic

frequencies of light

± The absorption spectrum of a gaseous

element can be measured




± A similar experiment can measure the energies

of the photons emitted by atoms in excited

states

± This gives an emission spectrum




± The characteristic patterns of energy gains and

losses provide information about atomic

structure





Quantisation of energy

± A photon with high enough energy can

cause an atom to lose on of its electrons

± This implies that absorption of a photon

results in an energy gain for an electron

in the atom

± Energy change for an atom equals theenergy change for an atomic electron

± Eatom = Eelectron = h



± Balmer was able to propose a single equation

that could describe the emission spectrum of

the hydrogen atom

± Bohr used the discovery that E = hv to

interpret Balmer¶s observations

n

1

n

1s103.29

2

2

2

1

115

¹¹ º

¸©©ª

¨v!




± He combined the two equations to form one

which could describe the energies

corresponding to the emission lines in the

spectrum of hydrogen

± A property that is restricted to specific values

is said to be quantised

2

1810182

n

J E

n

v

!




Energy level diagrams

± Atomic energy

transformations can

be represented

using an energy

level diagram




± A ball on a staircase shows some properties of

quantised energy states




4.3 Properties of electrons4.3 Properties of electrons

Electrons:

± All have the same mass and charge

± Have magnetic properties ± Display wave-particle duality

mu

h particle !P

Where m is mass and u is velocity

Suggests that every moving thing has a

wavelength associated with it







The Heisenberg uncertainty principle

± Because of their wave properties,

electrons are always spread out rather

than located in one particular place

± The position of a moving electron cannot

be precisely defined

± Electrons are described as delocalisedas their waves are spread out rather than

pinpointed




± Mathematically, the position and

momentum of a particle-wave are linked

± The Heisenberg Uncertainty Principle

states that the more accurately we knowposition, the more uncertain we are

about motion, and vice versa

± So instead, we identify a µprobable

location¶ of the electrons in an atom, not

an exact one



4.4 Quantisation and4.4 Quantisation andquantum numbersquantum numbers

The cloud of electrons that orbit thenucleus are referred to as orbitals

Each quantised property can beidentified, or indexed, using aquantum number

Each electron in an atom has three

quantum numbers that specify itsthree variable properties

A fourth number describes the spin




Principal quantum number

± Indexes energy for any atom containing

only a single electron

± The principal quantum number must be a

positive integer

± Is correlated with orbital size

± As n increases, the energy of theelectron increases, its orbital gets bigger

and it is less tightly bound to the atom




Azimuthal quantum number

± Indexes the angular momentum of the

orbital

± The value correlates with the number of

preferred axes in a particular orbital

± It thereby identifies the shape of the

electron distribution within the orbital ± Can be zero or any positive integer

smaller than n




Magnetic quantum number

± Objects with preferred axes have directionality

as well as shape

± The magnetic quantum number indexes the

restricted numbers of possible orientations



Spin quantum number

± All electrons have a property called spin

± This means that they can behave in one of two

ways in a magnetic field

± The spin quantum number indexes this

behaviour




The Pauli exclusion principle

± States that each electron in an atom has a

unique set of quantum numbers, which must

meet all the restrictions summarised below

A direct consequence is that any orbital can contain

a maximum of two electrons




4.5 Atomic orbital electron4.5 Atomic orbital electrondistributions and energiesdistributions and energies

The chemical properties of atoms aredetermined by the behaviour of their electrons

Because atomic electrons aredescribed by orbitals, the interactionsof electrons can be described in termsof orbital interactions

± There are two characteristics of orbitalsthat determine how electrons interact




Orbital electron distributions

± Distribution of electrons can be

described using electron density

± Orbitals describe the delocalisation of

electrons

± An atom that contains many electrons

can be described by superimposing theorbitals for all of its electrons to obtain

the overall size and µshape¶ of the atom



4.5 Atomic orbital4.5 Atomic orbitalelectronelectron

distributions anddistributions andenergiesenergies




± Orbital depictions

Orbital pictures provide maps of how an

electron wave is distributed in space

An electron density plot represents theelectron distribution in an orbital as a two-

dimensional graph

Electron density pictures can indicate the

three-dimensional nature of orbitals An electron contour drawing provides a

simplified orbital picture



± Orbital size

Experiments that measure the atomic radii provide

information about the size of orbitals

Orbitals get larger as the value of n increases




In any particular atom, all orbitals with the same

principal quantum number are similar in size




± Details of orbital electron distributions

The electron distributions of orbitals strongly

influence chemical interactions

The quantum number l = 0 corresponds to ans orbital

There is only one s orbital for each value of the

principal quantum number




The quantum number l = 1 corresponds to a

p orbital

For each value of n there are three different

p orbitals




The quantum number l = 2 corresponds to a

d orbital

There are five different d orbitals




Orbital energies

± All orbitals

correspondingto one level of nhave the sameenergy and are

said to bedegenerate




± The effect of nuclear charge





The energy of an orbital can be determined

by measuring the amount of energy required

to remove and electron completely

This is the ionisation energy (Ei)

H H+ + e- EiH = 2.18 x 10-18 J

He+

He2+

(g) + e-

EiHe+ = 8.72 x 10-18

J



± The effect of other electrons

In a multi-electron atom, each electron affects the

properties of all other electrons

A negatively charged electron in a multi-electronatom is attracted to the positively charged nucleus,

but it is repelled by the other negatively charged

electrons




± Shielding

Electron-electron

repulsion cancels aportion of the

attraction between the

nucleus and the

incoming electron

This partialcancellation is called

shielding





The effectiveness in shielding nuclear

charge decreases as orbital size increases

An electron¶s ability to shield decreases as n

increases

In multi-electron atoms, electrons with any

given value of n provide effective shielding

for any orbital with a larger value of n

The higher the value of the l quantum

number, the more that orbital is shielded by

electrons in smaller, lower-energy orbitals



4.6 Structure of the periodic4.6 Structure of the periodictabletable

The periodic table lists the elements in

order of increasing atomic number

This list is also in order of increasingnumber of atomic electrons

The elements are also placed in rows

called periods such that the columns

are formed with groups of elements

that have similar chemical properties




The aufbau principle and order of

orbital filling

± The ground state of an atom is the most

stable arrangement of its electrons

± We construct the ground-state

configuration by placing electrons in the

orbitals starting with the lowest in energyand moving progressively upward

± This is the aufbau principle




1. Each electron in an atom occupies themost stable available orbital. (Aufbau)

2. No two electrons can have identicalquantum numbers. (Pauli)

3. Orbital capacities are as follows:s: 2 electrons, p set: 6 electrons

d set: 10 electrons, f set: 14 electrons

4. The higher the values of n, the less stable

the orbital.5. For equal n, the higher the value of l , the

less stable the orbital






4.6 Structure of the periodic4.6 Structure of the periodic

tabletable



4.6 Structure of the periodic4.6 Structure of the periodic

tabletable




Valence electrons

± The chemical behaviour of an atom isdetermined by the electrons that are

accessible to an approaching chemical ± Accessible electrons are called valence

electrons

± Inaccessible electrons are called core

electrons ± Valence electrons participate in chemical

reactions, core electrons do not



4.7 Electron configurations4.7 Electron configurations

An electron configuration is a complete

specification of how an atom¶s electrons are

distributed in its orbitals

There are three common ways to representelectron configurations:

± Complete specification of quantum numbers

± Shorthand notation from which quantum numbers

can be inferred

± Diagrammatic representation of orbital energy

levels and their occupancy




n = 1, l = 0, ml = 0, m

s= + ½ 1s1

n = 1, l = 0, ml = 0, m s = + ½ 1s2

n = 1, l = 0, ml = 0, m

s= - ½

n = 1, l = 0, ml = 0, m s = + ½ 1s22s1

n = 1, l = 0, ml = 0, m

s= - ½

n = 2, l = 0, ml = 0, m

s= + ½



Electron configurations become longer as

the number of electrons increases

The element at the end of each period has

a noble gas configuration





Electron-electron repulsion

± For two of more degenerate orbitals, the

lowest energy situation results when

electrons occupy the orbitals that keepthem furthest apart

± Hund¶s rule states that the lowest energy

configuration involving orbitals of equal

energies is the one with the maximum

number of electrons in the same spin

orientation



Orbitals with nearly equal energies

± Some elements have ground-state

configurations different from the predictions of

the regular progression





Configurations of ions

± The electron configurations of atomic

ions are written using the same

procedure as for neutral atoms, takinginto account the proper no. of electrons

± Isoelectronic atoms and ions have the

same number of electrons




Guidelines for building atomic or ionic

electron configurations:

± Count the total number of electrons

± Fill orbitals to match the nearest noble

gas of smaller atomic number

± Add remaining electrons to the next

orbitals according to Hund¶s rule ± Look for exceptions and correct the

configuration if necessary




Magnetic properties of atoms

± Each electron with spin orientation +1/2

has a partner with spin orientation -1/2

± The spins of these electrons cancel each

other, giving a net spin of zero

± An atom or ion with all electrons paired is

termed diamagnetic ± An atom or ion with unpaired electrons is

called paramagnetic




Excited states

± When an atom absorbs energy it can

reach an excited state with a new

electron configuration

± Excited atoms are unstable and

spontaneously return to the ground-state

configuration, giving up excess energy

± Excited states have practical

applications, such as fireworks

4 8 P i di it f t i4 8 P i di it f t i



4.8 Periodicity of atomic4.8 Periodicity of atomicpropertiesproperties

Atomic radii

± Moving from left to right across a period,

orbitals become smaller and of lower energy

± Atomic size decreases from left to right and

increases from top to bottom of the periodic

table

± A convenient measure of atomic size is theradius of the atom









Ionisation energy

± The minimum amount of energy needed

to remove an electron from a neutral

atom is the first ionisation energy (Ei1)

± Ionisation energy increases from left to

right across each period and decreases

from top to bottom of each group

± Ionisation energy does not change much

for elements in the d and f blocks








± Higher ionisations

A multi-electron atom can lose more than one

electron, but ionisation becomes more difficult as

cationic charge increases





± Irregularities in ionisation energies

Ionisation energies deviate somewhat from smooth

periodic behaviour. These deviations can be

attributed to shielding effects and electron-electronrepulsion





Electron affinity

± The energy change when an electron is added

to an atom in the gas phase is called the

electron affinity (EEA)

± Both ionisation energy and electron affinity

measure the stability of a bound electron, but

for different species






± The aufbau principle must be obeyed

when an electron is added to a neutral

atom, so the electron goes into the

lowest energy orbital available ± Electron affinity tends to become more

negative from left to right across a period

of the periodic table

± Values for electron affinities remainnearly constant among elements

occupying the same group

4 8 Periodicit of atomic4 8 Periodicit of atomic




4 8 Periodicity of atomic4 8 Periodicity of atomic



Sizes of ions

± An atomic cation is always smaller than the

corresponding neutral atom

± An atomic anion is always larger than the

corresponding neutral atom




4.84.8Periodicity ofPeriodicity of

atomicatomicpropertiesproperties

4 8 Periodicity of atomic4 8 Periodicity of atomic



Cation stability

± Beyond group 1 and 2, the removal of all

valence electrons is usually not energetically

possible

± Other metallic elements form ionic

compounds with cation

charges ranging

from +1 to +3


4 9 Ions and chemical4 9 Ions and chemical



4.9 Ions and chemical4.9 Ions and chemicalperiodicityperiodicity

Anion stability

± Halogen atoms readily accept electrons

to produce halide anions

± Isolated atomic anions with charges

more negative than ±1 are always

unstable, but oxide and sulfide are found

in many ionic solids





Metals, nonmetals and metalloids

± Metals can form cations relatively easily

± All metals have similar properties

± The halogens and noble gases are

distinctly nonmetallic

± Elements in any intermediate column of

the p block display a range of chemicalproperties. Carbon is a nonmetal, silicon

and germanium are metalloids, tin is a

metal





s-block elements

± The electron configurations of any element in

groups 1 and 2 of the periodic table contains a

core of tightly bound electrons and one or twos electrons that are loosely bound





p-block elements

± The properties of elements in the p block

vary across the entire spectrum of

chemical possibilities ± Metallic character diminishes rapidly as

additional p electrons are added

± Many of the nonmetals combine withoxygen to form polyatomic oxoanions



SummarySummary

All atoms display certain basic

characteristics

Electromagnetic radiation in the form

of light may be used to probe the

structure of atoms

Light has both wave and particle-like

properties

All electrons display certain basic

characteristics



SummarySummary

The Heisenberg uncertainty principle

states that we cannot simultaneously

know both the position and the

momentum of any particle

Electrons bound to atoms have

quantised energies, while free

electrons can have any energy Quantised properties of an atom can

be indexed using quantum numbers



SummarySummary

The electron distribution of an orbitalcan be depicted in a variety of ways

The energy of an orbital in a one-

electron system is determined by boththe atomic number and the principalquantum number

The ground-state configuration of anatom is the lowest possible energyarrangement of the electrons



SummarySummary

The electron configuration of an atom

can be written by recording quantum

numbers for all electrons in the atom,

or using a compact configuration

Atoms or ions having unpaired

electrons are paramagnetic

Species having all electrons pairedare diamagnetic



SummarySummary

The size of an atom is determined by

factors such as nuclear charge, orbital

energy and electron distribution and

shielding

The nature of cations and anions

formed from neutral atoms is

determined by the electronconfiguration of the atom



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Chemistry-Ch04_Atomic Energy Levels