autumn lecture 2 (electron configs)

8/13/2019 Autumn Lecture 2 (Electron Configs)

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Lecture 2: The Arrangement of Electrons in Atoms(Chemistry in Context ch. 6)

Evidence for Electron ArrangementsAtomic Spectra

Successive Ionisation Energies

Comparison of First Ionisation Energies

Electron Arrangements

Quantum Numbers

Shapes of Orbitals

Pauli Exclusion Principle

Aufbau Principle

Electron Configurations

Writing Electron configurations

Electronic Configurations & the Periodic Table


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The Spectrum of Atomic Hydrogen

-model (theory) of structure of atom based onnegative electrons orbiting a compact, positive

nucleuselectrostatic attraction

-spectrum of atomic hydrogen:energy is lost by emission of light

problem why does spectrum of atomic hydrogen showemissions at certain wavelengths only?

Why not continuous range of wavelengths like a rainbow?


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long wavelength (l)low energy (E)

low frequency (f)

short wavelength (l)high energy (E)

high frequency (f)

Wavelength, Energy and Frequency Reminder


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(continuous spectrum)

e.g. light bulb, sunlight


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(line spectrum e.g. sparks, excited gases)

neon signs use similar glowing electrified noble gas


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The Bohr Atom

Bohr: hydrogens electron only adopts certain orbits of particular

energy (and distance from nucleus) electrons energy is quantised

Quantisationonly

particular discrete energy

levels are allowed. Like

climbing stairsyou cant

stand between steps!

Niels

Bohr


8/30(Lyman series)

high

energy

low

energy

-photon of light emitted as

excitedhydrogen atom loses

energy (when electron falls

from higher energy state tolower)

convergence limit

of Lyman Series


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Lyman series(UV)

electrons falling to groundstate

Balmer series(visible)electrons falling to n = 2state

n.b. electrons falling from:

n = state to n = 1stateemit radiation at the convergence limit

(equals to atoms ionisation energy):

H+(g) + e H(g)

groundstate

excited

states

converging energy levels


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Ionisation Energies (6.3)

-ionisation energyis energy needed to remove an electron from an

atom (to form a cation):

first

ionisation

second

ionisation

relatively easy

little energy

required

low value

all electrons

attracted to

positive nucleus

ionisations

areendothermic

more difficult

(attraction to cation)

more energy

requiredhigher value

M M+ M2+

N t i l


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Ionisation Energies (6.3)

-comparing succesive ionisation energies tells

us that atoms electrons are arranged in shells:

first I.E. second I.E. third I.E. (kJ/mol)

Al(Gp III) 577 1820 2740

Mg(Gp II) 736 1450 7740

Na(Gp I) 494 4560 6910

-electron removal from

outer (valence) shell (easy)-electron removal from

inner (core) shell (difficult)

Na has a single outer electron, Mg has 2 and Al has 3

Na reacts vigorously

with water


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Plotting Successive Ionisation Energies (6.4)

-the electron shells are more obvious if we remove all the electrons

of a large atom e.g. Na:

we can describe electron configuration of Na as 2-8-1

2 electrons in innermost

shell (n=1)

1 electron in outer shell (n=3)


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-Q: What elements successive ionisation

energies are plotted here?


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4 electrons in outer shell (n=3)


-Q: What elements successive ionisation

energies are plotted here?

we can describe electron configuration as 2-8-4

has 4 valence electrons (it is in Group IV) Si

2 electrons in

innermost

shell (n=1)


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-I.E. is highest for

stable noble gaselectron

configurations

-little steps

indicate sub-shells

period 2 period 3

Elements I.E. & Electron Subshells (6.5)

-evidence for shell structure also comes

from comparing first I.E. of many elements

-I.E. increases across a period

(more protons

added so electrons

pulled in closerby greater nuclear

charge)

1s

2s

2p3p

3s

4sn =2 shell n= 3 shell


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Electron shells house atomic orbitals (6.5)

- there are n2orbitals per shell i.e.

one orbital for n= 1

four orbitals for n= 2

nine orbitals for n= 3

1s

2s 2p 2p 2p

3s 3p 3p 3p 3d 3d 3d 3d 3d

Schrodingers Equation(1926)

solutions are wavefunctions (y)

called atomic orbitals

Erwin Schrodinger

Th H d 1 bit l (6 6)


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The Hydrogen1sorbital (6.6)

-dot plot shows electron density

in 1sorbital

(highest close to nucleus)

-small chance of finding electron

far away from nucleus

to draw shape of orbital we can draw region within

which there is 90% chance of finding electron

region is spherical:

(all sorbitals are spherical)


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2sorbitalhigher in energy than (further from nucleus)

2sorbitalis also spherical but there is a nodal surface where

electron cannot be found

1s 2s


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porbitals have two lobes:

-the threeporbitals within shells n= 2

are at right angles to each other i.e. 2px, 2py, 2pz

an eggtimer


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but within each shell all are same energy (degenerate) in one electron atoms.

name indicates the shell, shape of orbital and the orientation of its lobes.

e.g.

2py

indicates shape (symmetry) of orbital.

Determined by quantum number l

indicates orientation of

orbital. Determined byquantum number ml

indicates radius oforbital. Determined by

principal quantum

number n


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Electron Spin and ms

Each orbital can hold two electrons. These must be of opposite spin

(Pauli Exclusion Principle).

Pauli Exclusion Principleeach electron in any atom has a

unique set of four quantum numbers

-an electrons set of quantum numbers is like a unique address within the

atom - tells us shell number, type of orbital, orientation of orbital,

electron spin

fourth quantum number

electron spin

ms= + or -


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Building Up Principle

(6.7)

-to explain electronarrangement in all elements:

-add successive electrons to

hydrogen to form heavier

elements

-orbitals fill from lowest

energies to highest:

the Aufbau principle

-in many-electron atoms,

orbitals within each shell no

longer degenerate

e.g. 2porbitals higher in

energy than 2sorbitals


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-for element of atomic no. Z, we add Z electrons to fill its orbitals

-hydrogenhas a half-full 1s orbital its electron configurationas 1s1

Hehas a full 1sorbital1s2

Liis 1s22s1or because of its He core we can write [He]2s1

Similarly, Beis [He]2s2


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its more stable for carbons two 2pelectrons have parallel spinsand reside in two separate 2p orbitals of equal energy.

-only outermost shell contains valence electronsthose available to

take part in reactions

-at boron (Z = 5) one 2porbital is half occupied

-for C there is a choicesixth electron can either pair up with fifth or

else go into a newporbital


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-similarly N has three unpaired electrons

at oxygen electrons begin to pair up so O has only two unpaired electrons:

At Ne, n= 2 shell is full (atom has 8 valence electrons - a stable octet)

Ne forms no compounds


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-after Ne, period 3 elements (n= 3) starts to fill (same as 2nd period)

-the 3sorbital starts to fill first at sodium:

He noble

gas core

Ne noble

gas core


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fills first fills after

The Building-up Principle has a surprise after Argon (Z = 40):

4sorbital is lower in energy than3dso fills first.

K has configuration [Ar]4s1and NOT [Ar]3d1.


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fills first fills after

The Building-up Principle has a surprise after Argon (Z = 40):

4sorbital is lower in energy than3dso fills first.

K has configuration [Ar]4s1and NOT [Ar]3d1.

order of

filling


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2K(s) + H2

O(l) ??? (aq) + ??? (g)

-potassium reacts more violently with water than does Na (WHY?)

-the reaction produces which gas? (and which salt?)

autumn lecture 2 (electron configs)

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