1 Electronic Structure of Atoms & Periodic Table Chapter 4 CHEMISTRY - DMCU 1233 Fakulti Kejuruteraan Mekanikal, UTeM Lecturer: IMRAN SYAKIR BIN MOHAMAD

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Electronic Structure of Atoms & Periodic Table

Chapter 4

CHEMISTRY - DMCU 1233Fakulti Kejuruteraan Mekanikal, UTeM

Lecturer:IMRAN SYAKIR BIN MOHAMADMOHD HAIZAL BIN MOHD HUSINNONA MERRY MERPATI MITAN

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4.1 HISTORY OF ATOMIC MODEL

Contributor Model Explanation

John Dalton (1805)

Billiard Ball Model

Daltons atomic model was represented as a small united ball similar to a very tiny ball.

J. J Thomson (1897)

Plum Pudding Model

Thomson discovered the electron, a negatively charged particle. The atom was described as a sphere of positive charge with electrons embedded in it.

Ernest Rutherford

(1911)

Solar System Model

Rutherford discovered the proton, a positively charged particle in an atom. The proton and most of the mass of the atom were concentrated in the central region called the nucleus. The electrons moved in the spherical space outside the nucleus.

Neils Bohr (1913)

Bohr Model

According to Bohr, the electrons in an atom were not randomly distributed around the atomic nucleus, but moved around the nucleus in fixed orbits (shell). Each orbit formed a circle and had a fixed distance from the nucleus.

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4.1 HISTORY OF ATOMIC MODEL

Billiard Ball Model Plum Pudding Model

Solar System ModelBohr Model

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4.2 Quantum Numbers

• The Bohr model was 1-D model that used one quantum number to describe the distribution of electrons in the atom that representative of the size of the orbit, which was described by the principal quantum number (n).

• Meanwhile Schrödinger's model allowed the electron to occupy in 3-D space to describe the orbitals in which electrons can be found.

• Each electron in an atom is described by four different quantum numbers.

• The first three quantum number from Schrödinger's wave equations are the principal (n), angular (l), and magnetic (ml) quantum numbers describe the size, shape, and orientation in space of the orbitals on an atom.

• The fourth quantum number spin (ms) specifies how many electrons can occupy that orbital.

A quantum number describes the energies of electrons in atoms

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4.2 Quantum Numbers

Principal quantum number – ( n )

Angular momentum quantum number – ( l )

Magnetic quantum number – ( ml )

Spin quantum number – ( ms )

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Quantum Numbers (n, l, ml, ms)

Principal quantum number n

n = 1, 2, 3, 4, ….

n =1n=2

n=3

• Specifies the energy of an electron and the size of the orbital (the distance from the nucleus of the peak in a radial).

• All orbitals that have the same value of n are said to be in the same shell (level)

• The total number of orbitals for a given n value is n2

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Angular momentum quantum number l

for a given value of n, l = 0, 1, 2, 3, … n-1

n = 1, l = 0n = 2, l = 0 or 1n = 3, l = 0, 1, or 2

Shape of the “volume” of space that the e- occupies

l = 0 s orbitall = 1 p orbitall = 2 d orbitall = 3 f orbital


• Specifies the shape of an orbital with a particular principal quantum number.

• The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels).

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l = 0 (s orbitals)

l = 1 (p orbitals)

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l = 2 (d orbitals)

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l = 3 (f orbitals)

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Magnetic quantum number ml

for a given value of l ml = -l, …., 0, …. +l

orientation of the orbital in space

for l = 0 (s orbital) ml = 0if l = 1 (p orbital), ml = -1, 0, or +1

if l = 2 (d orbital), ml = -2, -1, 0, +1, or +2


• Specifies the orientation in space of an orbital of a given energy (n) and shape (l).

• This number divides the subshell into individual orbitals which hold the electrons.

• There are 2l+1 orbitals in each subshell.

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ml = -1 ml = 0 ml = 1

ml = -2 ml = -1 ml = 0 ml = 1 ml = 2

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ORIENTATION OF THE ORBITAL IN SPACE

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Spin quantum number ms

ms = +½ or -½


ms = -½ms = +½

Experimental arrangement for demo the spinning motion of electronsQ & A

• Specifies the orientation of the spin axis of an electron.

• An electron can spin in only one of two directions (sometimes called up and down).

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Quantum number

Symbol Meaning Range of values

Value examples

principal n shell 1 ≤ n n = 1, 2, 3, …

angular momentum

ℓsubshell

(s orbital is listed as 0, p orbital as 1 )

0 ≤ ℓ ≤ n − 1for n = 3:ℓ = 0, 1, 2 (s, p, d)

magnetic mℓ

Orbital(orientation of the subshell's shape)

−ℓ ≤ mℓ ≤ ℓfor ℓ = 2:mℓ = −2, −1, 0, 1, 2

spin ms

spin of the electron (−½ = "spin down",

+½ = "spin up")−s ≤ ms ≤ s

for an electron s = ½,so ms = −½, +½

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n l mlNumber of

orbitalsOrbitalName

Number ofelectrons

1 0 0 1 1s 2

2 0 0 1 2s 2

1 -1, 0, +1 3 2p 6

3 0 0 1 3s 2

1 -1, 0, +1 3 3p 6

2 -2, -1, 0, +1, +2 5 3d 10

4 0 0 1 4s 2

1 -1, 0, +1 3 4p 6

2 -2, -1, 0, +1, +2 5 4d 10

3 -3, -2, -1, 0, +1, +2, +3 7 4f 14

Relation between quantum number, atomic orbital and number of an electron

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Existence (and energy) of electron in atom is described by its unique Quantum Numbers

Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945)

No two electrons in the same atom can have identical values for all four of their quantum numbers


• Two electrons in the same orbital must have opposite spins.• Because an electron spins, it creates a magnetic field, which can be

oriented in one of two directions. • For two electrons in the same orbital, the spins must be opposite to each

other; the spins are said to be paired

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Paramagnetic

unpaired electrons

2p

Diamagneticall electrons paired

2p

• The substances are not attracted to magnets and are said to be diamagnetic.

• Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic.

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Shell – electrons with the same value of n

Subshell – electrons with the same values of n and l

Orbital – electrons with the same values of n, l, and ml

How many electrons can an orbital hold?

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How many 2p orbitals are there in an atom?

How many electrons can be placed in the 3d subshell?

Q & A

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Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.

1s1

principal quantumnumber n

angular momentumquantum number l

number of electrons in the orbital or subshell

Orbital diagram is shows the spin of the electron

H1s1

4.3 Electron configuration

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Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f

“Fill up” electrons in lowest energy orbitals

Aufbau principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states.

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Outermost subshell being filled with electrons

The order in which the electrons are filled in can be read from the periodic table in the following fashion

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H 1 electron H 1s1

Li 3 electrons Li 1s22s1

B 5 electrons B 1s22s22p1

C 6 electrons C 1s22s22p2

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The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule).

N 7 electrons N 1s22s22p3

O 8 electrons O 1s22s22p4

F 9 electrons F 1s22s22p5

Ne 10 electrons Ne 1s22s22p6

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What is the electron configuration of Mg?

What are the possible quantum numbers for the last (outermost) electron in Cl?

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Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Atoms gain electrons so that anion has a noble-gas outer electron configuration.

Of Representative Elements

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+1 +2 +3 -1-2-3

Cations and Anions Of Representative Elements

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Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]

O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

What neutral atom is isoelectronic with H- ?

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Electron Configurations of Cations of Transition Metals

When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.

Fe: [Ar]4s23d6

Fe2+: [Ar]4s03d6 or [Ar]3d6

Fe3+: [Ar]4s03d5 or [Ar]3d5

Mn: [Ar]4s23d5

Mn2+: [Ar]4s03d5 or [Ar]3d5

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Ion charges

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Name the orbital described by the following quantum numbers :

a. n = 3, l = 0b. n = 3, l = 1c. n = 3, l = 2d. n = 5, l = 0

Q & A session

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Give the n and l values for the following orbital a. 1s b. 3s c. 2p d. 4d e. 5f

What and the possible ml values for the following

types of orbital? a. s b. p c. d d. f

Q & A session

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How many possible orbital are there for n = a. 4 b. 10

How many electrons can inhabit all of the n = 4 orbital?

Place the following orbital in order of increasing energy: 1s, 3s, 4s, 6s, 3d, 4f, 3p, 7s, 5d, 5p

Q & A session

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Write electron configurations for the following atoms: a. H

b. Li+

c. N

d. F-

e. Ca

Q & A session

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Draw an orbital diagrams for atoms with the following electron configurations:

1s22s22p63s23p3

Q & A session

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When the Elements Were Discovered

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4.4 HISTORY OF THE PERIODIC TABLE

1. Antoine Lavoisier (1743–1794) Classify elements into four groups including light and heat, into metals and non-metals.

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2. Johann Dobereiner (1780–1849)• The first significant groupings of elements by place

certain elements in groups of three known as The Law of Triads.

• Founded that strontium had about the average properties of calcium and barium, and grouped these three together.

• Several more triad groups, including the halogen triad of chlorine, bromine, and iodine, and the alkali metal triad of lithium, sodium, and potassium.

• However, due to the inaccuracy of many measurements, including atomic weight, the relationship between large element groups could not be exacted

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3. John Newlands (1837–1898)• arranged known elements horizontally in

the ascending order of their atomic masses

• Each row consisted of seven elements. • Founded that elements with similar properties

repeated at every eighth element. • This arrangement was known as the Law of Octaves• However, this law was only obeyed by the first 17 elements.• There were no positions allocated for elements yet to be

discovered

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4. Lothar Meyer (1830–1895)• Plotted a graph of atomic volume against

atomic mass for all known elements. • Founded that elements with the same

chemical properties occupied the same relative positions on the curve.

• Showed that the properties of the elements were in a periodic pattern with their atomic masses.

• Proved that the properties of the elements recur periodically.

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5. Dmitri Mendeleev (1834–1907)• Showed that the properties of elements changed

periodically with their atomic mass.• Arranged the elements in the order of increasing

atomic mass and grouped them according to similar chemical properties.

• Able to predict the properties of undiscovered elements and left gap for these elements

• For examples correctly predicted the properties of the elements gallium, scandium and germanium which were only discovered later

• Mendeleevs table was used as a blueprint for the modern periodic table

Mendeleev’s periodic table

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6. Henry J. G. Moseley (1887–1915)• Based on the x-ray spectrum of elements studies,

he concluded that the proton numbers should be used as a basis for the periodic change of chemical properties instead of the atomic mass.

• Rearranged the elements in the ascending order of their proton numbers

• Similar to Mendeleev, Moseley left gaps for elements yet to be discovered.• produced a periodic table which was almost the same as Mendeleevs

periodic table. • Due to Moseley’s work, the periodic table was successfully developed and

being used today. • The modern periodic table is based on the arrangement of elements in the

ascending order of their proton numbers.

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7. Glenn Seaborg • discovered that the transuranium elements that

have atomic numbers from 94 to 102, resulting in the redesign of the periodic table

• Technically, both the lanthanide and actinide series of elements are to be placed between the alkaline earth metal and the transition metal.

• However, by doing this, the periodic table would be too wide. • Thus, the lanthanide and actinide series of elements were placed under the

rest of the periodic table. • Dr Seaborg and his colleagues were also responsible for identifying more

than 100 isotopes of elements.

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4.5 MODERN PERIODIC TABLE

The periodic table is a systematic classification of elements whereby elements with the same chemical properties are placed in the same group.

• The elements in the periodic table are arranged in rows called the periods and columns which are known as the groups

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Groups• There are 18 groups of elements in the periodic table. • Some of these groups have special names:

(a) Group 1 elements are called alkali metals.(b) Group 2 elements are called alkaline earth metals.(c) Group 3 to Group 12 elements are known as transition elements.(d) Group 17 elements are called halogens.(e) Group 18 elements are called noble gases.

• Each member of a group shows similar chemical properties although their physical properties such as density, melting point and colour show a gradual change when descending the group.

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Periods• There are seven rows from period 1 to period 7.• The elements are arranged horizontally in the ascending order of their

proton numbers in the periodic table.• The position of the period of an element in the periodic table is determined

by the number of shells occupied with electrons in the atom.• Period 1 has 2 elements only H and He, • Periods 2 and 3 have 8 elements each.• Periods 4 and 5, they have 18 elements each and they are called the

long periods. • Period 6 has 32 elements whereas the elements with proton number 58

to 71 are separated and are grouped below the periodic table known as the Lanthanide Series.

• Period 7 has 32 elements, the elements with proton number 90 to 103 are grouped below the periodic table known as the Actinide Series.

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Periodic Classification

Classification as metals and non-metals(a) Metals – a good conductor of heat and electricity.(b) Non-metals - a poor conductor of heat and electricity.(c) Metalloids – a intermediate between metal and non-metal properties

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ns1

ns2

ns2 n

p1

ns2 n

p2

ns2 n

p3

ns2 n

p4

ns2 n

p5

ns2 n

p6

d1 d5

d10

4f

5f


Ground State Electron Configurations of the ElementsThe similarity of the outer electron configuration (same type of valence electrons) makes the elements in the same group resemble one another in chemical behavior.

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Classification based on subshell filled with electron

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Representative elements (incompletely filled s and p subshells)

Transition metals (incompletely filled d subshells)

Noble gases (completely filled p subshells)

Actinides (incompletely filled 5f subshells)

Lanthanides (incompletely filled 4f subshells)

Zn, Cd, Hg (neither representative element nor transition metals)

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Periodic Trends

Periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of

the periodic table of elements

Elemental characteristicsa. Atomic radiusb. Ionization energyc. Electron affinityd. Electronegativitye. Metallic propertiesf. Non-metallic properties

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Periodic Trends

Atomic radius• The atomic radius is the distance between an atom's nucleus and its valence

electrons in an atom. • The atomic radius tends to decrease as one progresses across a period from

left to right because the effective nuclear charge increases, thereby attracting the orbiting electrons and reducing the radius.

• The atomic radius usually increases while going down a group due to the addition of a new energy level (shell).

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Periodic Trends

Ionic Radius• The ionic radius is different from the atomic radius of an element. • Positive ions are smaller than their uncharged atoms. • Negative ions are larger than their atoms.

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Periodic Trends

Ionization energy• The ionization potential is the minimum amount of energy required to remove

one electron from each atom.• Ionization energies increase moving from left to right across a period

(decreasing atomic radius). • Ionization energy decreases moving down a group (increasing atomic radius).

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Periodic Trends

Electron affinity• The electron affinity described as the energy gained by an atom when an

electron is added. • Electron affinities becoming increasingly from left to right• Electron affinities change little moving down a group, however they do

generally become slightly more positive (less attractive toward electrons)

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Periodic Trends

Electronegativity• Electronegativity refers to the ability of an atom to attract the electrons of

another atom to it when those two atoms are associated through a bond. • Electronegativity generally increases moving across a period and decreases

moving down a group. • Electronegativity plays a very large role in the processes of Chemical

Bonding.

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Periodic Trends

Metallic properties• Metallic property decreases across a period with increase in number of

valence electrons as well as a decrease in atomic radius, and it increases down the group with increase in number of shells and atomic radius.

Non-metallic properties• Non-metallic property increases across a period and decreases down the

group due to the same reason

Documents

1 Electronic Structure of Atoms & Periodic Table Chapter 4 CHEMISTRY - DMCU 1233 Fakulti Kejuruteraan Mekanikal, UTeM Lecturer: IMRAN SYAKIR BIN MOHAMAD