ATOMIC STRUCTURE

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atom structure

Presentation

Start

3- Electronic Configuration

2- Isotopic Symbol

1- Fundamental Particles

LEARNING OBJECTIVES

4- Ionisation Energy

6-Relative Atomic Mass

5- Mass Spectrometry

atom Structure

TIMELINE

Properties of Subatomic Particles

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Comprehension:

• a) Which of these three particles will be deflected the most by the electric field?
• b) In which direction will this particle be deflected?
• c) Explain your answer.

Atomic number and mass number

How many NEUTRONS?

Atoms of the same element that contain the same number of protons but different numbers of neutrons.

ISOTOPE

ISOTOPES OF HYDROGEN

State one similarity and one difference in the properties of these isotopes of HYDROGEN. Explain your answer.

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How many protons, neutrons and electrons?

In a neutral atom the number of positively charged protons in the nucleus equals the number of negatively charged electrons outside the nucleus. When an atom gains or loses electrons, ions are formed, which are electrically charged.

ATOMIC AND IONIC RADIUS

https://teachchemistry.org/classroom-resources/periodic-trends-simulation

electron arrangement

ELECTRON CONFIGURATION

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Summary of arrangement:

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ORBITALS

d-orbital

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s-orbital

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f-orbital

p-orbital

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Filling of atomic orbitals

Aufbau Principle

Hund's Rule

Pauli Exclusion Principle

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Electron Configuration

1. In finding the number of electrons refer to the number of protons: For a neutral atom: no. protons = no. electrons Ion: no. protons - charge = no. electrons; the total number of electrons to be accommodated in the orbitals.

Na

2. Fill the orbitals from the lowest energy (1s) upward; no orbital must contain more than 2 electrons.

3. For subshell that contains more than 1 orbital (p,d & f) place the electrons into different orbitals until all are singly occupied. Only then should further electrons in that sub-shell doubly occupying orbitals. .

4. 2 electrons sharing the same orbital must have opposite spin.

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Electron Configuration: Short-hand and Electrons-in-boxes diagram

[Ne] 3s1

1s2 2s2 2p6 3s1

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Try

Orbitals and Periodic Table

Check comprehension

a An element has the electronic configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 5p5. i Which block in the Periodic Table does this element belong to? ii Which group does it belong to? iii Identify this element. b Which block in the Periodic Table does the element with the electronic configuration 1s2 2s2 2p6 3s2 3p6 3d5 4s1 belong to?

ionisation: energy, 1st and successive ionisation

The 1st ionisation energy of an element is the energy needed to remove one electron from each atom in one mole of atoms of the element in the gaseous state to form one mole of gaseous 1+ ions.

1st ionisation

the amount of energy required to remove an electron from an atom or ion in the gas phase

successive ionisation energies

Factors Affecting Ionisation Energies

distance from the nucleus to the outermost electron

shielding effect

the nuclear charge

Succesive ionisation of Na

Using information from the successive energy values, one can deduce the arrangement of electrons in an atom.

WOW effect

Practice!

1. The successive ionisation energies of boron are shown in the table:

i Why is there a large increase between the third and fourth ionisation energies? ii Explain how these figures confirm that the electronic structure of boron is 2, 3.

2. For the element aluminium (Z = 13),a. Draw a sketch graph to predict the log10 of the successive ionisation energies (y-axis) against the number of electrons removed (x-axis).b. How does the graph provide evidence for the existence of three electron shells in an aluminium atom? c. Write an equation, including state symbols, to represent the 2nd ionisation energy of aluminium. d. Write the electronic configuration of an aluminium ion, Al3+, using 1s2 notation.

Investigate.

Why second ionisation energy of sodium is more than the second ionisation energy of magnesium?

On the 1st Ionisation trend for the 1st 20 elements why are the ionisation energies drop from Be to B, Mg to Al, N to O and P to S?

Don’t forget the trends in the radius (atomic & ionic)!

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Weighing Atoms = Mass Spectrometry

Mass spectrometry is an accurate instrumental technique used to determine the relative isotopic mass (mass of each individual isotope relative to carbon-12) and the relative abundance for each isotope. From this, the relative atomic mass of the element can be calculated.

Some uses of mass spectrometry include:

• carbon-14 dating
• detecting illegal drugs
• forensic science
• space exploration.

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What is relative atomic mass?

The relative atomic mass is the weighted average mass of naturally occurring atoms of an element on a scale where an atom of carbon-12 has a mass of exactly 12 units.

Most elements have more than one isotope. The Ar of the element is the average mass of the isotopes taking into account the abundance of each isotope. This is why the Ar of an element is frequently not a whole number.

Using mass spectra to calculate Ar

The mass spectrum of an element indicates the mass and abundance of each isotope present. For example, the mass spectrum of boron indicates two isotopes are present:

How can this be used to calculate the Ar of boron?

Calculating Ar

In a sample of magnesium, 79.0% of the magnesium atoms are 24Mg, 10.0% are 25Mg and 11.0% are 26Mg.

Calculating Ar

Beryllium has the electronic structure 1s2 2s2 and boron has the electronic structure 1s2 2s2 2p1. The fifth electron in boron must be in the 2p subshell, which is slightly further away from the nucleus than the 2s subshell. There is less attraction between the fifth electron in boron and the nucleus because: i the distance between the nucleus and the outer electron increases slightly ii the shielding by inner shells increases slightly iii these two factors outweigh the increased nuclear charge.

As the atomic number (number of protons) increases, the positive nuclear charge increases. The bigger the positive charge, the greater the attractive force between the nucleus and the electrons. So, more energy is needed to overcome these attractive forces if an electron is to be removed.

There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period. The ΔHi1 for lithium is much smaller than the ΔHi1 for helium. Helium has two electrons. These are in the first quantum shell. But lithium has three electrons. The third electron must go into the next quantum shell further away from the nucleus. So, the force of attraction between the positive nucleus and the outer negative electrons decreases because:i the distance between the nucleus and the outer electron increases ii the shielding by inner shells increases

• describe the structure of atoms, including the masses, electrical charges, and locations of protons, neutrons, and electrons. -understand the terms atomic and proton number; and mass and nucleon number
• -describe the distribution of mass and charge within an atom
• describe the behavior of beams of protons, neutrons, and electrons moving at the same velocity in an electric field define the term isotopes; and understand the notation of isotopic symbol, x/y A.
• state and explain why isotopes of the same element have the same chemical properties but different physical properties.

How it works?

X(g) X+(g) + e-

the spin-pairing of the electrons plays a part here:The electron removed from the oxygen is from the orbital that contains a pair of electrons. The extra repulsion between the pair of electrons in this orbital results in less energy being needed to remove an electron. So the IE for oxygen is lower, because of spin-pair repulsion.

1s2 2s2 2p6 3s1

The relative atomic mass (Ar) of an element is the mass of one of its atoms relative to 1/12 the mass of one atom of carbon-12.

Fig. 2.23 page 40 Coursebook

X(n-1)+(g) Xn+(g) + e-