unit 12: atomic structure ib chemistry ib topic 2 and 12

Unit 12: Atomic Structure

IB ChemistryIB Topic 2 and 12

IB Topics 2 and 12

2.1.1 State the position of protons, neutrons and electrons in the atom.

2.1.2 State the relative masses and relative charges of protons, neutrons and electrons.

2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an element.

2.1.4 Deduce the symbol for an isotope given its mass number and atomic number.

IB Topics 2 and 12

2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge.

HISTORY OF THE ATOMHISTORY OF THE ATOM

460 BC Democritus develops the idea of atoms

he pounded up materials in his pestle and

mortar until he had reduced them to

smaller and smaller particles which he

called

ATOMAATOMA

(Greek for indivisible)


1808 John Dalton

suggested that all matter was made up of

tiny spheres that were able to bounce

around with perfect elasticity and called

them

ATOMSATOMS

Dalton’s Atomic Theory (1808)

1. Elements are composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements.

2. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same.

3. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions.


1898 Joseph John Thompson

found that atoms could sometimes eject a

far smaller negative particle which he

called an

ELECTRONELECTRON

J.J. Thomson, measured mass/charge of e-

(1906 Nobel Prize in Physics)

e- charge = -1.60 x 10-19 C

Thomson’s charge/mass of e- = -1.76 x 108 C/g

e- mass = 9.10 x 10-28 g

Measured mass of e-

(1923 Nobel Prize in Physics)


Thompson develops the idea that an atom was made up of

electrons scattered unevenly within an elastic sphere

surrounded by a soup of positive charge to balance the

electron's charge

1904

like plums surrounded by pudding.

PLUM PUDDING

MODEL


1910 Ernest Rutherford

oversaw Geiger and Marsden carrying out

his famous experiment.

they fired Helium nuclei at a piece of gold

foil which was only a few atoms thick.

they found that although most of them

passed through. About 1 in 10,000 were

deflected and, to their surprise, some

helium nuclei bounced straight back.

1. atoms positive charge is concentrated in the nucleus2. proton (p) has opposite (+) charge of electron3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)

particle velocity ~ 1.4 x 107 m/s(~5% speed of light)

(1908 Nobel Prize in Chemistry)

atomic radius ~ 100 pm = 1 x 10-10 m

nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

Rutherford’s Model of the Atom


Rutherford’s new evidence allowed him to propose a

more detailed model with a central nucleus.

He suggested that the positive charge was all in a

central nucleus. With this holding the electrons in place

by electrical attraction

However, this was not the end of the story.


1913 Niels Bohr

studied under Rutherford at the Victoria

University in Manchester.

Bohr refined Rutherford's idea by

adding that the electrons were in

orbits. Rather like planets orbiting the

sun. With each orbit only able to

contain a set number of electrons.

Bohr’s Atom

electrons in orbits

nucleus

ATOMIC STRUCTUREATOMIC STRUCTURE

Particle

proton

neutron

electron

Charge

+ ve charge

-ve charge

No charge

1

1

nil

Mass

Subatomic Particles

Particle Mass

(g) Charge

(Coulombs) Charge (units)

Electron (e-) 9.1 x 10-28 -1.6 x 10-19 -1

Proton (p) 1.67 x 10-24 +1.6 x 10-19 +1

Neutron (n) 1.67 x 10-24 0 0

mass p = mass n = 1840 x mass e-

Subatomic Particles

For IB purposes we use:

1 atomic mass unit (amu) as the mass of the proton and neutron

5x10-4 amu for the mass of the electron.

Atomic number (Z) = number of protons in nucleus

Mass number (A) = number of protons + number of neutrons

= atomic number (Z) + number of neutrons

Isotopes are atoms of the same element (X) with different numbers of neutrons in the nucleus

XAZ

H11 H (D)2

1 H (T)31

U23592 U238

92

Mass Number

Atomic NumberElement Symbol

An ion is an atom, or group of atoms, that has a net positive or negative charge.

cation – ion with a positive chargeIf a neutral atom loses one or more electronsit becomes a cation.

anion – ion with a negative chargeIf a neutral atom gains one or more electronsit becomes an anion.

Na 11 protons11 electrons Na+ 11 protons

10 electrons

Cl 17 protons17 electrons Cl-

17 protons18 electrons

A monatomic ion contains only one atom

A polyatomic ion contains more than one atom

Na+, Cl-, Ca2+, O2-, Al3+, N3-

OH-, CN-, NH4+, NO3

-

13 protons, 10 (13 – 3) electrons

34 protons, 36 (34 + 2) electrons

Do You Understand Ions?

How many protons and electrons are in Al2713 ?3+

How many protons and electrons are in Se7834

2- ?

SUMMARYSUMMARY

1. The Atomic Number of an atom = number of

protons in the nucleus.

2. The Atomic Mass of an atom = number of

Protons + Neutrons in the nucleus.

3. The number of Protons = Number of Electrons.

4. Electrons orbit the nucleus in shells.

5. Each shell can only carry a set number of electrons.

Objectives

2.1.6 Compare the properties of the isotopes of an element.

2.1.7 Discuss the uses of radioisotopes.

Isotopes

Atoms with the same number of protons, but

different numbers of neutrons.

Atoms of the same element (same atomic

number) with different mass numbers Isotopes of chlorine

35Cl 37Cl17 17

chlorine - 35 chlorine - 37

Learning Check

Naturally occurring carbon consists of three isotopes, 12C, 13C, and 14C. State the number of protons, neutrons, and electrons in each of these carbon atoms.

12C 13C 14C 6 6 6

#p _______ _______ _______

#n _______ _______ _______

#e _______ _______ _______

Solution

12C 13C 14C 6 6 6

#p 6 6 6

#n 6 7 8

#e 6 6 6

Learning Check

An atom of zinc has a mass number of 65.

A. Number of protons in the zinc atom 1) 30 2) 35 3) 65

B. Number of neutrons in the zinc atom 1) 30 2) 35 3) 65

C. What is the mass number of a zinc isotope with 37 neutrons?

1) 37 2) 65 3) 67

Solution

An atom of zinc has a mass number of 65.

A. Number of protons in the zinc atom 1) 30

B. Number of neutrons in the zinc atom2) 35

C. What is the mass number of a zinc isotope with 37 neutrons?3) 67

Learning Check

Write the atomic symbols for atoms with the following:

A. 8 p+, 8 n, 8 e- ___________

B. 17p+, 20n, 17e- ___________

C. 47p+, 60 n, 47 e- ___________

Solution

16OA. 8 p+, 8 n, 8 e- 8

B. 17p+, 20n, 17e- 37Cl 17

C. 47p+, 60 n, 47 e- 107Ag 47

Atomic Mass on the Periodic Table

11Na

22.99

Atomic Number

Symbol

Atomic Mass

Atomic Mass

Atomic mass is the weighted average mass of all the atomic masses of the isotopes of that atom.

Example of an Average Atomic Mass

Cl-35 is about 75.5 % and Cl-37 about 24.5% of natural chlorine.

35 x 75.5 = 26.425

100

37 x 24.5 = 9.065 35.49 100

Which isotope is it?

1. A = 15, Z = 8

2. A = 36, Z = 17

3. A = 235, Z = 92

Which isotope is it?

1. A = 15, Z = 8

Answer : O

2. A = 36, Z = 17

Answer : Cl

3. A = 235, Z = 92

Answer : U

Radioisotopes

Radioactive isotopes of all elements can be produced by exposing the natural element to a flux of slow moving neutrons in a nuclear reactor.

This makes the nucleus capture an extra neutron.

The stability of a nucleus depends on the balance between the number of protons and neutrons.

When a nucleus has too many or too few neutrons it is radioactive and changes to a more stable nucleus by giving out radiation.

Radioisotopes

The radiation have different forms:

Alpha particles: emitted by nuclei with too many protons to be stable. They have 2 protons and 2 neutrons (the same as a helium nucleus).

Beta particles: emitted by nuclei with too many neutrons. They are electrons ejected from the nucleus as a neutron decays.

Gamma rays are a form of electromagnetic radiation.

Radioisotopes

These radioisotopes have many uses:

Generate energy in nuclear power stations

Sterilize surgical instruments in hospitals

Preserve food

Fight crime

Detect cracks in structural materials.

Carbon-14 Dating

The radioactive decay of carbon-14 is used to date carbon containing materials ex. fossils and archaeological objects.

14C is naturally present in the atmosphere and the abundance in living things is constant – it is absorbed from 14CO2 in the air.

When organisms die the 14CO2 steadily decreases as the radioactive isotope decays by releasing beta particles.

The half-life of the isotope is 5730 years so the amount of 14C relative to 12C is compared to find how old the item is.

Iodine-131as a Medical Tracer

Iodine-131 emits beta and gamma radiation and has a short half-life of 8 days. It is quickly eliminated from the body.

The activity of the thyroid gland can be monitored when a person takes a drink that contains iodine -131. This can also diagnose and treat thyroid cancer.

Iodine-125 emits beta radiation, has a half-life of 80 days and is used to treat prostate cancer by implanting it in the gland.

Cobalt-60 in Radiotherapy

Radiotherapy is the use of radioisotopes to treat cancerous cells. The treatment damages the DNA in the cell making it impossible for the cell to reproduce.

60Co is an example of this. It emits gamma radiation.

Radiotherapy is used to treat localized solid tumors like skin, brain and breast; and blood cancer like leukemia.

Healthy cells can recover is dosage is controlled.

Radioisotopes

What risks are associated with the use of radioactive material?

How are these materials controlled?

Mass Spectrometer

2.2.1 Describe and explain the operation of a mass spectrometer.

2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass using the 12C scale.

2.2.3 Calculate the non-integer relative atomic masses and abundance of isotopes from given data.

Mass Spectrometer

Mass Spectrometer

Separates particles according to their masses and records the abundance of each mass.

Process is as follows:

1. Substance is vaporized and a vacuum is maintained.

2. Gaseous atoms or molecules are ionized to have a positive charge by being bombarded by electrons.

Mass Spectrometer

3. Positive particles are accelerated by a large potential difference between 2 electrodes.

4. Fast moving particles pass through an electromagnet which deflects them according to their mass - the smaller particles will be deflected more.

Mass Spectrometer

5. Particles of a certain mass (adjustable) are detected by the detector plate.

6. The vacuum is there to prevent collision with other particles.

Mass Spectrum

Calculating Relative Molecular Mass

mass of isotope1 x % abundance1 / 100

+

mass of isotope2 x % abundance2 / 100

=

relative molecular mass of element, Mr


A mass spectrum of chlorine shows there to be 25% 37Cl and 75% 35Cl. Calculate the relative atomic mass of chlorine in this sample.


A mass spectrum of chlorine shows there to be 25% 37Cl and 75% 35Cl. Calculate the relative atomic mass of chlorine in this sample.

Solution:

(0.25 x 37) + (0.75 x 35) = 35.5

Calculating % abundance from atomic mass

Boron exists as 10B and 11B. Its relative atomic mass is 10.81

If x = # of 10B atoms then 100 – x = # of 11B atoms

Total mass = 10x + (100-x)11 = 10x + 1100 – 11x= 1100 – x

Average mass = total mass / number of atoms= (1100-x) / 10010.81 = 1100 – xx = 1100 – 1081 = 19.00

The abundances are 10B = 19.00% and 11B = 81.00%

Electron Arrangement

2.3.1 Describe the electromagnetic spectrum.

2.3.2 Distinguish between a continuous spectrum and a line spectrum.

2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels.

2.3.4 Deduce the electron arrangement for atoms and ions up to Z = 20.

Electromagnetic Spectrum

Light is a general word used to describe all wavelengths of electromagnetic radiation - from short gamma rays to long radio waves.

White light is made up of all the visible colors of light.

Electromagnetic Radiation

Electromagnetic radiation is commonly called light waves.

All light waves have the same speed (c) but have different wavelengths (λ) and therefore different frequencies (f).

Frequency and wavelength are inversely proportional since fewer long waves can pass the same point in space in a certain period of time.

They are related by the equation c = fλ

Electromagnetic Radiation

c = fλ

Where c is the speed of light

f is frequency measured in Herz (Hz)

is wavelength measured in meters (m)

c = 3 x 108 ms-1

Continuous Spectrum

If white light is passed through a prism we see the different colors that make up the light.

The different wavelengths blend into each other so this is called a continuous spectrum.

Evidence for orbitals

The best evidence for the existence of energy levels in an atom comes from studying emission spectra of elements.

When an element is excited it will often emit a light of a characteristic color ex. Red light in neon signs.

Colors

Gases can be excited by passing electricity through them at low pressure.

Some metals will show a distinctive color when heated in the bunsen flame.

Flame tests are commonly used to identify alkali metals.

Flame Colors

Lithium - red

Sodium - yellow

Potassium - lilac

Copper - green

Barium - yellow-green

Calcium - brick red

Line Spectrum

If the light that is emitted is looked at through a prism or diffraction grating, different lines of color can be seen.

These lines make up the line spectrum.

Each element has its own specific line spectrum.

Emission vs Absorption Spectra An emission spectrum shows lines of color

against a black background. The lines of color are the wavelengths of light that are produced.

An absorption spectrum shows lines of black on a colored background. The lines of black are the wavelengths where light is absorbed.

Planck’s Equation

When an atom is excited its electrons gain energy and they move to a higher energy level.

When returning to the lower energy level they lose this energy as light which we see as colors.

Only certain energy levels can be reached so only certain values of energy can be lost, therefore only certain frequencies of light can be emitted.

Planck’s Equation

The color of the light is measured by frequency and is related to energy by the equation:

∆E = hf

where h is Planck’s constant, E is energy and f is frequency (1Hz = 1s-1).

Planck’s constant = 6.62 x 10-34 m2 kg / s

Atomic Emission Spectrum of Hydrogen

This is the simplest emission spectrum.

Hydrogen only has 1 electron so there is no repulsion from other electrons.

The spectrum consists of bright lines.

The lines are in series named after the scientists who discovered them.

Hydrogen emission spectrum

Hydrogen Emission Series

Each series has a similar structure of lines that converge (get closer together) at higher frequencies / higher energies.

Each series corresponds to transitions in which the electron falls to a particular energy level.

The largest energy transition is from n = 2 to n = 1 and the lines are in the U.V. region (high energy)

All transitions to the n = 2 level are in the visible region.

Transitions to n = 3 are in the infrared region.

Hydrogen Emission Spectrum

Each series ends in a brief continuum at the high frequency end where the lines become too close together to be separated.

When an electron is at n = ∞, it is no longer in the atom and the atom has been ionized (positive ion formed).

In the case of the Lyman series this is the ionization energy of hydrogen.

Electron Configuration

12.1.1 Explain how evidence from first ionization energies across periods accounts for the existence of main energy levels and sub-levels in atoms.

12.1.2 Explain how successive ionization energy data is related to the electron configuration of an atom.

Trends in Ionization Energies

Ionization energy is the energy required to remove an electron from an atom to form a positive ion. Units: kJmol-1.

Ex. Al(g) Al+(g) + e-

The second ionization energy refers to removing another electron from a 1+ ion:

Al+(g) Al2+(g) + e-

Note that all of these changes are measured in the gas phase.


The successive ionization energies of Aluminum are shown below. Note the jump from the 3rd electron to the 4th and again from the 12th electron to the 13th.


There is an overall increase in ionization energy as more electrons are removed.

The first ionization involves separating an electron from a neutral atom.

The second ionization involves removing an electron from a positively charged ion.

As more electrons are removed the remaining electrons are more strongly attracted to the increasingly positively charged ion – so more energy is required.


The large increases in successive ionization energies indicate that the attractive force holding the electrons is becoming stronger.

This evidence led to the theory of different energy levels within an atom.

Electrons must be removed from outer energy levels first.

Then electrons in energy levels closer to the nucleus require are removed but more energy is required.


Atomic theory explains that an atom has discrete (separate) energy levels that hold a certain number of electrons.

As electrons are removed the sudden increases in ionization energies indicate an electron has been removed from an energy level nearer the nucleus.

If you look closely at ionization data you will see smaller increases within one energy level indicating the sub-levels that exist in atoms.

Energy Levels

12.1.3 State the relative energies of s, p, d and f orbitals in a single energy level.

12.1.4 State the maximum number of orbitals in a given energy level.

12.1.5 Draw the shape of an s orbital and the shapes of the px, py and pz orbitals.

12.1.6 Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 54.

Energy Levels

Electron configuration is a way to show where each electron in an atom exists.

We think of electrons as being in energy levels.

The lowest energy level is closest to the nucleus.

As we move further from the nucleus the energy increases.

Energy Levels

The main energy level is assigned a number with level 1 being closest to the nucleus. It is labeled n=1.

Each main energy level contains 1 or more sub-level which are labeled with the main energy number followed by the type of orbitals it contains.

We use the letters s, p, d and f to denote each type of orbital.

Ex. Energy level n = 2 contains sub-levels 2s and 2p.

Orbitals

An orbital is a region in space around the nucleus that an electron is likely to be found.

It is only a probability given by electron density - we cannot pinpoint the location of an electron because it is always moving very quickly.

Each orbital has a different 3-D shape and capacity for electrons.

Heisenberg’s Uncertainty Principle

It is possible to think of light waves as discrete particles – called photons.

In the same way electrons (or any object) can be thought of as behaving like waves.

In reality the behavior is somewhere between these two extremes.

Heisenberg said that it is impossible to make an exact measurement of both the position and momentum of any object at the same time.

Heisenberg’s Uncertainty Principle

Locating an electron in an atom is an example of this principle.

The act of trying to find the electron using radiation gives the electron a random “kick” sending it off in a random direction.

In other words, the more you know about an electrons location the less you know about its speed and vice versa.

Orbitals

The s orbital has a spherical shape and each s orbital can hold 2 electrons.

When an electron is in an energy level it spins in a particular direction.

When there are two electrons in an orbital together they must spin in opposite directions.

This is a result of the PAULI EXCLUSION PRINCIPLE.

Orbitals

p orbitals have 3 sub-levels and each can hold 2 electrons.

This gives a total capacity of 6 electrons.

The p orbitals have a dumb-bell shape in 3 dimensions.

Orbital Shapes

Electronic Structure

IB describes how many electrons are in each period using “electronic structure”.

They simply write the number of electrons in each period separated by a comma.

Ex. H is 1, He is 2 and Li is 2,1

Write the electronic structure for Be, Al, Kr

Using the periodic table to count up electrons will help you do this.

Electron Configuration

We use the orbitals to show how many electrons an atom or ion has and where they are located.

It is similar to a ticket for a sports event where section, row and seat number are given for each person.

How to write electron configuration

The first part of the electron configuration denotes which period (row) of the periodic table you are talking about.

The second part tells us the orbital, s,p, d or f.

The third part is a superscript number that tells us how many electrons are in the orbital.

For example:

1s1 is H 1s22s1 is Li

Learning Check

Write the electron configurations for the following:

1. He

2. C

3. F

Learning Check

Write the electron configurations for the following:

1. He 1s2

2. C 1s22s22p2

3. F 1s22s22p5

Rules for Electron Configuration

The AUFBAU PRINCIPLE states that orbitals fill up with electrons from the lowest energy first.

HUND’S RULE states that the lowest energy is achieved when the maximum number of electrons have the same spin direction within an energy level.

Ex. rather than . . .

IB expects you to use arrows in boxes to represent electrons – PIC!

Order of Filling Orbitals

Electron Configuration and the Periodic Table

The periodic table is arranged according to these sub-levels and the parts are referred as blocks.

Elements whose outer electrons occupy an s sub-level are in the s block.

Elements with outer electrons in the p orbitals make up the p block and so on for d and f blocks.

Periodic Table

The periodic table can be used to write electron configurations:

Ionization Energies… Again!

There is a general increase across a period when looking at first ionization energies.

As the number of protons in the nucleus increases it become more difficult to remove an electron.

There is then a decrease down to the next period.

Within the period there are small increases at the change between s and p sub-levels and half-way through the p sub-level.

Ion Configurations

Ions have different numbers of electrons from their uncharged atoms so the electron configuration must change too.

A positive ion has less electrons so the electron configuration will be shorter.

A negative ion has more electrons so the electron configuration will be longer.

Learning Check

Write the configurations of these ions:

1. Na+

2. Al3+

3. Cl-

4. O2-

Learning Check

Write the configurations of these ions:

1. Na+

1s22s22p6

2. Al3+

1s22s22p6

3. Cl-

1s22s22p63s23p6

4. O2-

1s22s22p6

unit 12: atomic structure ib chemistry ib topic 2 and 12

Documents

atom slide

chemistry slide

physics slide

electron slide

indivisible slide

relative number of atoms

idea of atoms

atoms positive charge