Topic 2 Atomic Structure. A Brief History 1807: John Dalton proposed atomic theory. – All matter was made up of small number of different kinds of atoms

Topic 2Atomic Structure

A Brief History• 1807: John Dalton proposed atomic theory.

– All matter was made up of small number of different kinds of atoms.

– Atoms are indivisible and indestructible, but could combine in whole numbers to form compounds.

• 1897: J. J. Thompson – When electricity is passed through a very low pressure tube,

rays from the negative electrode (Cathode) could be detected at the other end.

– Using magnets he bent the ray using magnets– Found that they had charge (neg.) and mass in a fixed ration

The electron was discovered

A Brief History Continued• 1909: Ernest Rutherford

– Fired alpha particles through gold foil only a few atoms thick.

– He expected them to only deflect in small angles according to Thomson’s model.

– Because some particles bounced back, he developed a different model which has a dense, positively charged nucleus as the core.

Brief History Con’t• Ernest Rutherford

• Accelerated alpha α particles (helium nuclei) at thing gold (Au) foil

• Helium nuclei is the center portion of an He molecule without electrons.(e-)

• Analogy: Like shooting a howitzer through a piece of paper

• They did something odd.

• Summary• alpha α particles are He

atoms without e-

• e- = electrons• Au = gold• He = Helium

• Drawing

Flash Animation

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf



Models• Plum pudding

• Negative ‘chocolate chips’ were thought to be surrounded by positive ‘cookie dough’

• Nuclear atom

• All mass around a small positive center. Neg. Electron cloud around it

IB Core Objective

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

• State: Give a specific name, value or other brief answer without explanation or calculation.

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

• Subatomic particles: Proton, neutron, and electron.• Nucleus: The center portion of the atom.

– Contains protons and neutrons

• Electrons: Occupy shells around the nucleus.• Orbital: Where an electron MIGHT be found• If the nucleus is one meter across, then the electrons would be about 10 kilometers across. • Much of the atom is empty space.

IB Core Objective• 2.1.2 State the relative mass and

relative charge of protons, electrons and neutrons.

• State: Give a specific name, value or other brief answer without explanation or calculation. (Obj 1)

2.1.2 State the relative mass and relative charge of protons, electrons and neutrons.

• Nucleus contains all the positive charge and nearly all the mass (>99.9%).

• amu stands for atomic mass unit

Symbol Mass Charge

N0 1 amu No charge

P+ 1 amu +1

e- 5 * 10-4 amu -1

IB Core Objective

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

• Define: Give the precise meaning of a word, phrase or physical quantity. (Obj 1)

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

• Atomic number (Z) is the number of protons

• Mass number (A) is the sum of the # of P+ and N0 (the nucleus)

• When writing these values, use the standard notation format below:

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

Isotopes• Many elements are composed of slightly

different types of atoms.• These are called isotopes.• Isotopes all have the same number of protons,

but differ in the number of neutrons.

IB Core Objective

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

• Deduce: Reach a conclusion from the information given. (Obj 3)

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

Using the Periodic Table• Periodic table helps us find

information on different elements.

• Atomic Mass: Tells you the AVERAGE weight of the atom

Atomic Number: Tells you number of protons


The following notation should be usedAX

Z= Number of protonsA=Atomic mass number (protons + neutrons)

Z


• Deduce the symbols using standard notation for the following elements:

H Cl

C Fe

IB Core Objective

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

• Calculate: Find a numerical answer showing the relevant stages in the working (unless instructed not to do so).

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

number, atomic number and charge.• Number of protons = atomic

number• Number of electrons =

Number of protons– If the atom is neutral (no

charge)• Number of neutrons =

(Atomic mass) – (Number of protons)


number, atomic number and charge.Calculating Neutrons

• The atomic weight MUST BE rounded. – 1.3 is rounded down 1– 1.6 is rounded up 2– 13.6 – 35.45 – Now you can use this weight to find the number of

neutrons!

14

35


number, atomic number and charge.

• What is the approximate number of neutrons in iron?

• 30 (remember, the atomic mass is an average based on the abundance (%) of isotopes.)

IB Core Objective

• 2.1.6 Compare the properties of the isotopes of an element.

• Compare: Give an account of similarities and differences between two (or more) items, referring to both (all) of them throughout.(Obj 3)

2.1.6 Compare the properties of the isotopes of an element.

• Carbon dating is done using the isotope C-14. The 14 is the atomic weight.

• # of protons =• # of electrons = • # of neutrons =

• What is the difference from the carbon on the periodic table?



• Elements may have multiple isotopes each with a different abundance that all contribute to the relative atomic mass of the element.

• Calculating Relative atomic mass:– 37Cl & 35Cl have relative abundances of 25% and

75% respectively, what is the relative atomic mass?

IB Core Objective

• 2.1.7 Discuss the uses of radioisotopes

• Discuss: Give an account including, where possible, a range of arguments for and against the relative importance of various factors, or comparisons of alternative hypotheses. (Obj 3)

2.1.7 Discuss the uses of radioisotopes

• Radioisotopes are isotopes which are radioactive.

• To be radioactive means the nucleus is unstable, and will emit subatomic particles or gamma rays (electromagnetic radiation).

2.1.7 Discuss the uses of radioisotopesHomework

Go onto the uaschemistry.pbworks site.In the 11th grade folder, go to the “Uses of

Radioisotopes” pageResearch a radioisotope that we use, and comment

what it is and how we use it on this page.If something has already been commented on, it cannot

be repeated, so better to start early!You can expand on someone else’s comments.These should be done by Tuesday, November 3rd

IB Core Objective

• 2.2.1 Describe and explain the operation of a mass spectrometer.

• Describe: Give a detailed account. (Obj 2)• Explain: Give a detailed account of causes,

reasons or mechanisms. (Obj 3)

2.2.1 Describe and explain the operation of a mass spectrometer.

• Sequence for Mass Spectroscopy

• 1) Vaporize• 2) Ionize• 3) Accelerate• 4) Deflect• 5) Detect• This needs to be done

in a vacuum to prevent molecules colliding.

IB Core Objective• 2.2.2 Describe how the mass

spectrometer may be used to determine relative, atomic masses using the 12C scale.

• Describe: Give a detailed account. (Obj 2)

2.2.2 Describe how the mass spectrometer may be used to determine relative, atomic masses

using the 12C scale.

Relative Atomic Mass (Ar)

The weighted mean mass of all naturally occurring isotopes of that element

relative to one twelfth of carbon-12.

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

scale.

• Deflection – The accelerated ions are deflected into the magnetic field. The amount of deflection is greater when:

• the mass of the positive ion is less• the charge on the positive ion is greater• the velocity of the positive ion is less• the strength of the magnetic field is greater


scale. • If all the ions are traveling at the same velocity and

carry the same charge, the amount of deflection in a given magnetic field depends upon the mass of the ion.

• For a given magnetic field, only ions with a particular relative mass (m) to charge (z) ration – the m/z value – are deflected sufficiently to reach the detector.


scale.

• Detection – ions that reach the detector cause electrons to be released in an ion-current detector.

• The number of electrons released, hence the current produced is proportional to the number of ions striking the detector.

• The detector is linked to an amplifier and then to a recorder: this converts the current into a peak which is shown in the mass spectrum.

IB Core Objective

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

• Calculate: Find a numerical answer showing the relevant stages in the working (unless instructed not to do so).

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

data.• Determine the

relative atomic mass from this data.

• First, what would your estimate be?

282930

z

z


data.

(28 x 92.21) + (29 x 4.79) + (30 x 3.09)100

=28.1


data.• A mass spec chart for a sample of neon shows that it

contains:

– 90.9% 20Ne

– 0.17% 21Ne

– 8.93% 22Ne

Calculate the relative atomic mass of neonYou must show all your working!

Ar= 20.18

IB Core Objective

• 2.3.1 Describe the electromagnetic spectrum.

• Describe: Give a detailed account. (Obj 2)

2.3.1 Describe the electromagnetic spectrum.

2.3.1 Describe the electromagnetic spectrum

The Chemical Fingerprint for Atoms• Energy levels are discrete levels where

electrons can exist

• When they are excited, elements will emit a characteristic color of light.

• What can we do to excite them?

http://jersey.uoregon.edu/vlab/elements/Elements.html




IB Core Objective

• 2.3.2 Distinguish between a continuous spectrum and a line spectrum.

• Distinguish: Give the differences between two or more different items.

2.3.2 Distinguish between a continuous spectrum and a line spectrum.

IB Core Objective

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

• Explain: Give a detailed account of causes, reasons or mechanisms. (Obj 3)

Chemical Finger Print

• When electrons fall from Excited State to ground state, a photon is emitted

•Every atom has a different number therefore each atom has its own fingerprint

Higher energy level(Farther from the

nucleus)Lower energy level

(Closer to the nucleus)

n =4

n =3

n =2

n =6

n =5

Atom nucleus has been omitted

Energy levels (ring proximity) not accurately

represented

n = 5n = 4n = 3

n = 2

n = 1

Energy in but not

enough to ionize

e- falling to the fist shell

(n =1) emit ultraviolet energy

‘Lyman series’

e- falling to the 2nd shell

emit visible light energy

‘Balmer series’

e- falling to the 3rd shell emit

Infrared energy ‘Paschen series’

IB Core Objective

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

• Deduce: Reach a conclusion from the information given. (Obj 3)

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

• Most stable energy levels are closest to the nucleus.

• Electrons fill the most stable levels before filling the higher ones.

• There is a maximum number each energy level can hold. The first holds two electrons, the second holds eight.

• The electrons in the valence shell determine the physical and chemical properties of the element.

Electron Arrangement

• The Bohr model (2,8,8,2)

• Li Be B CNe

IB HL Objective

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


12.1.1 Explain how evidence from first ionization energies across periods accounts for the existence of

the main energy levels and sub-levels in an atom.

• You will be explaining ionization energies to the class on Tuesday.

• You will all need to participate in helping teach this to the other students.

• You can pick how you teach them (as a class, each of you work with a small group, etc.)



Ionization Energy• Energy required to rip off one electron from an

element in its gaseous state.X(g) X+

(g) + e-



First Ionization EnergyThe energy needed to remove the first electron

from a neutral atom.

Na(g) → Na+(g) + e-



Second Ionization EnergyHmmm, what do you think this might be?

The energy needed to remove the second electron.

Na+ → Na2+(g) + e-

IB HL Objective

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




• Look at Table 7.2 on page 271 of your book.• Work with a partner(s) and explain the trends

in ionization energies that you see, and why you see them

IB HL Objective

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

• State: Give a specific name, value or other brief answer without explanation or calculation. (Obj 1)

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

• Electrons are arranged around the nucleus in specific energy levels and sub-levels.

• The different sub-levels differ in the shape of electron distribution.


S Orbitals• Look at page 231 at figure 6.16.• What is it showing?• What is it’s shape?


• All s orbitals are spherical and symmetric.• 1s orbital is the lowest-energy orbital.

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Topic 2 Atomic Structure. A Brief History 1807: John Dalton proposed atomic theory. – All matter was made up of small number of different kinds of atoms