Atomic Theory

2.1.1 State the position of protons, electrons and neutrons in the atom2.1.2 State the relative masses and relative charges of protons, neutrons and electrons2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an element2.1.4 Deduce the symbol for an isotope given its mass number and atomic number2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge.2.1.6 Compare the properties of the isotopes of an element2.1.7 Discuss the uses of radioisotopes

History of the atomDemocritus (400 BC) suggested that the material world was made up of tiny, indivisible particles atomos, Greek for uncuttable

Aristotle believed that all matter was made up of 4 elements, combined in different proportions Fire - Hot Earth - Cool, heavy Water - Wet Air - Light

The atomic view of matter faded for centuries, until early scientists attempted to explain the properties of gases

Re-emergence of Atomic TheoryJohn Dalton postulated that: All matter is composed of extremely small, indivisible particles called atoms All atoms of a given element are identical (same properties); the atoms of different elements are different

3. Atoms are neither created nor destroyed in chemical reactions, only rearranged4. Compounds are formed when atoms of more than one element combine A given compound always has the same relative number and kind of atoms

Atoms are divisible!By the 1850s, scientists began to realize that the atom was made up of subatomic particlesThought to be positive and negativeHow would we know this if we cant see it or touch it?

Cathode Rays and ElectronsMid-1800s scientists began to study electrical discharge through cathode-ray tubes. Ex: neon signs Partially evacuated tube in which a current passes through Forms a beam of electrons which move from cathode to anode Electrons themselves cant be seen, but certain materials fluoresce (give off light) when energised

Oh there you are!JJ Thompson observed that when a magnetic or electric field are placed near the electron beam, they influence the direction of flow opposite charges attract each other, and like charges repel.The beam is negatively charged so it was repelled by the negative end of the magnet

http://www.chem.uiuc.edu/clcwebsite/video/Cath.movMagnetic field forces the beam to bend depending on orientationThompson concluded that: Cathode rays consist of beams of particles The particles have a negative charge

Thompson understood that all matter was inherently neutral, so there must be a counter A positively charged particle, but where to put it It was suggested that the negative charges were balanced by a positive umbrella-charge Plum pudding model chocolate chip cookie model

Rutherford and the Nucleus This theory was replaced with another, more modern oneErnest Rutherford (1910) studied angles at which a particles (nucleus of helium) were scattered as they passed through a thin gold foilhttp://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

Rutherford expected Rutherford believed that the mass and positive charge was evenly distributed throughout the atom, allowing the a particles to pass through unhindereda particles

Rutherford explained Atom is mostly empty space Small, dense, and positive at the center Alpha particles were deflected if they got close enougha particles

Nucleus: Containing protons and neutrons, it is the bulk of the atom and has a positive charge associated with it Electron cloud: Responsible for the majority of the volume of the atom, it is here that the electrons can be found orbiting the nucleus (extranuclear)The modern atom is composed of two regions:

Major Subatomic ParticlesAtoms are measured in picometers, 10-12 meters Hydrogen atom, 32 pm radius Nucleus tiny compared to atom If the atom were a stadium, the nucleus would be a marble Radius of the nucleus is on the order of 10-15 m Density within the atom is near 1014 g/cm3

NameSymbolChargeRelative Mass (amu)Actual Mass (g)Electrone--11/18409.11x10-28Protonp++111.67x10-24Neutronno011.67x10-24

Elemental ClassificationAtomic Number (Z) = number of protons (p+) in the nucleus Determines the type of atom Li atoms always have 3 protons in the nucleus, Hg always 80

Mass Number (A) = number of protons + neutrons [Sum of p+ and n] Electrons have a negligible contribution to overall mass

In a neutral atom there is the same number of electrons (e-) and protons (atomic number)

Nuclear SymbolsEvery element is given a corresponding symbol which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and atomic number

Find the number of protons number of neutrons number of electrons atomic number mass number

IonsCation is a positively charged particle. Electrons have been removed from the element to form the + charge. ex: Na has 11 e-, Na+ has 10 e-Anion is a negatively charged particle. Electrons have been added to the atom to form the charge.ex: F has 9 e-, F- has 10 e-

IsotopesAtoms of the same element can have different numbers of neutrons and therefore have different mass numbers

The atoms of the same element that differ in the number of neutrons are called isotopes of that element

When naming, write the mass number after the name of the element

How heavy is an atom of oxygen?There are different kinds of oxygen atoms (different isotopes) 16O, 17O, 18O

We are more concerned with average atomic masses, rather than exact ones Based on abundance of each isotope found in nature

We cant use grams as the unit of measure because the numbers would be too small Instead we use Atomic Mass Units (u) Standard u is 1/12 the mass of a carbon-12 atom Each isotope has its own atomic mass

Calculating AveragesAverage = (% as decimal) x (mass1) + (% as decimal) x (mass2) + (% as decimal) x (mass3) + Problem:Silver has two naturally occurring isotopes, 107Ag with a mass of 106.90509 u and abundance of 51.84 % ,and 109Ag with a mass of 108.90476 u and abundance of 48.16 % What is the average atomic mass?

Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u)= 107.87 amu

If not told otherwise, the mass of the isotope is the mass number in amu

The average atomic masses are not whole numbers because they are an average mass value

Remember, the atomic masses are the decimal numbers on the periodic tableAverage Atomic Masses

Properties of IsotopesChemical properties are primarily determined by the number of electronsAll isotopes has the same number of electrons, so they have nearly identical chemical properties even though they have different masses.Physical properties often depend on the mass of the particle, so among isotopes they will have slightly different physical properties such as density, rate of diffusion, boiling pointThe isotopes of an element with fewer neutrons will have:Lower masses faster rate of diffusionLower densities lower melting and boiling points

Calculate the atomic mass of copper if copper has two isotopes69.1% has a mass of 62.93 amuThe rest (30.9%) has a mass of 64.93 amu

Magnesium has three isotopes78.99% magnesium 24 with a mass of 23.9850 amu10.00% magnesium 25 with a mass of 24.9858 amuThe rest magnesium 26 with a mass of 25.9826 amuWhat is the atomic mass of magnesium?More Practice Calculating Averages

RadioisotopesIsotopes of atoms that have had an extra neutron attached to their nucleus.Carbon-14 radioactive decay is used to measures the date of objects.After 5700 years the amount of 14C will be half its original value.Iodine-125 or 131 is used to monitor the activity of the thyroid gland (b/c the thyroid tends to absorb iodine)

Cobalt-60 produces gamma rays (intense radioactivity) and is used in radiation treatment of cancer.

Note: gamma rays are the shortest wavelength on the electromagnetic spectrum. They are the most dangerous and difficult to shield from.

2.2 The Mass Spectrometer

Mass SpectrometerThe mass spectrometer is an instrument used:To measure the relative masses of isotopesTo find the relative abundance of the isotopes in a sample of an element

When charged particles pass through a magnetic field, the particles are deflected by the magnetic field, and the amount of deflection depends upon the mass/charge ratio of the charged particle.

Mass Spectrometer 5 StagesOnce the sample of an element has been placed in the mass spectrometer, it undergoes five stages.Vaporisation the sample has to be in gaseous form. If the sample is a solid or liquid, a heater is used to vaporise some of the sample.

X (s) X (g)

or X (l) X (g)

Mass Spectrometer 5 StagesIonization sample is bombarded by a stream of high-energy electrons from an electron gun, which knock an electron from an atom. This produces a positive ion:

X (g) X + (g) + e- Acceleration an electric field is used to accelerate the positive ions towards the magnetic field. The accelerated ions are focused and passed through a slit: this produces a narrow beam of ions.

Mass Spectrometer 5 StagesDeflection 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

Mass SpectrometerIf all the ions are travelling 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.

Mass SpectrometerDetection ions that reach the detector cause electrons to be released in an ion-current detectorThe 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.

Atomic Structure Mass SpectrometerIsotopes of boron

Ar of boron = (11 x 18.7) + (10 x 81.3) (18.7 + 81.3)

= 205.7 + 813 100 = 1018.7 = 10.2 100

m/z value 11 10Relative abundance %18.781.3

Mass Spectrometer QuestionsA mass spec chart for a sample of neon shows that it contains:90.9% 20Ne0.17% 21Ne8.93% 22Ne

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

Mass Spectrometer Questions90.9% 20Ne0.17% 21Ne8.93% 22Ne

Ar= 20.18u

Mass Spectrometer Questions

Calculate the relative atomic mass of leadYou must show all your work!

Mass Spectrometer Questions1.5% 204Pb23.6% 206Pb22.6% 207Pb52.3% 208Pb

Ar= 207.24=

2.3 Electron Arrangement2.3.1 Describe the electromagnetic spectrum2.3.2 Distinguish between a continuous spectrum and a line spectrum2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20

Electromagnetic radiation.

Electromagnetic RadiationMost subatomic particles behave as PARTICLES and obey the physics of waves.

Electromagnetic Radiation

Wavelengths and energyUnderstand that different wavelengths of electromagnetic radiation have different energies. Waves have a frequencyc=c=velocity of wave (2.998 x 108 m/s)=(nu) frequency of wave, units are cycles per sec=(lambda) wavelength

Electromagnetic SpectrumIn increasing energy, ROY G BIV

Electromagnetic SpectrumLong wavelength --> small frequencyShort wavelength --> high frequency

Bohrs ModelWhy dont the electrons fall into the nucleus?Move like planets around the sun.In circular orbits at different levels.Amounts of energy separate one level from another.

Bohr postulated that:Fixed energy related to the orbitElectrons cannot exist between orbitsThe higher the energy level, the further it is away from the nucleusAn atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)Think of Noble gases

Those who are not shocked when they first come across quantum theory cannot possibly have understood it.

(Niels Bohr on Quantum Physics)

Atomic Line Emission Spectra and Niels BohrBohrs greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the LINE EMISSION SPECTRA of excited atoms.Problem is that the model only works for HydrogenNiels Bohr(1885-1962)

How did he develop his theory?He used mathematics to explain the visible spectrum of hydrogen gasLines are associated with the fall of an excited electron back down to its ground state energy level.http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

Spectrum of White Light

The line spectrum electricity passed through a gaseous element emits light at a certain wavelengthCan be seen when passed through a prismEvery gas has a unique pattern (color)

Line Emission Spectra of Excited AtomsExcited atoms emit light of only certain wavelengthsThe wavelengths of emitted light depend on the element.

Spectrum of Excited Hydrogen Gas

Line Spectra of Other Elements

Line spectrum

Continuous line spectrum

Bohr also postulated that an atom would not emit radiation while it was in one of its stable states but rather only when it made a transition between states. The frequency of the radiation emitted would be equal to the difference in energy between those states divided by Planck's constant.

Ehigh-Elow= h = hc/h=6.63 1034 J s = Plancks constantE= energy of the emitted light (photon) = frequency of the photon of light = is usually stated in nm, but for calculations use m.

This results in a unique emission spectra for each element, like a fingerprint.electron could "jump" from one allowed energy state to another by absorbing/emitting photons of radiant energy of certain specific frequencies.

Energy must then be absorbed in order to "jump" to another energy state, and similarly, energy must be emitted to "jump" to a lower state. The frequency, , of this radiant energy corresponds exactly to the energy difference between the two states.In order for the emitted energy to be seen as light the wavelength of the energy must be in between 380 nm to 750 nm

Bohrs TriumphHis theory helped to explain periodic law (the trends from the periodic table)Halogens (gp.17 or group VII) are so reactive because it has one e- less than a full outer orbitalAlkali metals (gp. 1 or group I) are also reactive because they have only one e- in outer orbital

DrawbackBohrs theory did not explain or show the shape or the path traveled by the electrons.His theory could only explain hydrogen and not the more complex atoms

Energy level populationsElectrons found per energy level of the atom.The first energy level holds 2 electronsThe second energy level holds 8 electrons (2 in s and 6 in p)The third energy level holds 18 electrons (2 in s, 6 in p and 10 in d) There is overlapping here, so when we do the populations there will be some changes.

That is as far as this course requires us to go!

Examples for group 1Li 2.1 Na 2.8.1 K 2.8.8.1

The Quantum Mechanical ModelEnergy is quantized. It comes in chunks.A quanta is the amount of energy needed to move from one energy level to another.Since the energy of an atom is never in between there must be a quantum leap in energy.Schrdinger derived an equation that described the energy and position of the electrons in an atom

Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms.He developed the WAVE EQUATIONSolution gives set of math expressions called WAVE FUNCTIONS, Each describes an allowed energy state of an e-E. Schrodinger1887-1961Quantum or Wave Mechanics

Heisenberg Uncertainty PrincipleThe problem of defining nature of electrons in atoms was solved by W. Heisenberg.He observed that one cannot simultaneously define the position and momentum (= mv) of an electron.If we define the energy exactly of an electron precisely we must accept limitation that we do not know exact position.W. Heisenberg1901-1976

A good site:http://www.chemguide.co.uk/basicorg/bonding/orbitals.html

Electron ConfigurationHL only12.1.3 State the relative energies of s, p, d, and f orbitals in a single energy level12.1.4 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 px, py and pz orbitals12.1.6 Apply the Aufbau principle, Hunds rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z=54.

S orbitals1 s orbital forevery energy level 1s 2s 3sSpherical shapedEach s orbital can hold 2 electronsCalled the 1s, 2s, 3s, etc.. orbitals

P orbitalsStart at the second energy level 3 different directions3 different shapesEach orbital can hold 2 electrons

The D sublevel contains 5 D orbitalsThe D sublevel starts in the 3rd energy level 5 different shapes (orbitals)Each orbital can hold 2 electrons

The F sublevel has 7 F orbitalsThe F sublevel starts in the fourth energy levelThe F sublevel has seven different shapes (orbitals)2 electrons per orbital

SummaryStarts at energy level

Electron ConfigurationsThe way electrons are arranged in atoms.Aufbau principle- electrons enter the lowest energy first.This causes difficulties because of the overlap of orbitals of different energies.Pauli Exclusion Principle- at most 2 electrons per orbital - different spinsHunds Rule- When electrons occupy orbitals of equal energy they dont pair up until they have to .

Phosphorous, 15 e- to placeThe first to electrons go into the 1s orbitalNotice the opposite spinsonly 13 more

The next electrons go into the 2s orbitalonly 11 more

The next electrons go into the 2p orbitalonly 5 more

The next electrons go into the 3s orbitalonly 3 more

The last three electrons go into the 3p orbitals.They each go into separate shapes3 unpaired electrons1s22s22p63s23p3

Orbitals fill in order Lowest energy to higher energy.Adding electrons can change the energy of the orbital.Half filled orbitals have a lower energy.Makes them more stable.Changes the filling order

Write these electron configurationsTitanium - 22 electrons1s22s22p63s23p64s23d2Vanadium - 23 electrons 1s22s22p63s23p64s23d3Chromium - 24 electrons1s22s22p63s23p64s23d4 is expectedBut this is wrong!!

Chromium is actually1s22s22p63s23p64s13d5Why?This gives us two half filled orbitals.Slightly lower in energy.The same principal applies to copper.

Coppers electron configurationCopper has 29 electrons so we expect1s22s22p63s23p64s23d9But the actual configuration is1s22s22p63s23p64s13d10This gives one filled orbital and one half filled orbital.Remember these exceptions

Electronic Structure of transition metals With the transition metals it is the 4s electrons that are lost first when they form ions:Titanium (Ti) - loss of 2 e-1s2 2s2 2p6 3s2 3p6 4s2 3d2 Ti atomTi2+ ionCr atomCr3+ ion1s2 2s2 2p6 3s2 3p6 3d21s2 2s2 2p6 3s2 3p6 4s1 3d5 1s2 2s2 2p6 3s2 3p6 3d3 Chromium (Cr) - loss of 3 e-

Electronic Structure - QuestionsCopy and complete the following table:

Atomic no.Mass no.No. of protonsNo. of neutronsNo. of electronsElectronic structureMg121s2 2s2 2p6 3s2Al3+2710S2-1616Sc3+2145Ni2+3026

Ionization Energy12.1.1 Explain how evidence from first ionization energies across periods accounts for the existence of main energy levels and sub-levels in atoms12.1.2 Explain how successive ionization energy data is related to the electron configuration of an atom

Ionization EnergyThe amount of energy required to completely remove an electron from a gaseous atom.An atom's 'desire' to grab another atom's electrons. Removing one electron makes a +1 ion.The energy required is called the first ionization energy. X(g) + energy X+ + e-

Ionization EnergyThe second and third ionization energies can be represented as follows:X+ (g) + energy X2+ (g) + e-X2+ (g) + energy X3+ (g) + e-More energy required to remove 2nd electron, and still more energy required to remove 3rd electron

Group trendsIonization energy decreases down the group.

Going from Be to Mg, IE decreases because:Mg outer electron is in the 3s sub-shell rather than the 2s. This is higher in energyThe 3s electron is further from the nucleus and shielded by the inner electronsSo the 3s electron is more easily removedA similar decrease occurs in every group in the periodic table.

Notice any trends? Any surprises?

General trend: Increasing I.E. as we go across a periodLook at the peak at Mg and the plateau between P and S. Can you explain why?

Why is there a fall from Mg to Al?

Al has configuration 1s2 2s2 2p6 3s2 3p1, its outer electron is in a p sublevel Mg has electronic configuration 1s22s22p63s2. The p level is higher in energy and with Mg the s sub level is full this gives it a slight stability advantage

Why is there a fall from P to S?

This can be explained in terms of electron pairing. As the p sublevel fills up, electrons fill up the vacant sub levels and are unpaired.This configuration is more energetically stable than S as all the electrons are unpaired. It requires more energy to pair up the electrons in S so it has a lower Ionization energy.There is some repulsion between the paired electrons which lessens their attraction to the nucleus. It becomes easier to remove!

Driving ForceFull Energy Levels are very low energy.Noble Gases have full energy levels.Atoms behave in ways to achieve noble gas configuration.

2nd Ionization EnergyFor elements that reach a filled or half filled sublevel by removing 2 electrons 2nd IE is lower than expected.Makes it easier to achieve a full outer shellTrue for s2 Alkaline earth metals form +2 ions.

3rd IEUsing the same logic s2p1 atoms have an low 3rd IE.Atoms in the aluminum family form +3 ions.2nd IE and 3rd IE are always higher than 1st IE!!!

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