Atomic Structure

GChem Chapter 3

Learning objectives1. Discuss the history of the current atomic model, including contributions of:

Dalton, Thomson, Rutherford, Moseley, Bohr.

2. Explain the laws of conservation of mass, definite proportions, and multiple proportions.

3. Describe the location, mass and charge of the three components of an atom.

4. Explain and calculate the atomic mass, atomic number and charge for any given atom.

5. Explain the concept of isotopes and calculate the average atomic mass for an element given natural abundances.

6. Write the AZX symbol for any given isotope.

7. Discuss the development of the quantum mechanical model of the atom.

8. Define the 4 quantum numbers.

9. Write the electron configuration and orbital diagram for any given element or ion.

10. Relate electron configuration to the arrangement of the periodic table.

The story of mankind’s search for the basic structure of matter stretches throughout the entire history of the human race.

People have always wondered about what things are made of and what makes one substance different from another.

The history of the atomic model

Ancient times- Around 400 BC, Greek philosophers developed the idea that all matter must be composed of tiny, indivisible particles which were termed “atoms”. They suggested that atoms of different substances were different and that atoms “hooked together” to form large scale matter.

Democritus is usually credited with the development of this idea.

The next 2208 years – Not much changed.

Throughout most of recorded human history, the Greek’s idea of “atoms” as the basis of matter was accepted but without any real experimental evidence.

Although much study was done during this time, little of it was rigorous and much was given mystical explanations – This was Alchemy. (the emerald tablet)

~ 1800 – More rigorous experimentation was done by various scientists. The results of these experiments were formulated into 3 laws:

The law of conservation of mass - Mass is not gained or lost in chemical reactions.Total mass of reactants = Total mass of products

The law of definite proportions - A compound will always be composed of the same percentage of each element by weight.

Examples:100g of NaCl → 39.4g Na and 60.6g Cl10g of NaCl → 3.94g Na and 6.06g Cl

CO2 will always be composed of 72.7% O and 27.3% C.

The law of multiple proportionsIf two compounds contain the same elements, a comparison of the mass of one element that reacts with a fixed mass of the other element will give a factor of a small whole number. Huhh??

Examples:H2O → 1g H and 7.94g O

H2O2 → 1g H and 15.88g O

A 2 to 1 ratio!!!!

Compound % O % N gO / gN Ratio

N2O 36.35 63.65 0.571 1

NO 53.32 46.68 1.142 2

NO2 69.56 30.44 2.284 4

The conclusion was that elements exist in “chunks” of mass. That for a given element, all chunks have the same mass.

John Dalton used the 3 laws to develop his atomic theory in 1808. This was the first theory that described the composition of matter based on experimental evidence.

Dalton’s Theory

1. Matter is made of indestructible atoms.2. All atoms of an element are identical.3. Atoms of different elements have different

properties.4. Atoms of different elements combine with

each other in whole number ratios.5. Chemical reactions are re-arrangements of

atoms.

Dalton’s Model – Billiard Ball Model

• J.J. Thomson- In a series of experiments with a cathode ray tube (CRT) in 1897, discovered that negatively charged particles of matter could be removed from atoms.

• This indicated that atoms were not indivisible but were composed of even smaller particles.

• The discovered particle was the electron.

The Plum Pudding Model - JJ Thomson 1904

Ernest Rutherford – 1911 – Proposed a new model of the atom based on the results of the Gold Foil experiment.

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

Rutherford Model of the atom

• Rutherford suggested that the atom is composed of a very small, dense, positively charged nucleus surrounded by an area of empty space containing the atom’s electrons.

Henry Moseley – 1913 – Discovered that the number of positive charges in an atom is equal to the element number. This indicated that there was a particle in the nucleus that was the source of + charge. In 1920, Rutherford named the particle “proton”.

Niels Bohr – 1913 – Considered that the Rutherford model of the atom was unstable and the spectra of atoms (discrete bands of light absorbed and emitted by atoms) to propose a new model with the electrons confined to specific energy levels (sometimes called shells or orbits).

• e- exist in specific energy levels in atoms and cannot exist between energy levels – The levels are quantized.

• e- can absorb specific amount of energy to jump to a higher energy level or release energy to drop to a lower level – quantum jumps.

• e- in lowest energy level = ground state e- jumps to higher energy level = excited state

• Specific numbers of e- can reside in each energy level.

The Bohr Model

• http://science.sbcc.edu/physics/flash/siliconsolarcell/bohratom.swf

• James Chadwick – 1932 – Explained the difference between the observed mass of atomic nuclei and the number of + charges (also considering spin) by proposing the presence of particles with masses similar to those of protons but with no charge -the neutron.

We now have a workable model of the atom: It is composed of three particles, proton, neutron and electron.

Name Symbol Mass Charge Location

Proton p+ 1.67x10-27 kg + nucleus

Neutron n0 1.67x10-27 kg None Nucleus

Electron e- 9.1x10-31 kg - Energy levels

The mass of a neutron is actually very slightly more than that of a proton however, in chemistry we generally consider them to be the same.

We use a convenient unit to express this mass, the amu (atomic mass unit).

An amu is defined as 1/12 the mass of a carbon-12 atom.

We generally consider the masses of both p+ and n0 to be 1 amu.

The mass of an e- is so much less than the other particles that we considered it to be zero in calculating the mass of an atom.

So the mass of an atom, in amu’s, is simply the number of protons plus the number of neutrons. This is sometimes called “mass number”atomic mass = #p+ + #n0

The total charge on an atom is determined by the number of p+ and e-. Since these particles have charges of equal magnitude and opposite sign, their charges cancel. When an atom has the same number of p+ and e- , the total charge must be zero – a neutral atom.

• An ION is a form of an atom where the number of e- does not match the number of p+.

• If the ion has more e- than p+ it will have a total negative charge.

• If the ion has less e- than p+ it will have a total positive charge.

Ex: O ion has 8 p+ and 10 e- . 2 more electrons than protons so the ion will have a charge of -2, (O-2)

Mg ion has 12 p+ and 10 e- . 2 more protons than electrons so the ion will have a charge of +2, (Mg+2)

Br ion has 35 p+ and 36 e- . 1 more electrons than protons so the ion will have a charge of -1, (Br-)

La ion has 57 p+ and 54 e- . 3 more protons than electrons so the ion will have a charge of +3, (La+3)

Ions can only form by the loss or gain of electrons.

Since protons determine the identity of the element, ions are never formed by losing or gaining p+.

This should make sense – Electrons are flying around the outer part of the atom – easy to gain or lose. Protons are locked into the nucleus at the center – cannot gain or lose.

• Writing symbols for atoms – the AZX method• X – the chemical symbol for the element, H, O, Ca• A – the mass number (atomic mass)• Z – the atomic number (number of p+)

• Additionally, the charge on an ion can be written in the upper right!

Examples: Write the AZX notation for the following:

1. 20 p+, 20 n0, 20 e-

2. 15 p+, 16 n0, 15 e-

3. 26 p+, 30 n0, 23 e-

4. 35 p+, 44 n0, 36 e-

• Summary:1. the number of protons is the atomic number and tells

you what element you have.13 p+ = Al

2. The number of electrons can be compared to number of protons to tell you if you have a neutral atom or an ion.

13 p+ and 10 e- = Al+3

3. The number of neutrons is added to number of protons to give the mass number (atomic mass).

13 p+ and 10 e- and 14 n0 = Al+32713

• The number of neutrons that atoms of a given element contains can vary.

• Differing numbers of neutrons do not change the identity of the element, however they will change the mass of the atom.

• Isotopes are versions of the same element with different numbers of neutrons and therefore different atomic masses.

Isotope p+ n0 Atomic massC-12 6 6 12

C-14 6 8 14

Mg-24 12 12 24

Mg-25 12 13 25

U-235 92 143 235

U-238 92 146 238

Isotopes of hydrogen

Practice:Write the AZX symbols for the isotope with:

21 p+ and 24 n0

17 p+ and 20 n0

82 p+ and 125 n0

41 p+ and 52 n0

The atomic mass for an element given on the periodic table is the weighted average of the masses of all the naturally occurring isotopes for that element.

The average takes into account both the mass of the isotope and it’s natural abundance (what percentage of it occurs in nature).

• Pb has 4 naturally occurring isotopes:• Atomic Number = 82 so 82 p+ in each• Pb-204 204-82 = 122 no

• Pb-206 206-82 = 124 no

• Pb-207 207-82 = 125 no • Pb-208 208-82 = 126 no

Isotope Natural AbundancePb-204 1.4%Pb-206 24.1%Pb-207 22.1%Pb-208 52.4%

To calculate Avg At Mass: A weighted average!204 x 0.014 = 2.856 multiply the isotope mass206 x 0.241 = 49.65 by it’s natural abundance207 x 0.221 = 45.75 (move the decimal point 2208 x 0.524 =108.99 places)

207.25 add the results!

• Average atomic mass practice:Uranium has three common isotopes. If the

abundance of 234U is 0.01%, the abundance of 235U is 0.71%, and the abundance of 238U is 99.28%, what is the average atomic mass of uranium?

Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%), 48Ti (73.4%), 49Ti (5.5%), 50Ti (5.3%). What is the average atomic mass of titanium?

Electron ConfigurationQuantum Mechanical View

The Bohr model is very useful but has real limitations and in some ways does not agree with observations.

Suggested that electrons could be considered as particles or waves – just as Einstein had proposed about light.

Only certain “orbits” are possible because these correspond to multiples of the wavelength for the electron. The orbit can be thought of as a standing wave.

Louis DeBroglie : In 1924 proposed a new model to explain the problems with Bohr’s model.

Erwin Schrodinger: In 1926, derived an equation that described the position of an electron as a probability. The equation can be used to produce a set of four quantum numbers for each electron that describes its probable location in an atom.

Wolfgang Pauli – 1925 – The Pauli exclusion principle states that no two electrons in an atom can have the same set of 4 quantum numbers.

• Werner Heisenberg – 1927. The Heisenberg uncertainty principle states that the uncertainties of the position and momentum of an electron are inversely proportional.The more accurately that one is known, the less accurately the other can be known.

The Cat HotelRule 1: Cats are lazyRule 2: Cats do not like

each other

5th

4th

3rd

2nd

1st

• A set of 4 quantum numbers for an electron in an atom gives the “address” for that electron.

• Where an e- resides in an atom is called an orbital. Each energy level contains 1 or more orbitals.

• The quantum numbers describe the energy level, the shape and orientation of the orbital and the spin of the e-.

L 3

L 2

L 1

• Quantum Numbers are used to describe the location of an electron in an atom.

• Four quantum numbers are needed for each electron and no electrons in an atom can have the same set on QN’s.

• The principal QN is is identified by the letter n and gives the main energy level of the electron.

• The principal QN can have the values 1,2,3,4,5……• The second QN is identified by the letter l and gives the

shape of the electron orbital.• The l QN can have the values 0,1,2,3…n-1• The m l QN gives the orientation of the orbitals and can

have the values - l …-3,-2,-1,0,1,2,3…+ l • The ms QN gives the spin of an electron and has the values

of +1/2 or -1/2.

For N=1 l = _____ This is an _____ orbital which holds _____ e-

For N=2 l = _____ This is an _____ orbital which holds _____ e-

And l = _____ These are _____ orbitals which hold _____ e-

For N=3 l = _____ This is an ______ orbital which holds _____ e-

And l = _____ These are _____ orbitals which hold _____ e-

And l = _____ These are _____ orbital which hold _____ e-

For N=4 l = _____ This is an _____ orbital which holds _____ e-

And l = _____ These are ______ orbitals which hold ____ e-

And l = _____ These are ______ orbital which hold _____ e-

And l = _____ These are ______ orbitals which hold _____ e-

• Orbital diagrams:

• Translates to electron configurations1s 2s 2p 3s 3p 4s 3d 4p 5s

1s 2s 2p 3s 3p

• To write orbital diagrams and electron configurations:

1. Decide how many electrons the atom has. For neutral atoms this is the same as the atomic number. For + ions, 1 e- is lost for each + charge. For – ions 1 e- is gained for each - charge.

2. Starting with the 1st energy level (n=1) add 2 e- to each orbital until all are used up. Remember, for sets of orbitals (p,d,f) e- do not double up until they have to.

3. Use a filling diagram to fill in the correct order.

Write the orbital diagram and e- configuration for the following: O

Si

Ti

Sn

Pm

IonsGround state versus excited state

Shortcut for writing e- configurationsFind the noble gas that has a lower atomic number that is nearest to the element.Write the symbol for the noble gas in brackets.Write the remainder of the e- configuration between the noble gas and the element:Ex: Cd1s22s22p63s23p64s23d104p65s24d10 The underlined is the configuration for Kr so we can write:[Kr]5s24d10

Additional Content

• Quarks!