1 Atomic Structure Periodic Table - B - Atomic Theory

8/17/2019 1 Atomic Structure Periodic Table - B - Atomic Theory

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Atomic Structure & the Periodic Table

B) Atomic Theory

CHE101 Chemistry Unit 1

2015 – 2016Ms. T. Jackson

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Topics

• The Scientific Process & Theoretical Change

• Atomic Theory

– Dalton 1807

–

Thomson 1897 – Rutherford 1909

– Bohr 1913

• Atomic Structure

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Atomic Theory

• An atom is the smallest unit of matter

– From the Greek ‘atomos’ meaning “cannot be split”

– Atoms are basic building blocks of chemistry

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HO


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Atomic Models / Theories

• Dalton 1807

• Thomson 1897

• Rutherford 1909

•

Bohr 1913

• For each model/theory, describe:

– The model

– Experiments that produced the model

– Limitations of the model

– How the limitations were resolved by the next model/theory

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Atomic Theory Development Timeline

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1807Dalton

•theorised

atoms are

“hard spheres”

1897

Thomson

•

electrons detected•“plum pudding” model

with embedded electrons

1800 1900 2000

1900

Goldstein

•protons detected

1909

Rutherford

•“planetary model”

•electrons orbit nucleus

1926

Schroedinger

•wave equation

for electron orbitals

1913

Bohr

•quantum theory model

•electron energy levels

quantised

1913

Mosely•neutron existence

hypothesised1932

Chadwick

•neutrons detected1924

deBroglie

•theorised wave-particle

duality


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Dalton 1807“Hard Spheres”

• Conceptualized atoms as “hard spheres”

• Assumptions of Dalton’s atomic theory:

1. Matter consists of tiny particles called atoms, which areindestructible & indivisible (i.e. cannot be split)

2. All atoms of the same element are identical in mass andchemical properties. They differ from the atoms of all otherelements.

3. Atoms can combine in simple whole number ratios to formcompounds

4. A chemical reaction consists of rearranging atoms from onecombination into another. Atoms are not created, destroyed,or broken into smaller pieces by any chemical reaction. i.e.theindidvidual atoms remain intact

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Dalton 1807

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H O

Hard spheres with

different masses &

chemical properties


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Dalton 1807Evidence for Dalton’s Model

1. a) The particulate theory of matter explains the observed

differences in behaviour of solids, liquids, and gases in terms

of packing, motion, etc.

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solid liquid gas


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b) Diffusion experiments support the particulate nature of

matter. E.g.

i. purple KMnO4 crystals placed into a beaker of water purple

solution

ii. chlorine mixing with air in a gas jar although it is denser than

air

iii. random (Brownian) motion of smoke particles observed in a

smoke cell

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2. The previously theorised law of conservation of mass: “the

total mass remains constant during a chemical reaction”

was supported by Dalton’s atomic theory.

Since every atom has a definite mass (according to postulate2) & a chemical reaction should only rearrange the chemical

combinations of atoms (postulate 4), the mass must remain

constant.

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E.g. mercury + oxygen product

2.53 g + 0.20 g 2.73 g


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3. The previously theorised law of definite proportions (or

constant composition): “a pure compound, whatever its

source, always contains definite or constant proportions of

the elements by mass”

From postulate 3, a compound is a type of matter containing

the atoms of two or more elements in definite proportions.

Because atoms have definite mass, compounds must have

the atoms in definite proportions by mass.

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E.g. 1.0000 g of sodium chloride always contains

0.3934 g sodium and 0.6066 g chlorine


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4. Dalton’s atomic theory predicted the law of multipleproportions: “when two elements form more than onecompound, the masses of one element in these compounds for a fixed mass of the other element are in ratios of smallwhole numbers”

Deducing this law from atomic theory helped to convincechemists of the validity of the theory.

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E.g. carbon + oxygen compound 1 + compound 2

1.000 g excess 1.000 g C : 1.000 g C :1.3321 g O 2.664 O

CO + CO2


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Dalton 1807

Inaccuracies/Problems with Dalton’sModel

• Atoms were later found to consist of further particles i.e. they

could be divided into smaller parts

• Atoms can be destroyed by nuclear reactions

• Atoms of the same element can have different masses -

isotopes

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Thomson 1897“Plum Pudding” Model

• “Plum-pudding” model with

electrons embedded in a sea of positive charge

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n+


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• Discharging electricity through gases at low pressure – Cathode rays deflected by electric and magnetic fields

– Charge/mass ratio e/m = -1.76 x 1011 Ckg-1 same regardless of type of gas orelectrode used

– Same negatively charged particles (electrons) present in all matter

– Atoms are not indivisible. They contain charged sub-atomic particles 15

1) 2 electrodes from

high voltage source

sealed into an evacuated

glass tube

2) high voltage current

turned on

3) beam of rays given

off by negatively charged

electrode (cathode)

4) cathode rays deflected

by electric field towards

positively charged plate

Thomson 1897Experimental Evidence


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Thomson 1897Inaccuracies/Problems with the Model

• This model could not explain the deflection of (alpha)

particles by metal foil (Later discovered that most of the atom

is actually empty space with electrons ‘orbiting’ the positively

charged nucleus) [see Rutherford model]

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Rutherford 1909“Planetary” Model

• “Planetary model” with negatively charged electrons orbiting

a positively charged nucleus

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n

+n+

electron

nucleus


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Rutherford 1909Experimental Evidence

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•Bombarding thin metalfoils (e.g. gold) with alphaparticles

– Most passed through withno interaction

– A few ~ 1 in 8000 were

scattered at large angles;some were sentbackwards towards thesource

– Atom mainly emptyspace with mass &

positive chargeconcentrated (>99.5%) intiny central nucleus whilenegatively chargedelectrons orbit thenucleus (like planetsorbiting the sun)


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Rutherford 1909Experimental Evidence (Geiger & Marsden)

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• Illustration of theorisedexplanation for scatteringof particles by the nuclei

of metal atoms

– Only particles thatcollide with the positivelycharged nucleus aredeflected

– The vast majority passthrough the spacesbetween the nuclei


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Rutherford 1909Inaccuracies/Problems with the Model

• If electrons are negatively charged, and the nucleus is

positively charged, why don’t the electrons spiral into the

nucleus?

• Could not explain atomic and emission spectra, i.e. why do

atoms absorb or emit light of certain frequencies? [see Bohr

model]

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Bohr 1913Quantization of Electronic Energy Levels

• Quantum theory based model

• Max Planck had recently suggested that in certain systems

energy can be absorbed or emitted in certain specific

amounts i.e. in separate packets of energy called ‘quanta’

• Bohr applied this to the atom & postulated the existence of

discrete energy levels within the atom, i.e. electrons can only

orbit the nucleus at certain distances depending on energy

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electron

nucleus


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Bohr 1913Features of the Bohr Model

• Electrostatic force between nucleus & orbiting electronscancelled out by outward force due to orbital motion [soelectrons do not spiral into the nucleus]

• Electron in a given orbit can only have a certain amount of

energy (i.e. the energy is quantized) & the orbit can only havea certain radius

• If the electron absorbs a “quantum” of energy (a photon oflight energy), it moves to an orbit with a higher energy levelthat is further away from the nucleus. It is in an “excited”

energy state• An excited electron emits energy to return to its stable

“ground state” orbit

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• For an electron to move from an orbit E1 to one of energy E2,

the light absorbed must have a frequency given by Planck’s

equation:

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h = E2 – E1

or E = h (Also, c = l

So, DE = hc/l)

where = frequency

h = Planck’s constant = 6.63 x 10-34

Jsc = speed of light (or electromagnetic

radiation) in vacuo = 3 x 108 ms-1

l = wavelength

Orbit of energy E2Orbit of energy E1

Energy emitted = E2 – E1

Frequency of light emitted

= = (E2 – E1)/h

Energy absorbed = E2 – E1

Frequency of light emitted

= = (E2 – E1)/h


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• Bohr assigned quantum numbers to the orbits

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Orbit of lowest energy (closest to the nucleus) n = 1

Next E level n = 2

etc.

n = 1

n = 2n = 3n = 4n = 5

n = 1

n = 2

n = 3n = 4

ionized

Energy

Energy Level Diagram


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Bohr 1913Experimental Evidence

• Atomic Absorption and Emission Spectra – If white light is shined through a prism, a continuous spectrum is

produced containing all visible wavelengths or frequencies

– In contrast if the light is shined through the sample of an element,radiation is absorbed at certain frequencies producing a discontinuous

coloured spectrum (an atomic absorption spectrum) – The fact that the same discontinuous absorption spectrum is obtained

for a particular element supports the theory that electronic energylevels are quantized

– Similarly, if electrical or thermal energy is passed through a gaseoussample of an element e.g. if the sample is heated to a sufficiently hightemperature and then allowed to cool, radiation is emitted only atcertain frequencies producing a discontinuous discrete line spectrum(an atomic emission spectrum)

– The fact that the same discrete lines are seen in the atomic emissionspectra for a particular element supports the theory that electronicenergy levels are quantized

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Bohr 1913Experimental Evidence – White Light

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all colours seen

continuous spectrum


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Bohr 1913Experimental Evidence – Atomic Absorption Spectra

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Sample

Discontinuous

black lines on a

bright background

absorption spectrum


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E.g. Na in a flame yellow flame

E.g. H2 gas tube excited with

electric discharge reddish-pink

glow

Bohr 1913Experimental Evidence – Atomic Emission Spectra

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coloured lines on a

black background

emission spectrum


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Hydrogen Emission Spectrum

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•Hydrogen’s emission spectrum has been extensively studied

• Lines in different regions of the electromagnetic spectrum

• Lines named after discoverers [note: some series overlap]

• In each series, lines become closer together as frequency

increases until at high frequencies the lines coalesce orconverge to form a continuum of light


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Hydrogen Emission Spectrum

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n = 1

n = 2

n = 3n = 4

E

Energy Level Diagram

n = 5

Emission

Spectrum Lyman

series

Balmer

series

DE = h

for each line

• Hydrogen emission spectrum

transitions:

– Lyman: UV

transitions to n = 1 (ground state)

from any other orbital

– Balmer: visible light

transitions to n = 2

–Paschen, Brackett, Pfund, IRto n = 3, n = 4, n = 5 respectively


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Bohr 1913Inaccuracies/Problems with the Bohr Model

• Explained the emission spectrum of a simple atom like

hydrogen, but failed to explain the spectra of more advanced

atoms. (Resolved by Schroedinger’s wave equations in 1926)

[ See later]

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Other Contributions

• Goldstein (1900) & Moseley (1913) – detected protons as separate entities – theorised that since atoms are neutral # protons = # electrons

– theorised that since the mass of the electron is negligible and themass of the protons adds up to less than the total mass of the atom

neutrally (zero) charged particles (i.e. neutrons) must exist

•Chadwick (1932) – experimental detection of the neutron

– bombarded Be with alpha particles

– produced a stream of neutrally charged particles that had enoughmass to pass through several cm of solid lead neutrons detected

•

Schroedinger (1926) [see later] – applied deBroglie’s wave-particle duality theory to electrons

– atomic orbitals described in terms of probability densities

– suborbitals of electrons theorised

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References

• Chemistry for CAPE by Susan Maraj and Arnold Samai. Caribbean

Educational Publishers 2011

• A-Level Chemistry by E. N. Ramsden. Nelson Thornes Ltd 2000

• CAPE Chemistry: A Caribbean Examinations Council Study Guide Unit 1

by Norris et al. Nelson Thornes Ltd 2012

• Oxford Revision Guides AS & A Level Chemistry Through Diagrams by

Michael Lewis. Oxford University Press 2012

• General Chemistry by Ebbing and Gammon. Houghton Mifflin Company

2005

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