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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Dalton 1807Evidence for Dalton’s Model
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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Dalton 1807Evidence for Dalton’s Model
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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Dalton 1807Evidence for Dalton’s Model
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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Dalton 1807Evidence for Dalton’s Model
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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Bohr 1913Features of the Bohr Model
• 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 1913Features of the Bohr Model
• 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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