Dalton's Atomic Theory: Limitations and Modern Updates

Limitations of Dalton's Atomic Theory

• Dalton's atomic theory was a key early step in understanding matter's inner structure.
• It provided a foundation for studying matter in the 19th century and held sway for roughly 100 years.
• Later research by scientists like J.J. Thomson, Rutherford, and Bohr significantly advanced the knowledge of atomic structure.
• Dalton's theory could explain chemical combination by mass, but not the law of gaseous volumes.
• It couldn't explain differing masses, sizes, valencies among atoms of different elements.
• The theory did not clarify how same or different elements combine to form molecules.
• It didn't address the nature of binding forces between atoms/molecules in solids, liquids, or gases.
• It lacked distinction between ultimate particles of elements and compounds.

Modified Dalton's Atomic Theory (Modern Atomic Theory)

• Dalton's atomic theory has been updated due to research in chemistry and physics,.
• The modern theory retains key points that explain laws of chemical combination.
• Atoms are no longer indivisible but have a complex structure with electrons, protons, and neutrons.
• Atoms of the same element can have different atomic masses and are called isotopes.
• Atoms of different elements can have the same atomic masses and are called isobars.
• The ratio in which different atoms combine in a molecule is fixed and integral, but not always simple.
• Example: Cane sugar (C12H22O11) has a 12:22:11 ratio.
• Atoms are the smallest particles participating in a chemical reaction, regardless of their subatomic components.
• Atoms are no longer indestructible. Nuclear reactions can transform one element's atoms into another and are called transmutation.
• Example: Nitrogen can become oxygen through alpha-ray bombardment.
• Example: Uranium-235 can convert to plutonium-239 using neutrons.
• During nuclear reactions, mass converts into energy, following Einstein's equation (E=mc²).
• Atoms can be interconverted, and mass can change into energy.
• Atoms remain unaffected in ordinary lab reactions and cannot be created or destroyed.

Avogadro's Hypothesis/Law/Principle

• Gay-Lussac's Law states gases combine in simple volume ratios.
• Dalton's Atomic Theory states elements combine in simple whole number ratios.
• Berzelius attempted to connect Dalton's Atomic Theory with Gay-Lussac's Law.
• Berzelius suggested a relationship between gas volume and the number of atoms.
• Berzelius Hypothesis: Equal gas volumes under the same conditions contain equal atom numbers.
• Experimentally, one volume of hydrogen and chlorine yields two volumes of hydrogen chloride gas.
• Applying Berzelius' Hypothesis, it states one compound atom is made of 1/2 an atom of hydrogen and 1/2 an atom of chlorine, contradicting Dalton's idea that atoms are indivisible.
• Avogadro distinguished between atoms and molecules.
• An atom, per Avogadro, is the smallest particle of an element in a chemical reaction and which may or may not exist independently.
• A molecule, per Avogadro, is the smallest particle of a substance that is capable of independent existence.
• Avogadro's hypothesis: Equal gas volumes under the same conditions contain equal molecule numbers.
• Avogadro's Law explains reactions between gases and is the following:
• Experimentally, one volume of hydrogen and chlorine make two volumes of hydrogen chloride gas.
• Applying Avogadro's Law, one molecule of hydrogen chloride gas contains 1/2 a molecule of hydrogen and chlorine.
• This does not contradict Dalton's atomic theory as molecules contain one or more atoms.

Applications of Avogadro's Law

• Atomicity Calculation: Atomicity is the number of atoms per molecule of an element.
• Oxygen (O2) atomicity is two, and ozone (O3) atomicity is three.
• Molecular Mass/Vapor Density Relationship:
• Vapor density = (Mass of gas volume at STP) / (Mass of same volume of H2 at STP)
• Vapor density = (Mass of one gas molecule) / (Mass of two H atoms)
• Molecular mass = 2 × Vapor density
• Vapor density is also called 'relative density'.
• Mass/Volume Relationship:
• Molecular mass = 2 × Vapor Density = 2 × (Mass of 1L gas at STP) / (Mass of 1L H2 at STP)
• Molecular mass = (2 / 0.089) × Mass of 1L gas at STP = 22.4 × Mass of 1L gas at STP
• 22.4 liters of any gas at STP weighs the same as its molecular mass in grams.
• This is called the Gram-Molecular Volume (G.M.V) Law.

Relative Atomic Masses

• Determining atomic masses was a key interest after discovering matter is made of atoms.
• Atoms are too small for direct weighing.
• Finding the mass of a single atom through weighing a large sample isn't possible due to the inability to count the atoms contained.
• Avogadro's hypothesis states that equal volumes of different gases under the same temperature and pressure contain an equal number of molecules.
• For equal volumes of two gasses under same conditions, the ratio of their masses equates to mass ratio of their molecules.
• For hydrogen and oxygen, masses are always in a 1:16 ratio.
• An oxygen molecule is 16 times heavier than a hydrogen molecule.
• Because a hydrogen molecule contains two atoms and an oxygen molecule also contains two atoms, an oxygen atom is 16 times heavier than a hydrogen atom.
• It follows that if the atomic mass of hydrogen is one, the relative atomic mass of oxygen is 16.
• In the beginning, the atomic masses of elements were relative to hydrogen (set to 1).
• Because most elements had fractional atomic masses in comparison to hydrogen, oxygen (set to 16) became the standard.
• Carbon (set to 12) then became the accepted standard.
• Relative mass of hydrogen is 1.008 and that of oxygen is 15.9994.
• The atomic mass of an element is how many times heavier its atom is compared to an atom of carbon (set to 12).
• Atomic masses are relative and expressed in atomic mass units (amu).
• 1 amu is 1/12th of the mass of a carbon-12 atom.
• The atomic mass of hydrogen is 1.008 amu, and oxygen is 15.9994 amu (or 16 amu).
• Before 1961 the reference for atomic masses was oxygen.
• Chemists (using the chemical scale) set oxygen to 16, while physicists (using the physical scale) set only the isotope 16O to 16 and to solve the disparity, scientists settled on using the isotope 12C.
• The unified scale now expresses atomic masses using the 12 6C scale.
• The symbol 'u' is now used instead of amu in the unified scale but amu is still used.
• Atomic mass of hydrogen can be written as 1.008 u and oxygen as 16 u.
• A 'mass spectrometer' has helped determined the accurate atomic masses of elements.
• If an element has isotopes, the atomic mass used is the average value.
• For ordinary chlorine, the average atomic mass is 35.5 u.
• The average must be included in the definition of atomic mass.
• The atomic mass of an element is the average relative mass of its atoms as compared with an atom of carbon-12 taken as 12.
• A = Σ (pi Ai)/100 : where p is the percent abundance of the isotope its the atomic mass, A.
• Average relative atomic mass can be calculated from the 'fractional abundances' of isotopes.
• Fractional abundance is the fraction of the total atoms comprised within that isotope.
• Average atomic mass is the sum of the fractional abundance times the isotopes.
• For neon: Average atomic mass equals 20.179 u.
• Carbon contains three isotopes (12, 13, and 14 u), carbon-12 is the reference for the atomic mass scale.