Lecture 1- Atomic theory 1 Flashcards
pure sample
- contains particles, or units, of one specific atom of molecule
- FU=ionic
- can separate with chemical or nuclear change
Mixture
- contains particles, or units, of more than one specific atom or molecule
- each have their own properties
- can separate by undergoing physical change
Democritus
- all matter made up of very small, individual particles called atoms
Antoine Lavoisier
- father of chemistry
- law of conservation of matter
- gave oxygen its name and showed it was involved in combustion reactions
Aristotle
4 elements
- sided with him over Democritus
Joseph Proust
- law of definite proportions
Law of definite proportions
- joseph proust
- different pure samples of the same compound always contain the same proportions of each element by mass
- ex. water is always 88.8% oxygen and 11.2% hydrogen by mass
John Dalton
- law of multiple proportions
- atomic molecular theory
- if elements were made of tiny indivisible particles, then molecules of a particular compound would always be composed of equal numbers of each type of element
- thus, each element in a given compound would account for a consistent percentage of that compounds mass
Daltons 4 postulates
- Each element is composed of tiny particles called atoms
- All atoms of a given element are identical and all atoms of different elements are different (wrong because isotopes)
- Atoms are not crested or destroyed in chemical reactions. They are dissociated, combined, or recombined
- Compounds are formed when atoms of different elements combine with one another. A given compound always contains the same relative numbers and types of atoms
JJ Thompson
- used cathode ray tube to discover electrons
- stream of particles made a green glow and were attracted towards the positive side
- calculated charge to mass ratio
- plum pudding
Earnest Rutherford
- Gold foil experiment
- Beta particles toward positive= negative- much deflection
- Alpha towards negative= positive- more mass, didn’t reflect as much
- gamma= neutral- highspeed photon
Gold foil experiment
- Nucleus is a very small positively charged core containing protons and neutrons
- negatively charged electrons are extremely tine and occupy the vast majority of the atoms volume
sub atomic particles
Proton- +1, 1amu
Electron- -1, .0005amu
Neutron- 0, 1amu
Mass number
Equal to the number of protons plus the number of neutrons
atomic number
number of protons
Isotopes
- isotopes of an element have the same number of protons, but different number of neutrons
- exhibit identical chemical behavior
Atomic masses
- mass of one c-12 atom is exactly 12 amu
- all others measured on a scale relative to that
- average of each isotope on periodic table
Average atomic mass
- Atomic mass given to you on the periodic table is an average of all the isotopes of that element
- average mass of any large number of atoms of a given element is always the same for that given element
Mass Spectrometry
- Used to compare the masses of isotopes
- Atoms form a pure sample of an element are ionized and accelerated through a magnetic field
- Isotopes with smaller masses experience a greater degree of deflection
How mass spectrometry works
- isotope in a gas state
- hit with an exact ionizing energy
- measures mass to charge ratio
- accelerated through an electric field
- Deflected through a magnetic field
- recorder measures frequency and mass to charge ratio
mass number vs atomic mass
- Mass number= mass of each particle
- atomic number= average mass of all isotopes of the atom
average atomic mass how to find
(mass X abundance) + (mass X abundance)= average atomic mass
Formula weight
the basic unit in a network sold
- FU=ionic compounds
molecular weight
molecules that can exist independently
- molecular compounds
Formula and Molecular weights
the sum of the atomic masses of the atoms in a chemical formula
how to find formula weight
regular add all the elements
Mass percent
(# atoms of the element)(elements atomic mass) over (formula weight of compound) X 100
Law of multiple proportions
- John Dalton
- When two elements form more than one compound, the different masses of element A that combine with a fixed mass of element B can always be expressed in a small whole number ratio
- The small whole number ration tells us how many times more A is contained in the more massive compound. It must be a whole number, as compounds can only bye build out of whole atoms
- Ex. H2O2 has twice as many oxygen atoms so its equal
Mass ratio of elements in a pure sample of a compound
- the ratio of isotopes of a certain element in any pure sample of a given compound is constant
- if we could calculate the average mass of all the atoms of an element in a pure sample, we would obtain the average atomic mass of that element, which is given on the periodic table
- as the ratio of the different atoms of each element that form a given compound is constant, the ratio of the masses of the constituent elements in a pure sample of that compound will also be constant
does data from mass spectrometry indicate that modern scientists have to make modifications to Daltons model?
Yes because data from mass spectrometry indicates that modern scents have made modifications to Daltons model because Dalton said that each atoms of each element is identical, but mass spectrometry says that the same elements can have atoms of different masses
atomic molecular theory
- Dalton
- the ratio/percent of the same elements in the same compound will always remain constant.
- when two elements form more than one compound, the different masses of A will combine with a fixed mass of elmetn B that can be expressed in a whole number ratio. No matter the size of the compound, the percent of an element in the sample will always be the same because the ratios of elements in molecules/units remains constant. That is why the two samples of CO2 would have the same percent of Carbon
why was it that the vast majority of alpha particles that traveled through the gold foil in Rutherford’s experiment experienced no deflection?
because the alpha part ices were significantly larger than the electrons and didnt encounter anything large enough to significantly deflect their math (with the exception of the few deflected by the positive atoms in the nucleus)
experiment that proved atoms contain positive, negative, and neutral particles
Ernest Rutherfords experiment proved that atoms contain positive, negative, and neutral particles. He placed radioactive material in a lead casing and emitted particles through a magnetic field. He determined that the Beta particle carried a negative chafe since it went towards the positive side, the alpha particle carried a positive charge since it went towards the negative charge, and the gamma carried a neutral chafe since it was attracted to neither and goes straight through.
experiment that proved electrons carry a negative charge
JJ Thompson with cathode ray tube experiment. Used a stream of part ices that emitted a green glow when connected to power and used an electric field. The particles moved towards the positive charge, proving the particles had a negative charge
