Modeling Atoms 1.1 

Modeling Atoms 1.1 


An atom is the smallest particle of an element that retains its identity in a chemical reaction


Early Atomic Theory

  • Existence of atoms as indivisible units of matter has been suggested since 2,400 years ago

  • Early ideas lacked experimental evidence and did not explain chemical behavior

  • Many scientists have used scientific methods over the past few centuries to contribute to what we know of atoms today

  • Atomic theory supported by experimental evidence begun with John Dalton in 1803

  • Studied the ratios in which elements combine in chemical reactions

  • Incorporated the observations of other scientists when forming his postulates

  • A model of an atom using Dalton’s postulates/ atomic theory would show a tiny, indivisible sphere


Atoms

  • Atoms of the same element are identical

  • Atoms of any one element is different from any other element

  • Atoms of different elements physically mix together

  • Atoms of different elements can chemically combine together in simple whole-number ratios to form compounds

  • Chemical reactions occur when atoms are separated, joined, or rearranged.


Discovering Subatomic Particles

  • Construction of atomic models using Dalton’s postulates marked the beginning of atomic theories

  • Much of Dalton’s theory is accepted today, but we now know that atoms are divisible

  • Atoms can be broken down into smaller, fundamental particles (subatomic particles), including electrons, protons, and neutrons


Thomson’s Discovery of Electrons

  • In 1897, Thomson discovered electrons (negatively charged subatomic particles)

  • He passed electrical currents through electrodes in sealed gas tubes

  • Electrons become charged and resulted in a glowing beam or cathode ray that traveled from one end to the other 

  • Cathode rays are attracted to positively charged metal plates and are repelled by a negatively charged metal plate

  • Thomson deduced that cathode rays (electrons) are beams of negatively charged particles due to the fact that opposite charges attract and like charges repel. 


Thomson’s Experiment

  • Observed that cathode rays are deflected by a magnet

  • As a result of this observation, he concluded that the mass of the negative charge was very small compared to the magnitude of the charge.

  • A cathode ray travels from the cathode to the anode

  •  A magnet held near the cathode ray causes ray to change direction

  • One electrode, called the cathode, is negatively charged

  • One electrode, called the anode, is positively charged


Thomson’s Atomic Model

  • Based on his discovery of electron properties, he constructed an atomic model that was the next major development in atomic theory

  • Thought that electrons were distributed throughout a positively charged material

  • This model is known as the “Plum-Pudding Model”

  • Electrons are stuck into a lump of positive charge in this model


Protons and Neutrons

  • Atoms have no net electric charge

  • Can, however, lose negatively charged electrons

  • The part of an atom that remains when an electron is lost is positively charged

  • Evidence for positively charged particles was founded in 1886 by Eugen Goldstein

  • Observed a cathode ray tube and discovered rays traveling in the opposite direction of cathode rays

  • Concluded those rays were composed of positively charged particles

  • Goldstein’s and others’ experiments led to the discovery of protons (positively charged subatomic particles)

  • James Chadwick confirmed the existence of neutrons (subatomic particles with no charge) in 1932


Discovering the Nucleus

  • Scientists wondered how the subatomic particles were put together in an atom as a result of the discovery of subatomic particles

  • Many scientists favored Thomson’s model of the atom

  • Ernest Rutherford tested Thomson’s atomic model in 1911

  • Test involved beaming alpha particles, helium atoms that have lost their electrons and positively charged, at a sheet of gold foil


Rutherford’s Experimental Design


  • A beam of alpha particles strikes a sheet of gold foil surrounded by a fluorescent screen

  • Most of the particles pass through the foil with no deflection or very little deflection


  • Within the gold foil, alpha particles that approach the nucleus closely are deflected

  • This deflection was observed for a small fraction of the particles

  • Concluded that most alpha particles pass through the foil because the atom is mostly empty space

  • Some particles are deflected by a small positively charged region with most of the atomic mass that he called the nucleus

  • Based on this experiment, Rutherford constructed a revised atomic model known as the nuclear atom

  • In this model, the atom is mostly empty space

  • All of the positive charge and almost all the mass are concentrated in a small region he called the nucleus (positively charged central core of an atom and is composed of protons and neutrons)

  • The electrons are distributed around the nucleus and occupy most of the volume


  • After discovering the nucleus, Rutherford used existing ideas about the atom and proposed an atomic model in which the electrons move around the nucleus (like planets moving around the sun)


Visualizing the Atom

  • Matter is made atoms that are too tiny to be seen without a microscope

  • Each atom is made up of smaller subatomic particles

  • Positively charged protons (p+) and uncharged neutrons (n0) are in the nucleus (the dense central core of an atom)

  • The mass of an atom a proton and neutron is about 1 atomic mass unit (amu), which is approximately 1830 times larger than the mass of an electron (about 0.0005 amu)

  • Negatively charged electrons (e-) are in the electron cloud that surrounds the nucleus which accounts for most of an atom’s mass

  • The electron cloud is a model that is used to describe the region negative charge surrounding the nucleus and is primarily empty space

  • Atoms do not really have a electron cloud, but it is useful when discussing atoms and their properties




Atomic Number

  • An element is the simplest form of matter that has a unique set of properties

  • The number of protons is what makes one element different from another

  • The number of protons in the nucleus of an atom is called an element’s atomic number

  • Atoms can be described as electrically neutral particles because they have no net charge

  • For an atom to be neutral, the number of protons must equal the number of electrons


The Periodic Table

  • There are 118 elements in the periodic table, all with different atomic numbers

  • A periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties

  • The elements are listed in order from left to right and top to bottom by atomic number

  • Organization of the periodic table allows you to easily compare the properties of one element (or group of elements)

  • Elements above and below each other tend to have similar properties


Chemical Symbols

  • An element’s name is often not convenient to use

  • Each element can be represented by a one/ two chemical symbol

  • These chemical symbols are used to represent the elements in the periodic table and chemical formulas for compounds

  • When writing the formula for a compound, you combine the symbols of the elements that make up the compound


Mass Number

  • When describing the structure of an atom, most of its mass is from protons and neutrons

  • Electrons have almost no mass

  • The total number of protons and neutrons is called the mass number


Representing Atoms

  • The composition of any atom can be represented in shorthand notation using the element’s chemical symbol, atomic number, and mass number

  • Atoms can also be represented by using the name and mass number of an element


Mass Number and Neutrons

  • The number of neutrons in an atom can be calculated using the atomic number and mass number of an atom

  • Number of neutrons = mass number - atomic number



Isotopes

  • All atoms of an element have the same number of protons

  • Atoms of the same element may have different numbers of neutrons

  • Isotopes are atoms that have the same number of protons but different number of newtons

  • Since isotopes have different numbers of neutrons, they also have different mass numbers 


Comparing the Atomic Masses

  • It is not convenient to measure the mass of a single atom because the mass is so small

  • It’s more convenient to compare the relative masses of atoms using an reference isotope

  • A reference isotope is an isotope whose mass can be measured accurately and is used as the basis for a scale to compare the masses of other atoms


Isotope Abundance and Atomic Mass

  • In nature, most elements occur as a mixture of two or more isotopes

  • Each isotope of an element has a fixed mass and a natural percent abundance

  • An isotope’s natural percent abundance is the percentage of the isotope found in a naturally occurring sample of an element

  • The average atomic mass (atomic mass) of an element is weighted average of the masses of its isotopes

  • A weighted average mass reflects the masses and the relative abundance of isotopes as they occur in nature


Calculating Atomic Mass

  • Isotopes that are more common in nature have greater importance in atomic mass calculations

  • The average atomic mass of an element can be calculated using the isotopic composition

  • The resulting sum after multiplying the mass of each isotope by its natural abundance (expressed as decimal) and then adding the products is the weighted average mass of the atoms of an element as they occur in nature

  • The atomic mass is used most in scientific calculations


Understanding Atomic Theory


Chemistry throughout the ages


Ancient Times

  • Ancient humans

  • Ancient Egyptians

  • Greek philosophers

  • Ancient Humans: fire and cooking

  • Egyptians: Embalming techniques for mummification; extracting metals from ores; perfumes; paints; and make-up

  • Greek Philosophers: Debated about what matter is made of; Atomos (indivisible particles) or the four elements (fire, water, air, earth)


Middle Ages

  • Alchemists

  • Development of many techniques still used today

  • Combined science with mystical rituals

  • Philosopher's Stone; transmuting lead into gold


17th to 18th Centuries

  • Robert Boyle

  • Antoine-Laurent Lavoisier

  • Development of the scientific method led to the reproduction of experiments due to the measurements rigorously taken

  • Discovery of many scientific laws (e.g., Gas Laws, Laws of Conservation of Mass)


1803

  • John Dalton

  • Combined his research with the known laws to propose the first modern Atomic Theory

  • 5 postulates:

  • Elements are small, indivisible particles

  • All atoms of the same element are identical

  • Atoms of different elements are different

  • Atoms of different elements can combine with each other in simple, whole number ratios to form compounds

  • Chemical reactions occur when atoms are separated, joined or rearranged, However, atoms of one element are NOT changed into atoms of another element by a chemical reaction.


1897

  • J.J Thomson

  • Cathode RayTube experiments

  • 3 Experiments

1)  Bent the rays with a magnet to show the negative charge was contained in the rays

2) Obtained a nearly perfect vacuum to show the ray would bend in an electric field, therefore the ray was negatively charged to charge ratio

3) Bent the rays in a magnetic field to determine the mass to charge ratio. This showed that the particles were nearly 2000 times smaller than a hydrogen atom.

  • Conclusion: Plum Pudding Model


1908 - 1911

  • Ernest Rutherford

  • Gold Foil Experiment - was designed to confirm the Plum Pudding Model

  • Nuclear Atom


Properties of Atoms

  • Neutral Atom 

  • Are neutral (no overall charge)

  • The number of protons (p+) = The number of electrons (e-)

  • Ion

  • Are not neutral (has overall charge)

  • The number of protons ≠ The number of electrons


  • Cations

  • Positive charge

  • More protons than electrons (lost electrons) 

  • Anions

  • Negative charge (gained electrons)

  • More electrons than protons


  • Atomic Number (Z) 

  • It often has the symbol Z due to it having to do with charges (old symbol)

  • Number of protons

  • Whole number on the periodic chart

  • Periodic table is in order of atomic number

  • Each element gains a proton, which also means it gains an electron


  • Period = Horizontal row

  • Group = Vertical column

  • Elements in the same group tend to have similar properties


  • Mass Number (A)

  • Mass Number = Protons + Neutrons

  • Neutrons = Mass Number - Protons

  • Not on the periodic table

  • Sometimes called the isotope mass

  • Whole number (not on the periodic chart)

  • Isotopes

  • Same element, different number of neutrons (i.e., different mass numbers)


Representing Atoms

  • Hyphen notation

  • Name hyphen mass number

  • Nuclear symbol

  • First letter capitalized; second letter lowercase


Atomic Mass

  • Mass of One Atom

  • Can’t measure the mass of one atom; it's easier to measure the relative masses

  • Carbon-12 defined as having a mass of 12 amu (atomic mass unit)

  • 1 amu is defined as the mass of one proton and neutron (actual mass is approximately 1.67 x 10^-27 grams)

  • The mass of electrons is negligible compared to protons and neutrons


(Average) Atomic Mass

  • Decimal number on the periodic chart

  • The average mass of all neutral isotopes of a particular element


Calculating Average Atomic Mass

  • Atomic Mass = “the sum of” (fraction abundance)(mass of isotope)

  • Fraction abundance = % abundance / 100

  • Mass of Isotope = Use the Atomic mass, but if it isn't given then use mass number

  • Do NOT take the masses of isotopes and divide by the number of isotopes

  • Always round the atomic mass to the hundredths (two places behind the decimal point)

  • When determining the most common isotope, look at the one that is the closest to the average atomic mass.