Important discoveries about the atom

Matter - In chemistry, matter is any physical substance that has mass and takes up space by having volume.

Atom - An atom is the smallest unit into which matter can be divided.

• Atoms are also the smallest unit of matter that still have the characteristic properties of their respective element.

• Atoms are often referred to as the basic building block of life.

Molecule - A molecule is a combination of two or more atoms.

Compound - A compound is created when two or more elements are held together by chemical bonds in a constant whole number ratio.

• All compounds are molecules but a substance that contains only one type of atom is not a compound, it’s just a molecule.

Law of Conservation of Matter - In chemical reactions, matter is neither created nor destroyed.

Law of Constant Composition - Each pure chemical compound always has the same percentage composition of each element by mass.

Dalton’s atomic theory - John Dalton created a scientific theory on the nature of matter and atoms, consisting of three major components.

1. All matter is composed of tiny, indivisible particles, called atoms, that cannot be destroyed nor created.

2. Each element has atoms that are identical to each other in all of their properties, and these properties are different from the properties of all other atoms.

3. Chemical reactions are simple arrangements of atoms from one combination to another in small whole-number ratios.

Scientific theory -  A theory is a new prediction that may be tested by experiments to support or disprove the theory.

• Scientists can never prove a theory to be true.

• Experiments may be used to support a theory, but not to prove it

Law of Multiple Proportions - When two elements can be combined to make two different compounds, and if samples of these two compounds are taken such that the masses of one of the elements in the two compounds are the same in both samples, then the ratio of the masses of the other element in these compounds will be a ratio of small whole numbers.

Electric current can cause chemical reactions to occur, which demonstrates the electric nature of elements.

Cathode ray tube - A cathode-ray tube (CRT) is a specialized vacuum tube in which images are produced when an electron beam strikes a phosphorescent surface.

• Sir William Crookes, the scientist who developed cathode ray tubes in the 1870s, originally thought cathode rays to be negatively charged molecules instead of electrons

Electrons - An electron is a basic unit of positive charge in the atom.

• Electrons were discovered in 1879 by J. J. Thompson, when he determined that cathode rays were fundamental parts of matter he called electrons.

• He determined that their charge to mass ratio was  -1.76 x 10^8 coulombs per gram.

Oil drop experiment - This was an experiment performed by Robert Millikan, which allowed him to calculate the charge of an electron to be -1.60 x 10^-19 columbs.

• Using his own findings and J.J. Thompson’s charge to mass ratio, he calculated the mass of the electron to be 9.11 x 10^28 gram.

• This information led to the creation of the “plum pudding model“ of the atom.

Plum Pudding Model - The plum pudding model of the atom had electrons bathed in a sea of positive charges, similarly to plums embedded in a pudding.

Alpha particles - Alpha particles are complicated, but the only thing you need to remember about them right now is that they are positively charged particles.

Beta particles -  Beta particles are complicated, but the only thing you need to remember about them right now is that they are negatively charged particles.

Gold foil experiment - This was an experiment in which heavy alpha particles were aimed at a thin gold foil and while most particles went through the foil with no visible effect, some were deflected from their path and some bounced back in the direction they came from.

• This experiment was conducted by Ernest Rutherford, Hans Geiger, and Ernest Marsden.

• This experiment led to the creation of the nuclear model of the atom.

Nuclear Model - The nuclear model was a model of the atom that has an extremely small, dense, positively charged nucleus surrounded by empty space that is sparsely occupied by electrons.

• This model is most commonly used in general chemistry classrooms because while it is not fully accurate, it is simple and accurate enough for students to learn in a basic chemistry classroom.

Proton - A proton is a basic unit of positive charge in the atom.

• The proton has a positive charge which is exactly equal in magnitude to the electron charge.

• The proton has a mass of 1.67 x 10^-24 gram.

Neutron - The neutron is the third major particle that marks up the atom, but has no charge.

Each element, when heated or sparked with electricity, gives off characteristic colors.

• The colors consist of discrete wavelengths of light (line spectra) and not a uniform rainbow found when white light is separated by a prism.

• The line spectra of most elements and compounds are very complex.

Spectroscope - A spectroscope is a machine that is used to produce and record the light/color spectra of a particle for examination.

In 1900, Max Planck described light as packets, or quanta, of energy called photons.

Niels Bohr created a theory that electrons move around the nuclei in circular orbits and that electrons only exist in certain “allowed orbits.”

• This aided, along with Max Planck’s theory, aided Bohr  in developing the solar system model of the atom.

Solar System model - The solar system model of the atom describes atoms as consisting of a nucleus with a number of electrons in orbits around that nucleus, similar to a solar system.

• The solar system model is often referred to as the Bohr model, after the scientist that introduced it.

In IN 1924, Louis de Broglie suggested that if light can be considered particles, then small particles like electrons may have the characteristics of waves.

• This inspired Erwin Schrodinger to apply the equations for waves to the electrons in an atom and create the wave mechanical theory of the atom.

Wave Mechanical model - The wave-mechanical model suggests that electrons do not follow a precise orbit around the nuclear.

• The position of the electron in the wave-mechanical model is described by a probability of where it will be located.

Uncertainty principle - The uncertainty principle states that the position and momentum of any particle cannot both be known at exactly the same time.

• As one is known more precisely, the other becomes less certain.

• This principle was discovered  by Werner Heisenberg in the 1920s.

Atomic Models:

1. Solid Particle model (400 BC)

2. Plum Pudding model (1909)

3. Nuclear model Rutherford (1910)

4. Solar System model Bohr (1913)

5. Wave-Mechanical model Schrodinger (1927)

Light and the atom

Ground state - An atom usually exists in the lowest possible energy state, which is called the ground state.

Excited state - When an atom has more energy than in the ground state, it reaches an excited state.

• This means that when objects are heated/charged/energized in any way, the atoms gain energy and the atoms reach excited state.

• When in excited state, the electrons in the atom move up a shell.

• When an atom loses energy in going from an excited state to a ground state, light may be emitted. This explains the “characteristic colors“ given off by elements when they are heated or sparked with electricity, as mentioned above.

Visible light is only a small part of the electromagnetic spectrum.

Electromagnetic spectrum -

All electromagnetic radiation may be considered as waves that are defined by their wavelengths and frequencies.

Wavelength - The wavelength is the distance between two repeating points on a sine wave.

• Wavelength has units of meters and any necessary appropriate prefix.

Frequency - The frequency is defined as the number of waves that pass a point in space in one second.

• Frequency has units of reciprocal seconds, also known as hertz.

The wavelength and frequency of light are inversely proportional to each other.

• Wavelength x Frequency = Speed of Light

The speed of light is 3.0 x 10^8 meters per second.

The energy of electromagnetic waves is proportional to the frequency and inversely proportional to the wavelength.

• Planck’s constant - The proportionality constant, h, has a value of 6.63 x 10 ^-34.

• E = hv = h(c/wavelength)

Atomic Structure

86-90

The most stable position for the electron in the atom is the first level since the electron has the lowest possible amount of energy.

Principal Energy Levels (Shells) - The positively charged nucleus is surrounded by one or more principal energy levels or electron clouds.

• Principal energy levels may also be called the principal shells.

• The principal energy level, or principal shell, nearest the nucleus is assigned the number 1, and each succeeding energy level is numbered with consecutive integers.

• The largest element known needs only seven principal energy levels to hold all of its electrons.

• The number of the principal energy level is given the symbol n.

• Since the principal energy levels become larger the further they are from the nucleus, they can hold correspondingly more electrons.

Sublevels (Subshells) - Each principal energy level within an atom contains one or more sublevels or subshells.

• The number of sublevels possible in each principal energy level is equal to the value of n for that energy level.

• For the 118 known elements, only four sublevels are actually used.

• Sublevels are corresponding letters of s, p, d, and f.

• A sublevel will not exist unless the atom has enough electrons to occupy at least part of the sublevel.

• The table below shows the sublevels possible for each energy level.

• To distinguish one sublevel from another, chemists usually combine the principal energy level number with the sublevel letter in order to indicate in which principal energy level the sublevel is located.

  • For example, “4p” indicates a p sublevel in the fourth principal energy level.

Orbitals - An orbital is a region o space that has a high electron density.

• Each sublevel of the atom contains one or more electron orbital.

• Each orbital contains a maximum of two electrons.

• Orbitals are designated as s, p, d, or f, according to the sublevel they are in.

• Each sublevel has a different number of orbitals, as seen in the table below.

• Orbitals also have different shapes, based on the sublevel

  • An s orbital has a spherical shape, a p orbital has a dumbbell shape, a d orbital is shaped more like a 4-leaf clover, and an f orbital is more tetrahedral.

Electrons will always fill the orbitals with less energy first.

Their sequence, in order of lowest to highest energy, can be seen below.

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d

This sequence is not memorized, it can easily be derived for any given element using the periodic table.

Electron configuration - Each element has its own unique sequence of orbitals that can, once again, be derived using the periodic table.

• The sequence is listed like one would expect, 1s 2s 2p… except each subshell also has a superscript that represents the amount of electrons present in that specific subshell.

  • For example, Na has an electron configuration of 1s2 2s2 2p6 3s1.

• Sequences can be written out fully but as once goes down the periodic table, the sequence can get longer, so instead an abbreviated electron configuration can be written, where the sequence is instead written as the last noble gas before the element being recorded, in brackets, and then all the sublevels that make up the last principal energy level.

  • For example, the full electron configuration for Fe is 1s2 2s2 2p6 3s2 3p6 4s2 3d but it can be abbreviated as [Ar] 3d64s2

• Sometimes, there are exceptions to the energy level order. In order to maintain charge stability within an atom, certain atoms will “unfill“ their previous s sublevel to fill their new d sublevel. AP students don't have to memorize which atoms have these expectations, but are expected to be able to explain why this occurs.

Hund’s Rule - All p, d, or f orbitals in a sublevel must be filled with one electron before a second electron is allowed to pair in any orbital.

• In other words, electrons will only begin to pair up if every orbital in the sublevel is first occupied with one electron.

Orbital Diagrams - Diagrams used to represent orbitals in the valence shell of the electron, where arrows are used to represent the electrons .

• Orbital diagrams are used mainly to describe the valence electron since all of the inner electrons will be paired.

• Orbital diagrams can also be drawn in staggering heights to show the energy differences between orbitals.