UNIT 2

Structure and History of the Atom

History of the Atomic Theory:

DEMOCRITUS, (~400BC) ancient Greek philosopher, first to suggest that matter existed in a very small fundamental particle. He hypothesized that matter could be subdivided again and again until it would finally reach a particle which could not be divided. This smallest “invisible” particle he called “atomos.”

History/Atomic Theory continued:

ARISTOTLE hypothesized that matter was “continuous” in nature. That is, matter could be subdivided again and gain an infinite number of times and you would never reach a basic particle which could not be divided. This thought was accepted for the following 200 years. Both Democritus & Aristotle’s views were NOT based on experimental fact.

History/Atomic Theory continued: JOHN DALTON (1808) was an

Englishman who was the first to develop and publish a theory about how atoms looked and behaved. He thought the atom was a solid sphere, much like a billiard ball.

Dalton’s Atomic Theory:○ All matter is composed of extremely small

particles called atoms.○ Atoms of a given element are identical in size,

mass and properties. (This was later disproved by the discovery of isotopes).

○ Atoms cannot be subdivided, created or destroyed. (This was later disproved by the discovery of protons, neutrons and electrons).

○ Atoms of different elements combine in simple whole-number ratios to form compounds.

○ Atoms are not made or destroyed in a chemical reaction; they are joined, separated or rearranged to form a new substance

J. J. THOMPSON (1897) in his CATHODE RAY TUBE EXPERIMENT, discovered electrons.

CATHODE RAY TUBE EXPERIMENT

Electron Stream w/o applied electric field Electron Stream after applying field, + on top, - on bottom

In this experiment, particles were deflected away from negatively charged plates & toward positively charged plates. Since positives charges attract negative charges, he concluded that atoms contain negatively charged particles called electrons.

Thompson proposed that the atom has a mixture of positive charges and negative charges. This model was similar to a chocolate chip cookie. The “chips” represent electrons lying in surrounding “dough” made of protons.

History/Atomic Theory continued: ERNEST RUTHERFORD (1911)

worked for Thomson in his lab for a while. He performed the GOLD FOIL EXPERIMENT which proved that the atom as mostly space and that all of the positive charge was located in a very small central nucleus.

GOLD FOIL EXPERIMENT

In this experiment, alpha particles were sent through a thin sheet of gold foil. Most of the particles went straight through the foil. Some particles were slightly deflected & a few were deflected back toward the source. The reason for the deflection was the positive charge of the nucleus.

Note: If an atom was the size of Pirate Stadium, the nucleus would be the size of your class ring sitting on the 50 yard line.

History/Atomic Theory continued: NIELS BOHR (1913)

Danish scientist was the first to suggest the “planetary model” of the atom. Electrons move around the nucleus in a set path called orbits much as the planets orbit the sun. Each ring or level represents a certain amount of energy. Electrons farther from nucleus have greater energy.

History/Atomic Theory continued: JAMES CHADWICK (1936) British

chemist who discovered the neutron. Scientists had looked for another subatomic particle (other than protons and electrons) for many years but couldn’t find them. Chadwick found out that the neutron is approximately the same mass as the proton but had NO electrical charge.

ERWIN SCHRODINGER – an Austrian physicist, along with Werner Heisenberg and Louis de Broglie, postulated the quantum (wave) mechanical model of the atom which we believe is true today.

History/Atomic Theory continued:

Figure 2-6 In this model, the electrons

do not actually “orbit” the nucleus but are found only in definite areas, called orbitals, based on the amount of energy they have. Like the particles in a light wave, electrons move so rapidly that their exact location is unknown – this creates a general “cloud” of electrons around the nucleus.

THE ATOM PROTON- positively-charged particle which is

found in the nucleus; it is considerably more massive than the electron. Proton mass is about 1.6726 x 10-24 g.

ELECTRON- very light negatively-charged particle which is found somewhere outside the nucleus. It weighs only 1/1837 that of a proton, but its negative charge is as powerful as the positive charge of a proton.

NEUTRON- a particle found in the nucleus which is approximately the same mass as a proton but does not have an electrical charge associated with it – it is neutral.

THE ATOM ATOMS ARE ALWAYS NEUTRAL PARTICLES—that

is, they contain the same number of protons as electrons (and it doesn’t matter how many neutrons they have)

IONS – particles which do not have the same number of protons and electrons so therefore they do have an electrical charge associated with them. If they have more protons than electrons, they will have a positive charge. If they have more electrons than protons, they will have a negative charge. # of Protons – # of Electrons ( e-) =

charge of ion Example 2-1:

If an element has 12 protons & 10 electrons, what is the charge of this ion?

If an element has 17 protons & 18 electrons, what is the charge of this ion?

The Atom

Atomic Number – the number of protons found in the nucleus of an atom. On the Periodic Table it is the large whole number in the upper corner. Notice that the only thing which makes an element different from another one is the number of protons it contains. The number of protons identifies an element.

THE ATOM Example 2-2: An atom with 6 protons is

___________________; an atom with 5 protons is __________________ and an atom with 7 protons is ___________________.

THE ATOM MASS NUMBER – the

number of protons and neutrons found in the nucleus of an atom, this DOES NOT show on the Periodic Chart. Used most often to identify an isotope of an element.

# of neutrons + # of protons = mass #

THE ATOM Example 2-3:

If Calcium’s atomic number is 20 and mass number is 46, how many protons and neutrons are present?

If Neon’s atomic number is 10 and mass number is 22, how many protons and neutrons are present?

THE ATOM ISOTOPE – atoms with the

same number of protons but different numbers of neutrons. This difference in neutrons also results in different masses. For instance, “regular” Carbon has 6 protons and 6 neutrons. It is called Carbon-12. But Carbon-14, the radioactive carbon used in carbon dating, has 6 protons and 8 neutrons.

Example 2-4: ○ What isotope has 19

protons and 20 neutrons?

○ What isotope has 26 protons and 31 neutrons?

THE ATOM ATOMIC MASS – the mass of a single atom.

Atoms are too small to measure mass in grams, so a system of measurement was produced. The mass of a single atom is found by comparing it to the standard Carbon-12 atom. The unit of measurement for a single atom is Atomic Mass Units (amu).

Note: Atoms are very small – typically 1 x 10-8 cm in diameter. 1.0g of lead contains 2.9 x 1021 atoms of lead. By comparison, the earth’s entire population is only 5 x 109 people

THE ATOM AVERAGE ATOMIC MASS – The atomic

mass that shows on the Periodic Table for each element is really an average of the masses of all the isotopes of that element, weighted by their percentage of abundance. Every element on the Periodic Table has at least 3 isotopes; some of them have 20 or 30 or more isotopes. Atomic mass is closest to the mass of the most abundant isotope.

THE ATOM Equation to find the average atomic mass of an

element: Avg. Atomic Mass = (% isotope #1) (mass of #1) +

(% isotope #2) (mass of #2) + (% of isotope 3) (mass of #3) +……

Example 2-5: Naturally occurring chlorine is 75.53% chlorine-35 and 24.47% chlorine-37. What is the average atomic mass which should be placed on the Periodic Table for the element?

Example 2-6: The element neon consists of three isotopes with masses of 19.99, 20.99 and 21.99 amus. These three isotopes are present in nature to the extent of 90.92%, 0.25% and 8.83% respectively. Calculate the atomic mass of neon.

CHEMICAL CONFIGURATION

is a way of writing atoms or ions which gives you lots of information.

Mass # is on top, element symbol in middle, and atomic # on the bottom:

CHEMICAL CONFIGURATION Example 2-7: Write the chemical configuration for a

particle containing 5 protons, 6 neutrons and 5 electrons.

Write the chemical configuration for a particle which contains 8 neutrons, has a mass number of 14 and contains 6 electrons.

Write the chemical configuration for a particle which contains 18 electrons, has a mass number of 32 and has a -2 charge.

CHEMICAL CONFIGURATION

Sometimes you are not given enough information to write the entire chemical configuration for a particle but you can at least identify as to what element it is and whether it is an atom or a positively-charged or negatively-charged ion.

Example 2-8: What particle has +2 charge and has 12 protons? What particle has 44 protons and 44 electrons? What particle has a -3 charge and contains 36 electrons?

ELECTRON LOCATION

Electrons are distributed around the nucleus in specific ways and each electron has an "address" that consists of 4 parts – energy level, sublevel, orbital, & spin.

Energy Level - indicates how far from the nucleus the electron is located. Energy levels closest to nucleus have least energy. The number of electrons in any energy level is 2n2 (n=energy level)

ELECTRON LOCATION

Example 2-12: How many electrons

could the 2nd energy level hold?

How many electrons could the 3rd energy level hold?

ELECTRON LOCATION Sublevel - Each

energy level is divided into sublevels which are designated by the letters s, p, d, f.

Orbital – regions of different shape & energy that have a high probability for finding electrons

ELECTRON LOCATION Each sublevel has a set number of orbitals related

to positions in 3-D space. s sublevel - 1 orbital; spherical shaped p sublevel - 3 orbitals; dumb-bell shaped d sublevel - 5 orbitals; f sublevel - 7 orbitals;

ELECTRON LOCATION Each orbital holds only 2 electrons Spin – Like Earth, an electron in an orbital

can be thought of as spinning on an internal axis. The electrons within an orbital must

have opposite spins.

Sublevel # of orbitals Total e- s 1 2 p 3 6 d 5 10 f 7 14

Rules for Arranging Electrons

Aufbau Principle - electrons are placed in orbitals of lowest energy first. The orbital with the lowest energy is the 1s orbital. The 2s orbital is the next highest in energy, then the 2p orbitals.

Pauli Exclusion Principle - an orbital may hold only two electrons & they must have opposite spins. Arrows are used to indicate the direction of the electron spin.

Rules for Arranging Electrons

Hund’s Rule - when filling orbitals of equal energy, one electron enters each orbital until all contain an equal amount. Then orbitals receive a second electron.

ELECTRON ARRANGEMENT

Knowing the rules for electron arrangement allows us to write a description of where in an atom the electrons are actually located.

ELECTRON ARRANGEMENT Example 2-13: Write the electron

configuration for each of the following: Hydrogen helium lithium beryllium boron oxygen

ELECTRON ARRANGEMENT

ELECTRON ARRANGEMENT Example 2-14: A certain element's electron

configuration is known to be 1s22s22p63s23p5. What element are we describing?

ELECTRON ARRANGEMENT

An Orbital Diagram provides the same information with more detail. In this representation, the number of electrons in each orbital & each sublevel is indicated. The spin of each electron is also shown.

Example 2-15: Draw the orbital notation for fluorine (F - 9 electrons)

ELECTRON ARRANGEMENT

Example 2-16: Draw the orbital notation for calcium (Ca - 20 electrons)

ELECTRON ARRANGEMENT Order of Electron Fill: Just about the time you think you are

"getting the hang of this" there is an exception. After the 3p fills up you would think the 3d would fill, but it is not so. Electrons will fill the areas which require the least amount of energy, and it takes less energy to fill the 4s than it does to fill the 3d. To help you know the filling order of electrons, use the diagonal rule or "tree".

ELECTRON ARRANGEMENT Example 2-17: Write the electron configuration for iron

(Fe) based on the diagonal rule.

Draw the orbital diagram.

ELECTRON ARRANGEMENT Example 2-18: According to the Diagonal Rule,

which is higher-energy electron:

3d or 4sHow do you know?

5d or 6pHow do you know?

6p or 7sHow do you know?

ELECTRON ARRANGEMENT Example 2-19: In the element chlorine,

which electron is the highest-energy electron?

Which electron is physically farthest from the nucleus?

How many unpaired electrons are in chlorine?

PERIODIC TABLE and ELECTRON CONFIGURATION An element’s placement on the Periodic Table is

related to the arrangement of electrons in an atom. A. Each horizontal row (called a period) corresponds to an energy level. The periodic table can be broken into sections that represent the orbitals. And if you read the periodic table like a paragraph from left to right, the order of electron fill is provided.

Note: The "s" and "p" electrons will have an energy level number same as the period they are on. The "d" electrons will have an energy level ONE LESS than the period they are on & "f" electrons will have an energy level two less than the period they are on

PERIODIC TABLE and ELECTRON CONFIGURATION

PERIODIC TABLE and ELECTRON CONFIGURATION Example 2-20: Use the Periodic Chart to

read the electron configuration for Na (#11)

Example 2-21: Use the Periodic Chart to write the electron configuration for …

Al (#13) Ar (#18) Ni (#28) Zn (#30) Br (#35)

PERIODIC TABLE and ELECTRON CONFIGURATION Electron configurations are sometimes written using noble-

gas notation. In this form of notation, instead of writing out the entire series of energy levels and sublevels, you can use an abbreviation. First identify the noble-gas that comes before the element. Write the symbol of the noble-gas in brackets & continue writing the energy level, sublevel, and number of electrons until you reach your element.

For example: Electron Configuration of ○ Si 1s22s22p6 3s23p2

Noble-gas Notation of

○ Si [Ne] 3s23p2

PERIODIC TABLE and ELECTRON CONFIGURATION Example 2-22: Write the electron

configuration using noble-gas notation for….

O (#8) P (#15) Cu (#29) Br (#35)