Chapter 8

Atoms and Periodic Properties

The Atom

• The concept that matter is made up of atoms dates back to the Greeks 2,500 years ago.

• Democritus in 460-362 BC came up with the atomic model in which the atom was indivisible.

• In the early 1800’s John Dalton reintroduced the concept of the atom.

Dalton’s Atomic Theory

•

Indivisible particles

called atoms make up all matter.

• •

All the atoms of an element are exactly alike in shape and mass.

• The atoms of different elements differ from one another in their masses.

• Atoms chemically combine in definite whole number ratios to form chemical compounds.

• Atoms are neither created or destroyed in chemical reactions.

Bolded items were later found to be untrue

Discovery of the Electron

• Thomson discovered the existence of the electron in 1897.

• He placed charged metal plates on each side of a beam from a high voltage electrical source in a vacuum and found that the beam was deflected away from the negative plate.

• Since like charges repel, this meant that the beam was made up of negatively charged particles.

The Nucleus

• Ernest Rutherford discovered the nucleus in 1907.

• Alpha particles from a radioactive source were allowed to move as a narrow beam of fast moving particles through a small opening in a lead container directed to a very thin gold foil.

• The alpha particles were detected by plates that produced small flashes of light when struck by alpha particles.

• Most of the alpha particles went right through the foil. Some of the particles were deflected at different angles, some even backwards.

Conclusions of Rutherford’s Experiment

• The positive alpha particles that were repelled was because they encountered massive positive charge concentrated in a small region of the gold atom.

• He concluded that the atom contains a tiny, massive and positively charged nucleus surrounded by electrons.

• The volume of the atom is mostly empty space.

The Nucleus

• The nucleus contains protons and neutrons.

• The protons are positively charged while the neutrons do not have a charge, they are neutral.

• The number of protons is equal to the atomic number of the element.

• The number of protons is equal to the number of electrons, since all atoms are neutral.

• All atoms of the same element have the same atomic number, and hence number of protons.

• Every element has a different atomic number.

• There are 113 different elements.

Isotopes

• The neutrons in the nucleus, along with the protons, contribute to the mass of the atom.

• The number of neutrons in an atom of a given element may vary.

• Atoms of the same element which have different number of neutrons are called isotopes.

• Isotopes have the same number of protons and electrons but different number of neutrons.

Atomic Mass Unit

• The unit of mass used for atoms is the atomic mass unit (AMU) • This is defined as 1/12 of the mass of C-12, the isotope of carbon which contains 6 protons and 6 neutrons.

• The mass number of atoms is equal to the number of protons plus the number of neutrons expressed in AMU.

• The atomic weight or atomic mass of an atom is the weighed average of all the mass numbers of the isotopes of that element. Both the mass number and the % abundance of each isotope of the element are taken into account to arrive at the atomic weight. For example, the element chlorine has a Cl-35 (75%) and a Cl-37 (25%) isotope. The atomic mass found in the periodic table for chlorine is 35.45.

Protons, Neutrons, Electrons

Particle Symbol proton p + Charge Mass +1 1 amu neutron n o electron e 0 1 amu -1 0 amu The number of protons = number of electrons.

The number of neutrons varies. The atomic number is the number of protons, the mass number is the number of protons plus neutrons.

Electrons (Total of 7) Bohr Model of the Atom Atomic weight = 14.0067 amu 14 N 7 Protons (total of 7) Neutrons (total of 7)

Bohr Theory of the Atom

• An electron can revolve around an atom only in specific allowed orbits.

• An electron in an orbit does not emit radiant energy as long as it remains in the orbit.

• An electron gains or loses energy only by moving from one allowed orbit to another. The energy an electron has depends on which orbit it occupies The only way that an electron can change its energy is to jump from one orbit to another in “quantum leaps”.

• An electron must acquire energy to jump from a lower orbit to a higher one, and the acquisition of energy must be sufficient for the jump to occur all at once.

• The energy that the electron acquires comes from high temperatures or from electrical discharges.

• An electron jumping from a higher to a lower orbit gives up energy in the form of light. A single level of the two orbits. This is a

photon quantum

is emitted when a downward jump occurs and the energy of the photon is exactly equal to the difference in the energy of energy.

A photon possesses a quantum of energy. The frequency is proportional to the amount of energy the photon possesses.

Bohr Orbits n=1 n=4 n=3 n=2 Nucleus The further away from The nucleus the higher the frequency and energy of the electron.

Quantum Mechanical Model of the Atom

• This is highly mathematical treatment of where the electrons are located in the atom.

• The Bohr model of the atom considered electrons as particles in circular orbits that could be only certain distances from the nucleus.

• The quantum mechanical model considers the electron as a wave and considers the energy of its harmonics.

• In the Bohr model the location of the electron was certain, in an orbit, in the quantum mechanical model the electron is a spread out wave and its location is uncertain.

• The Bohr model dealt with orbits, the quantum mechanical model deals with energy levels.

Quantum Mechanical Model of the Atom

• Quantum mechanics indicates the probability of finding an electron in a certain region in space at a given instant.

• There are shapes associated with the different energy levels where electrons may be.

90% of the time an electron in a particular energy level will be found in a particular region in space which has specific boundaries. Increasing energy levels have the same shape but occupy larger areas which reach further away from the nucleus.

The Periodic Table

• All the elements are located in the periodic table.

• The rows are called periods and are horizontal. There are rows 1-7 • The columns are called groups or families and are vertical. There are groups 1-8 A and the B groups.

• The table is arranged with increasing atomic numbers going left to right and coming around to the next row.

• Elements in the same group have similar chemical properties.

Periodic Table

• Elements are identified with a symbol consisting of 1 or 2 letters. The first one is always capitalized, the second one, if present, is lower case.

• The symbols may come from latin or other language.

• The number above the symbol is the atomic number of the element.

• The number below the symbol is the atomic weight of the element.

• The elements to the left of the staircase are metals and to the right are nonmetals.

• The majority of the elements are metals and solids.

Groups in the Periodic Table

• Group IA is called the alkali metals with the exception of hydrogen. These are the most reactive metals. They don’t occur by themselves in nature, only as compounds. They are called the alkali metals because they react violently with water to form alkaline or basic solutions. Sodium is an example of this. Group IA are all solids.

• Group IIA is called the alkaline earth metals and they are the second most reactive group of metals. Magnesium is an example of this. Group IIA are all solids.

• Group VIIA is called the halogens. These are the most reactive nonmetals. Fluorine and chlorine are greenish poisonous gases. Bromine is a reddish-brown liquid and iodine is a dark purple solid. They are used as disinfectants and have other uses as well. They react with metals to form salts. e.g. NaCl, sodium chloride, which is table salt.

Groups in the Periodic Table

• Group VIIIA is the noble gases. These are the most stable of all elements, not reactive at all, and they are all gases. • Groups IA-VIIIA are the representative elements and the B groups are the transition elements, or transition metals. The transition metals are all solids except for mercury, which is a liquid. Many of these metals are very common, such as copper, silver, and gold, iron. Transition metals are usually not very reactive and can sometimes be found in nature in mines as elements.

• The two rows of elements along the bottom of the periodic table are called the inner transition elements. The first row is the lanthanides or rare earths. The second row is the actinides. All elements after uranium (U) are manmade.

Metals, nonmetals, and Metalloids

• Metals are shiny, all are solids except for mercury and all conduct electricity. They can be drawn into wires (ductile) and can be bent (malleable). About 80% of the elements are metals.

• Nonmetals are dull, there are solids and gases and one liquid, bromine. They do not conduct electricity and the solids are brittle, they break rather than bend. Consider a piece of charcoal, made up of C atoms.

• Metalloids are the elements on either side of the staircase. They are semiconductors and they are all solids. They have properties intermediate between metals and nonmetals. They are used to make computer chips, e.g., silicon.

Electrons in Atoms

• n=1 (energy level 1) can have a maximum of 2 electrons.

• n=2 (energy level 2) can have a maximum of 8 electrons.

• n=3 (energy level 3) can have a maximum of 18 electrons • n=4 (energy level 4) can have a maximum of 32 electrons.

The Periodic Table

• Elements in the same group have the same number of electrons in the last energy level. These are called valence electrons.

• The number of valence electrons for the Group A elements is equal to the group number. e.g. C is in group IVA and has 4 valence electrons, Br is in group VIIA and has 7 valence electrons.

• The valence electrons determine the chemical properties and reactivity of the elements, and also some of the physical properties.

Electrons in Atoms

Period 1: H (1 e ) and He (2e ) n=1 only Period 2: n=1(2e ) and n=2 (1-8 e ) Period 3: n=1 (2e ), n=2 (8e ), n=3 (1-8e ) IA IIA IIIA IVA VA VIA VIIA VIIIA

Electrons in Atoms

Period 1: H (1 e ) and He (2e ) n=1 only Period 2: n=1(2e ) and n=2 (1-8 e ) Period 3: n=1 (2e ), n=2 (8e ), n=3 (1-8e ) All elements want 8e in last energy Level, like the noble gases (group VIII)

Electron dot structures

• The group number of the representative elements determines the number of electrons in the valence shell (valence electrons). These are the electrons in the highest energy level of the atom.

• Valence electrons can be represented using dot structures. • Atoms of representative elements form ions by losing or gaining electrons so that they will have 8 valence electrons, like the stable noble gases.

• Groups IA-IIIA will lose 1, 2, or 3 electrons so that they will go to the previous energy level and have 8 electrons there.

• Groups V-VIIA will gain 3, 2, or 1 electrons so that they will end up with 8 valence electrons.

Ions

• Ions are atoms that have gained or lost electrons so that the total number of electrons are no longer the same. The number of protons is still the same.

• Every time an electron is lost the charge becomes +1. Only metals will lose electrons. Every time an electron is gained the charge becomes -1. Only nonmetals will gain electrons. If more than one is lost or gained it will be an additional +1 or -1 for each.

• The elements in groups IA-IIIA form +1, +2, and +3 ions respectively. These are all metals. The group number gives you the positive charge of the ion formed.

• The elements in groups V-VIIA will form -3, -2, and -1 ions respectively. These are all nonmetals.

• 8-group number gives you the negative charge of the ion formed.

• The elements in group IVA will gain electrons if they are nonmetals and lose if they are metals. The metalloids do not tend to gain or lose electrons.

Ions formed by Representative Elements

Element Group Valence e Ion Formed Na IA 1 Na + Mg IIA 2 Mg 2+ Al IIIA 3 Al 3+ N VA 5 N 3 O VIA 6 O 2 Cl VIIA 7 Cl Ne VIIIA 8 None

Element Names to Know

• Groups IA, IIA, VIIA, and VIIIA • Periods 1, 2, 3, 4.

• Groups IB and IIB

Practice Exercises

• p. 226-228 Applying Concepts. # 2, 3, 4, 5, 6, 8, 9, 10, 13, 14, 15, 23, 24, 25, 26, 27, 28, 29, 30, 31.

• P. 229-230 Parallel Exercises Group A.

# 12, 13, 14, 15, 16, 17, 18, 19. New Book: p. 243-246 # 3, 6, 7, 9, 10, 11, 12, 13, 17, 19, 22, 23, 24, 25, 26, 27, 28, 29, 30, 33, 39, 40, 41, 42, 43, 44, 48, 49, 50

• • • • • • • •

Review for Chapter 8

Dalton’s atomic theory Thomson’s experiment: Discovery of electrons.

Rutherford’s experiment: The nucleus of the atom: protons and neutrons.

• • • Symbols, charges and masses for protons, neutrons and electrons.

Quantum Mechanical Model of the atom.

The periodic table: arrangement, groups, periods, metals, nonmetals and metalloids.

Atomic number, mass number and atomic mass. Atomic mass unit (AMU), #p + , #e , #n o .

• Electron dot structures for elements in groups IA-VIIIA, valence electrons.

Isotopes Symbols for atoms (isotopes) showing atomic number and mass number.

• • Numbers of electrons in energy levels 1-4.

Numbers of electrons in energy levels for the first 18 elements.

Bohr model of the atom.

Orbits, energy levels, electrons absorbing and emitting energy as a photon (a quantum of energy), and corresponding to the colors of the visible spectrum.

• • • Ions formed by elements in groups IA-VIIIA.

Metals, Semimetals and Nonmetals.

Alkali metals, alkaline earth metals, halogens, noble gases, representative elements, transition elements, inner transition elements, lanthanides (rare earths) and actinides.