Science
1. Determine the number of valence electrons of the different
elements using the Periodic Table.

To determine the valence electrons of elements in Groups 1 or 2, the valence
electron is the same number as their group number.

For elements in groups 3 to 12 there is no rule to determine the valence electron,
you need to go by each energy level until it is reached the outermost energy level
with the electrons.

And for the elements in groups 13 to 18 they have 10 fewer valence electrons as
their group number. However, helium just has 2 valence electrons.
2. Describe Electronegativity and determine how it changes across the
Periodic Table.
Electronegativity is defined as how strongly an atom is able to tug on bonding
electrons. The greater the electronegativity of an element the greater its ability to pull
electrons towards itself when bonded.
Electronegativity increases from left to right in the periodic table. It also increases if it
goes up. Fluorine is the most electronegative element, the farther an element is the less
electronegative it is.
Note: Noble gases are not included as electronegative elements.
3. Explain the different concepts that made up the Quantum Theory.
Theories
 Max Planck: said that light energy is quantized the same way matter is. He also said
that light energy beams are not continuous steams of energy, instead they are small
packages of energy called quantum.

Albert Einstein: said that when energy hits a metallic surface, small packages of
energy are released, if this energy acts as light they are called photons.

Neils Bohr: An electron has more potential energy when it is farther from the
nucleus. When an atom absorbs energy from a photon, an electron is the one
acquiring the energy. Because this electron has gained energy it moves away from
the atom. When an electron loses energy it moves closer to the atom, and the
energy lost is emitted as a photon. An atom has a limited number of energy levels,
an electron moving around the nucleus is restricted to certain distances, determined
by the amount of energy it has.

Loius de Broglie: said that the slower an electron moves, the more its behavior is
that of a particle of mass; the faster it move, the more its behavior is that of a wave
energy. Electrons move as waves or as particles.
Electron excitation:
4. Explain why Noble Gases do not normally make chemical bonds with
other elements.
Noble gases usually don’t make chemical bonds because they have 8 valence electrons
and 8 is the perfect number, so they don’t need to receive or to give electrons.
5. Define the following terms: orbital, shells, electronic configuration,
photon, quantz, electron, ion, chemical bond, covalent, ionic and
metallic bonds, orbital diagram, energy levels, valence electrons,
Electronegativity.

Orbital: is an energy wave with a specific shape.

Shells: is a region of space about the atomic nucleus within which electrons may
reside.

Electronic configuration: is the form in which electrons distribute in the different
orbitals.

Photon: are small packages of energy that act as light.

Quantz:

Electron: a negatively charged particle in an atom.

Ion: an electrically charged particle created when an atom either loses or gains one
or more electrons.

Chemical Bond: the force of attraction between two atoms that holds them together.

Covalent bonds: a chemical bond in which atoms are held together by their mutual
attraction for two or more electrons they share.

Ionic bonds: is a bond that forms when electrons are transferred from one atom to
another atom.

Metallic bonds: is a bond formed by the attraction between positively charged metal
ions and the electrons in the metal.

Orbital Diagram: is a diagram using arrows to show the occupation of electrons on
each shell.

Energy levels: are levels in which electrons are distributed.

Valence electrons: are the electrons in the outer-most energy level and are the
responsible of chemical behaviour.

Electronegativity: the ability of an atom to attract a bonding pair of electrons to itself
when bonded to another atom.
6. Be able to determine the electron configuration (unabbreviated and
abbreviated) of different elements in the Periodic Table.
Steps to give the electronic configuration unabbreviated:
• Write down the number of electrons the element has
• Determine the orbitals it occupies (s-p-d or f)
• Determine the shells it occupies (1-7)
• Write the shell (level) and orbital (sub-level) and the number of electrons occupied
• S: space for 2
• P: space for 6
• D: space for 10 (when using d, decrease a level)
• F: space for 14 (when using f, decrease two levels)
Ex: 84Po 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p4
Steps to give the abbreviated configuration:
 Write in square brackets the last noble gas before the element
 Continue with the rest of the electronic configuration
Ex: [Xe] 6s24f145d106p4
7. With a given electron configuration, be able to determine the
element in the Periodic Table.
If it is in the unabbreviated way, read the configuration of the last level with the last
letter.
If it is in the abbreviated way reed from the last Noble gas until the last configuration.
Note: the last configuration is the element, not the next one, neither the last one.
8. Identify the different orbitals in the Periodic Table (s, d, f and p).
S-2 electrons
P- 6 electrons
D-10 electrons
F-14 electrons
9. Draw the orbital diagram of elements in the Periodic Table.
Steps:
 Determine the number of electrons the element has
 Locate the element in the periodic table
 Make the diagram using spaces and arrows
 Each space should only contain 2 electrons
 Always start with the up-arrows.
 For S it is just needed 1 space
 For P it is needed 3 spaces
 For D it is needed 5 spaces
 For F it is needed 7 spaces
Ex: 20Ca: 1S2 2S2 2p6 3S2 3p6 4s2
__↑↓_ __↑↓_ __↑↓_ __↑↓___↑↓_ __↑↓_ __↑↓_ __↑↓_ __↑↓_
__↑↓_
10. With a given orbital diagram, be able to determine the element in
the Periodic Table.
The only way to determine the element of a given orbital diagram is by counting the
spaces and arrows, if the last space is 1, its in S, if there are 3 spaces, its P; if there are
5 spaces the element is in D and if there are 7 spaces, the element is in F.
11. Determine the type of chemical bond, using the Periodic Table
(ionic, covalent and metallic).
To determine what type of chemical bond is the given one, using the periodic table you
should check if the 2 elements that make the bond are: metal or non metal. If both
elements are metals, the bond is metallic; if both are non metals it is a covalent; if one is
a non metal and the other one a metal it is an ionic bond.
12. Use the Lewis structures and electron dot structures for elements,
ions and compounds.
Lewis for ions:
 chemical symbol surrounded by the valence e structure placed within square brackets
 with a subscript indicate the change of an ion
Example:
Na+1: [Na.]
Cl -1: [ ]
Lewis for Ionic compounds (metal + non metal)
 The overal charge of the compound must be equal to zero, the # of e - gained by the
other atom
 The Lewis structure of each ion should work to build the L.S. of the compound.
Example: LiF
Li loses 1e- and form Li+1
F gains 1e- and form [ ]-1
LiF [Li.]+1 [
]-1
Lewis for covalent compounds (non metal + non metal)
 Electrons are shared between atoms to form a covalent bond; the # of e - depends to
provide stability.
 Electrons in the L.S. are paired to show the bonding pair of e-.
 Sometime it is circled.
__:Shared electrons
Example: H2O;
..
..
H O .. H
..
..
H__O__
H
..
13. Clasifique los compuestos de acuerdo al número de elementos
que contiene (binarios, terciarios y cuaternarios).
Clasificación:
 Binarios: (2 elementos) Ej:H2O
 Terciarios: (3 elementos) Ej: C6H12O6 (glucosa) se pueden llamar ternarios también
 Cuaternarios: (4 elementos) Ej: compuestos orgánicos,
-NH3
14. Clasifique los compuestos de acuerdo a la composición que tienen
(hidruros, óxidos, hidrácidos, hidróxidos, sales).
Clasificaciones:






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Hidruros: (metal + hidrógeno). Ej.: KH
Hidrácido: (hidrógeno + no metal). Ej.: HCl (ac): disuelto en agua
Hidróxido: (metal +OH). Ej.: NaOH
Oxido Metálico: (metal + oxigeno). Ej.: NaO
Oxido no metálico: (no metal + oxigeno). Ej.: CO2
Sal binaria: (metal + no metal). Ej.: KF
Sal no binaria: (no metal + no metal) que no sea con H2 u O2. Ej.: CN