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Chemistry I
Unit III
The atom & the periodic table
Ms. Claudia Barahona
October 2014
Stage 3
• Development of atomic theories
• Subatomic particles
• Electron configuration
• Organization of elements on the
periodic table
• Periodic properties (trends):
– Atomic size
– Ionization energy
– Electronegativity
Activity 3.1 The atom
• What is the atom composed of?
• Composed by subatomic particles:
• Protons : p+
• Electrons: e-
• Neutrons: n
Subatomic particles
• Particles that are smaller than the atom.
• Protons and neutrons make up the nucleus of an atom.
Note: Amu (Atomic Mass Unit) is defined as one-twelfth of the mass of
the carbon atom with six protons and six neutrons.
Name Symbol Elctrical charge Mass (amu)
Proton
Electron
Neutron
Name Symbol Elctrical charge Mass (amu)
Proton p+
Electron e-
Neutron n
Name Symbol Elctrical charge Mass (amu)
Proton p+ +1 1
Electron e- -1 0.000549
Neutron n 0 1
• Atomic number:
Is equal to the number of p+ in an atom.
• Mass number:
Is equal to the number of p+ and n in
the nucleus of an atom.
• Atomic mass or atomic weight:
Weight average mass of all the
natural occuring isotopes of an element
Information in the periodic table
Smallest particle of an element that retains the characteristis of
the element.
Atom
Atomic theories: Theories that try to explain the structure of the atom:
• Dalton
• Thomson
• Bohr
• Ruhtherford
• Schrodinger
• Decomposition process in which unstable atomic nuclei will spontaneously decompose to form nuclei with a higher stability, resulting in a release of high energy radiation.
Radioactivity
Radioactivity
Benefits Risk
Used in nuclear medicine. Radiotheraphy Irradiation
Can burn the surface of the skin. In contact with cells in the body may cause genetics mutations.
Fire detective systems Generates wastes which we do not know how to manage or destroy. Nuclear power stations
• Ex: P-32 (Tx of Leukemia), C-14 (Radiocarbon dating),
Au-198 (Liver imaging and carcinoma), I-123 (Thyroid, brain and porstate cancer).
Modern periodic table
• Period: Each horizontal row.
• Group/family: Each vertical column, this elements will have similar properties.
Modern periodic table
Representative elements
Transition Metals
Periodic trends
Representative elements
Transition Metals
Acquisition of knowledge
• Isotopes:
Are atoms of the same element that have different mass number but the same chemical behavior.
Atomic symbol for writing Isotopes: AZX X= chemical symbol for the element Z= Atomic number A= Mass number
• Atomic mass or atomic weight: Weight average
mass of all the natural occuring isotopes
of an element
• Atomic number: Is equal to the number
of protons in an atom.
Information in the periodic table
Representative elements
Transition Metals
Exercise # 1 Subatomic particles
• Is each of the following statements true or false?
If false, explain your reason.
a. Protons are heavier than electrons.
True
b. Protons are attracted to neutrons.
False. p+ are attracted to e-
c. Electrons are so small that they have no electrical charge.
False. e- have a -1 charge
d. The nucleus contains all the protons and neutrons of an atom.
True
Exercise # 2 Subatomic particles
• Using the periodic table, state the atomic number, number of protons, and number of electrons for an atom of each of the following elements:
Element Atomic
number p+ e-
Nitrogen
Magnesium
Bromine
Representative elements
Transition Metals
Subatomic particles
• Using the periodic table, state the atomic number, number of protons, and number of electrons for an atom of each of the following elements:
Element Atomic
number p+ e-
Nitrogen 7 7 7
Magnesium 12 12 12
Bromine 35 35 35
Exercise # 3 Subatomic particles
• Consider an atom that has 79 electrons.
a. How many protons are in its nucleus?
b. What is its atomic number?
c. What is its name, and what is its symbol?
Representative elements
Transition Metals
Subatomic particles
• Consider an atom that has 79 electrons.
a. How many protons are in its nucleus?
79
a. What is its atomic number?
79
a. What is its name, and what is its symbol?
Subatomic particles
• The number of protons gives atoms their identity.
• Atomic number (Z)
#of protons
• Mass number (A)
• #protons + #neutrons
Information in the periodic table
• Atomic mass or atomic weight: Weight average
mass of all the natural occuring isotopes
of an element.
• Isotopes:
Are atoms of the same element that have different mass number but the same chemical behavior.
Atomic symbol for writing Isotopes: AZX X= chemical symbol for the element Z= Atomic number A= Mass number
Identify protons and neutron in isotopes
State the number of p+, e- and n in each of the following isotopes of Carbon (C).
Exercise# 4 Isotopes
Isotope Protons Electrons Neutrons Protons Electrons Neutrons
6 6 6
6 6 7
6 6 8
Identify protons and neutron in isotopes
Write the symbol for each of the following isotopes:
a) A nitrogen atom with 8 neutrons.
b) An atom with 20 protons and 22 neutrons.
c) An atom with mass number 27 and 14 neutrons.
Exercise# 5 Isotopes
Representative elements
Transition Metals
Write the symbol for each of the following isotopes:
a) A nitrogen atom with 8 neutrons.
157N
a) An atom with 20 protons and 22 neutrons. 42
20Ca
a) An atom with mass number 27 and 14 neutrons.
2713Al
Exercise# 5 Isotopes
Review activity 3.2
Atomic theories
• Theories that try to explain the
structure of the atom.
– John Dalton
– J. J. Thomson
– Ernest Rutherford
– Niels Bohr
– Erwin Schrodinger
Evolution of the atomic theories
Billiard ball model- a small, solid sphere
Evolution of the atomic theories
Law of Conservation of mass:
States that the total mass present before a chemical reaction
is the same as the total mass present after the chemical
reaction; thus, mass is conserved. The law of conservation
of mass was formulated by Antoine Lavoisier (1743-1794).
This law was a result of his combustion.
Evolution of the atomic theories
Law of Constant composition or Law of definite
proportions:
Formulated by Joseph Proust (1754-1826).
States that if a compound is broken down into its constituent
elements, the masses of the constituents will always have
the same proportions, regardless of the quantity or source of
the original substance.
Evolution of the atomic theories
Law of Constant composition or Law of definite
proportions:
Jon Berzelius did experiments with about 2000 compounds
Berzelius prepared and purified the necessary reagents,
developed the techniques to perform the analyses, and
collected data on the relative weights of atoms of 43
elements.
That confirmed John Dalton's atomic theory as well as
Proust's law showing that separate elements always
combined in whole-number proportions.
*Also introduced the symbolism with which chemical
formulas are still written
Evolution of the atomic theories
How was Dalton wrong in his proposal?
Not all atoms of the same element are exactly alike
(isotopes).
Atoms are made up of subatomic particles.
J.J. Thomson
• Discovered the electron in a series of experiments using cathode-ray tube.
• In 1904 Thomson suggested a model of the atom as a sphere of positive matter in which electrons are positioned by electrostatic forces.
“Plum-pudding”
model
Ernest Rutherford • Rutherford performed a series of experiments
with radioactive alpha particles.
• He found that while most of the alpha particles passed right through the gold foil, a small number of alpha particles passed through at an angle (as if they had bumped up against
something) and some bounced straight
back
• Rutherford's experiments suggested
that gold foil, and matter in general, had
holes in it!
Niels Bohr
• Devised the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus—similar in structure to the solar system.
Erwin Schrodinger
• Schrödinger model describes the probability that an electron can be found in a given region of space at a given time. This model no longer tells us where the electron is; it only tells us where it might be.
• Introduced “wave
mechanics” as a
mathematical model.
Evolution of the atomic theories
Evolution of the atomic theories
Electronic configuration
Energy level: Specific energy that an electron has (bound by the electric field of the nucleus).
e- in the lower energy levels are usually closer
to the nucleus.
Sublevel: Group of orbitals of equal energy within principal energy level.
Orbital: Region around the nucleus where e-s of a certain energy are more likely to be found. (s, p,d and f)
Electronic configuration
List of the number of electrons in each sublevel within an atom, arranged by increasing energy.
Electronic configuration
Orbital´s maximum number of e-
• S= 2
• P= 6
• d= 10
• f = 14
Electronic configuration
Representative elements
Transition Metals
Electronic configuration
• To contruct the electronic configuration of an atom do the following:
1. Determine the number of electrons in the atom.
2. Put electrons moving from the lowest energy levels to the highest energy orbital available, starting with 1s (holds a maximum of two electrons).
3. Fill in the orbitals according to the number of electrons in the atom.
Electronic configuration
• Example: Write the electronic configuration of Lithium atom.
1. Electrons involved:
atomic number 3
2. Begin with the 1s sublevel.
1s2
3. Fill in the needed orbitals, until all the electrons are being positioned.
1s2 2s1
Electronic configuration
• Write the electron configuration for :
a. Nitrogen atom
b. Silicon atom
c. Chlorine atom
Representative elements
Transition Metals
Electronic configuration
• Write the electron configuration for :
a. Nitrogen atom
1s2 2s2 2p3
b. Silicon atom
1s2 2s2 2p6 3s2 3p2
c. Chlorine atom
1s2 2s2 2p6 3s2 3p5
Electronic configuration
• List of the number of electrons in each sublevel within an atom, arranged by increasing energy.
• Chemistry book chapter 3 Pg 81
Orbital diagram
• Boxes represent the orbitals and half arrows represent electrons.
Electronic configuration
Maximum number of e-
• S= 2
• P= 6
• d= 10
• f = 14
Aufbau´s principle
• e-s fill orbitals starting at the lowest available energy state before filling higher states (1s before 2s).
Pauli exclusion principle
• States the an orbital can hold up to maximum of 2 e-´s, which are seeing as spinning on its axis, which generates a magnetic field.
• An orbital can hold 0, 1, or 2 electrons only, and if there are two electrons in the orbital, they must have opposite (paired) spins.
Hund´s rule
• When filling sublevels other than s, electrons are placed in individual orbitals before they are paired up.
Representative elements
Transition Metals
Blocks in the periodic table
Valence electrons
• Electrons in the outermost energy level.
• Given by the group number (representative elements).
Valence electrons
• Electrons in the outermost energy level.
• Given by the group number (representative elements).
Oxidation number
• Shows the total number of e-´s which have been removed from an element (+) or added to an element (-).
Periodic trends
• Elctron configuration of atoms are an important factor in physical and chemical properties of the elements.
• Periodic properties increases or decreases across a period, and then the trend is repeated again in each successive group.
Periodic trends
• Electron afinity: The ability of an atom to attract additional electrons.
• Electronegativity: The relative ability of an element to attract electrons in a bond.
• Ionization energy: Energy needed to remove the least tight bound electron from an atom in gaseous (g) state.
• Atomic radius: Distance from the nucleus to the energy level that contains the valence (outermost) electrons.
Periodic trends
Writing formulas
1. Identify the cation and anion or polyatomic ion.
2. Balance the charge.
3. Write the formula, cation first, using the subscript from the charge balance.
Subscripts in formulas
• The subscripts in the formula represent the number of positive and negative ions that give an overall charge of zero.
Check activity 3.3
Check End of chapter