Periodic Table and Elements

Periodic Table and Elements

• The periodic table organizes elements by atomic number, electron arrangement, and chemical properties.
• Elements are pure substances with one type of atom, represented by a unique symbol, possessing distinct chemical properties.
• There are 118 elements on the periodic table.
• The first 94 elements occur naturally, while the remaining 24 are synthetic.

Compounds and Classification of Elements

• A compound is defined as a pure substance composed of two or more different elements.
• Metals are good conductors of electricity and heat, typically shiny, malleable, and ductile.
• Nonmetals are poor conductors, brittle in solid form, and exist in various states.
• Metalloids exhibit properties of both metals and nonmetals.
• Metalloids are located along the stair-step line, also known as a zig-zag line, on the periodic table, between metals and nonmetals.
• Metals can be found to the left of the staircase line.
• Nonmetals can be found to the right of the staircase line.
• The elements that are classified as Metalloids are Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po), and Astatine (At).

Chemical Families

• Chemical families are columns of elements with similar properties on the periodic table.
• Alkali Metals (Group 1) include Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), Francium (Fr).
• Alkali Metals are reactive with water, soft, have low melting points, and possess one outer electron, forming +1 ions.
• Reactivity increases as you move down the group of Alkali Metals.
• Alkaline Earth Metals (Group 2) include Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), Radium (Ra).
• Alkaline Earth Metals are reactive, but less so than alkali metals, harder, have higher melting points, and possess two outer electrons, forming +2 ions.
• They react with water, but not as violently as alkali metals.
• Noble Gases (Group 18) include Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn).
• Noble gases are non-reactive (inert) due to full outer electron shells, making them stable, colorless, and odorless at room temperature.
• Halogens (Group 17) include Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).
• Halogens are reactive, especially with alkali and alkaline earth metals.
• They have seven outer electrons, one short of a full shell, and form salts with metals.
• Reactivity decreases as you go down the group of Halogens.

Periodic Trends

• Elements in the same horizontal row (period) show trends of increasing or decreasing reactivity.
• Group 1 alkali metals are more reactive than their Group 2 neighbors.
• Group 17 halogens are more reactive than Group 16 elements in the same row.

History of the Periodic Table

• The periodic table was originally developed in 1869 by Dmitri Mendeleev.
• There were only 63 known elements when it was developed,.
• Mendeleev arranged elements by increasing mass, and organized elements with similar properties in the same column.
• Mendeleev predicted the existence and properties of undiscovered elements, which were later found to match his predictions.

Scientific Theory

• A scientific theory is defined as the best understanding of a phenomenon, based on scientific evidence or reasoning. It is not a guess.

Democritus' Atomic Theory

• Democritus proposed that matter could be divided into smaller pieces until an indivisible particle, called an "atom," was reached.
• Democritus suggested that without any experimental evidence atoms possess different sizes, are in constant motion and are separated by empty spaces.

Aristotle's Theory

• Aristotle rejected the idea of the atom.
• He believed all matter is made up of four basic substances: Earth, Water, Air, Fire.
• He believed all matter possess the qualities of: dry, wet, cold, and hot.
• This theory was accepted for almost 2000 years.

Dalton's Atomic Theory (Billiard Ball Model)

• John Dalton proposed that all matter is made of tiny, indivisible atoms.
• All atoms of an element are identical.
• Atoms of different elements are different.
• Atoms are rearranged in chemical reactions but are never created or destroyed.

Thomson's Atomic Theory (Plum Pudding Model)

• J.J. Thomson discovered the electron through cathode ray tube experiments
• J.J. Thomson's model proposed that atoms contain negatively charged electrons evenly distributed within a positively charged sphere.

Rutherford's Atomic Theory (Nuclear Model)

• Ernest Rutherford conducted the gold foil experiment, which led to the discovery of a small, dense, positively charged nucleus.
• Rutherford proposed that the center of the atom possesses a positive charge, called the nucleus.
• The nucleus contains most of the atom’s mass but occupies a very small space.
• The nucleus is surrounded by a cloud of negatively charged electrons, most of the atom is empty space.

Bohr's Model

• Electrons orbit the nucleus like planets around the sun.
• Each electron in an orbit has a definite amount of energy.
• The farther the electron is from the nucleus, the greater its energy.
• Electrons cannot reside between orbits, but they can jump between orbits, releasing energy as light.
• Each orbit can hold a certain maximum number of electrons.
• The maximum number of electrons in the first, second, and third orbits is 2, 8, and 8, respectively.

Chadwick's Discovery

• James Chadwick discovered the neutron in 1932.
• Chadwick proposed an atom with a tiny, dense nucleus containing positively charged protons and neutral neutrons.
• The mass of a neutron is about the same as that of a proton.
• Negatively charged electrons circle rapidly through the empty space around the nucleus.
• A neutral atom has the same number of protons as electrons.

Atomic Structure

• Atoms consist of protons, neutrons, and electrons.
• Each element has a unique number of protons, neutrons, and electrons, giving it unique properties.
• Atomic Number = # of protons.
• Mass Number = # of protons + # of neutrons.
• To find the number of neutrons, subtract the atomic number from the mass number.
• Isotopes are two or more forms of an element with the same number of protons but different numbers of neutrons.
• Isotopes of an element have the same chemical and physical properties.
• All atomic masses listed on the periodic table are averages taken from the percent abundance of each isotope of an element

Molecules and Compounds

• A molecule is a group of two or more atoms bonded together, like water (H₂O).
• A chemical formula shows the elements in a compound and the ratio of atoms, such as CO₂ for carbon dioxide.
• A molecular element consists of molecules made up of only one type of atom, e.g., oxygen (O₂).
• Molecular compounds are formed when two or more different atoms bond together, e.g., carbon dioxide (CO₂).

Ions

• An ion is defined as an atom or molecule with an electric charge due to the loss or gain of electrons.
• A cation is a positively charged ion, for example, Na⁺ (sodium ion).
• An anion is a negatively charged ion, for example, Cl⁻ (chloride ion).
• Ionic compounds are formed from the electrostatic attraction between cations and anions, e.g., NaCl (sodium chloride).
• Atoms combine to form compounds to achieve full outer energy levels, similar to noble gases, through gaining, losing, or sharing electrons, which is also referred to as forming chemical bonds.

Ionic Bonds

• Ionic bonds occur between a metal and a nonmetal.
• Metals lose electrons to become positive ions (cations).
• Non-metals gain electrons to become negative ions (anions).
• Opposite charges attract each other and form a bond
• Example: Sodium (Na) + Chlorine (Cl) → Sodium Chloride (NaCl).

Covalent Bonds

• Covalent (Molecular) Bonds occurs between two non-metals.
• Atoms share electrons to fill their outer shell
• Example: Hydrogen (H) + Oxygen (O) → Water (H₂O).
• Atoms form bonds to become more stable, aiming to have a full outer shell like noble gases.