Ch 2. Periodic Properties
Chapter 2: Periodic Properties
Introduction to the Periodic Table
The periodic table is an organized chart that presents all known chemical elements, arranged by their atomic number, electron configuration, and recurring chemical properties. It serves as a valuable tool in chemistry, providing insights into the relationships between different elements and their behaviors.
Periodic Properties Overview:
This chapter discusses significant trends relevant to the Dental Admission Test (DAT), including:
Atomic radius: The size of an atom, which influences various physical and chemical properties.
Ionization energy: The energy required to remove an electron from an atom, indicating how tightly electrons are held.
Effective nuclear charge: The net positive charge experienced by valence electrons, which affects atomic size and ionization energy.
Electron affinity: The energy change when an electron is added to an atom, reflecting the atom's tendency to gain electrons.
Electronegativity: The measure of an atom's ability to attract and bond with electrons in a chemical bond, affecting molecular formation and behavior.
Key Terms:
Groups: Vertical columns in the periodic table where elements share similar chemical properties due to having the same number of valence electrons.
Periods: Horizontal rows in the periodic table representing elements with the same number of electron shells.
Atomic Number: Represents the number of protons in an atom's nucleus; for example, nitrogen has an atomic number of 7, indicating it has 7 protons and 7 electrons in its neutral state.
Groups Explained:
Group 1 (excluding hydrogen): Known as alkali metals, these elements (like lithium and sodium) are characterized by their high reactivity and formation of strong bases.
Group 2: Composed of alkaline earth metals, such as magnesium and calcium, which are also reactive but less than alkali metals.
Groups 3-12: Encompasses transition metals like iron and copper, which have variable oxidation states and exhibit complex behaviors in bonding.
Groups 13-17: Includes metalloids (like silicon) that have properties intermediate between metals and non-metals, and non-metals like carbon and oxygen.
Group 17: Contains halogens such as fluorine and chlorine, known for their high reactivity, especially with alkali metals.
Group 18: Noble gases like helium and neon, recognized for their lack of reactivity due to a full valence shell.
Oxidation States:
These states relate to the electrons that an atom gains, loses, or shares during bonding and chemical reactions. Transition metals frequently display multiple oxidation states; for instance, manganese can exhibit values ranging from +2 to +7, influencing its chemical behavior.
Inner Transition Metals
Period 6: Lanthanides, which include elements like cerium and neodymium, are known for their applications in technology, including catalysts and magnets.
Period 7: Actinides, consisting of elements like uranium and plutonium, are known for their radioactive properties and use in nuclear energy.
Diatomic Atoms and Metal Characteristics
Diatomic Atoms: These are often unstable as single atoms and are commonly found bonded in pairs. Key diatomic elements include hydrogen (H₂), nitrogen (N₂), fluorine (F₂), oxygen (O₂), iodine (I₂), chlorine (Cl₂), and bromine (Br₂).
Metallic Character Trends: Generally, metallic characteristics increase as you move from the right to the left across a period, and from top to bottom in a group.
Properties of Metals vs. Non-Metals:
Metals: Typically exhibit malleability, ductility, and good conductivity of electricity and heat. They tend to react with acids and form basic oxides. In bonding, metals generally lose electrons, forming positive ions (cations).
Non-Metals: Usually brittle, dull in appearance, and poor conductors of heat and electricity. They tend to form acidic oxides and typically gain electrons in chemical reactions, forming negative ions (anions).
Periodic Trends
Atomic Radius
Definition: Defined as half the distance between the nuclei of two identical atoms bonded together, the atomic radius gives insight into the size of an atom.
Trends: The atomic radius increases from right to left across a period due to a decrease in effective nuclear charge pulling electrons closer, and increases down a group as additional electron shells are added.
Effective Nuclear Charge
Definition: It represents the net positive charge experienced by electrons due to the shielding effect of inner-shell electrons.
Trends: The effective nuclear charge increases from left to right across a period, as protons are added without increasing electron shielding, and increases up a group as there are fewer inner shells providing shielding.
Isoelectronic Series
Definition: A group of atoms or ions that have the same electron configuration but different numbers of protons.
Properties: When comparing isoelectronic species, anions (which gain electrons) will have a larger radius than their neutral atoms, while cations (which lose electrons) will have a smaller radius.
Ionization Energy
Definition: The energy required to remove an electron from a gaseous atom, reflecting the atom's strength to hold onto its electrons.
Trends: Ionization energy increases from left to right across a period due to a higher effective nuclear charge, and increases going up a group as atom size decreases, leading to stronger attraction between nucleus and electrons.
Electron Affinity
Definition: The energy change that occurs when an electron is added to a neutral atom in the gaseous state. Positive values indicate a release of energy, while negative values signify energy input required for electron addition.
Trends: Electron affinity generally increases from left to right across a period as nuclear charge increases, leading to a greater attraction for added electrons, and increases up a group because of reduced shielding allowing stronger attraction by the nucleus.
Electronegativity
Definition: A measure of an atom's ability to attract and hold onto electrons when bonded to another atom.
Trends: Electronegativity increases from left to right across a period due to rising nuclear charge, and up a group due to smaller atomic size which allows electrons to be more tightly attracted. The most electronegative element is fluorine.
Summary of Key Trends
Atomic Radius: Increases right to left and down.
Effective Nuclear Charge: Increases left to right and up.
Ionization Energy: Increases left to right and up.
Electron Affinity: Increases left to right and up.
Electronegativity: Increases left to right and up.