Chemistry Chapter 5 Flashcards

Dmitri Mendeleev and the Periodic Table

• Dmitri Mendeleev is recognized as the "Father of the Periodic Table" for organizing elements by their properties.
• He utilized cards to categorize elements based on atomic mass and characteristics, revealing periodic patterns.

Classification of Elements

• Mendeleev classified elements by increasing atomic mass, leading to regular similarities in chemical and physical properties.
• His early periodic table was arranged by mass numbers instead of atomic numbers, which was a key distinction from modern tables.

Discovery by Henry Moseley

• Henry Moseley discovered that arranging elements by atomic number revealed a clearer pattern among them.
• This revelation established the modern periodic law.

Periodic Law and Table

• The Periodic Law states that the physical and chemical properties of elements are functions of their atomic number, with similar properties appearing at regular intervals.
• The current Periodic Table organizes elements by atomic number, grouping those with similar properties in columns.

Noble Gases and Specific Element Groups

• Noble gases, found in the last vertical column, include helium, neon, argon, krypton, xenon, and radon, known for their chemical inertness.
• Lanthanides are elements with atomic numbers 58 to 71, while actinides range from atomic numbers 90 to 103.

Periodicity

• Periodicity refers to the repeatable properties observed in groups on the periodic table due to electron arrangements.
• Horizontal rows are labeled periods, corresponding to the number of principal energy levels filled.

Metals, Nonmetals, and Metalloids

• Metals, located left of the staircase on the table, include alkali metals (Group IA) and alkaline earth metals (Group IIA).
• Nonmetals are to the right of the staircase, including halogens (Group VIIA) and noble gases (Group VIIIA).
• Metalloids share properties of both metals and nonmetals, positioned along the staircase.

Transition Elements

• Transition elements, found in the middle (Groups 3-12), exhibit unique properties due to unfilled d orbitals and can form various ion charges.

Atomic Structure and Trends

• Atomic radius measures half the distance between nuclei of bonded atoms, influenced by energy levels.
• Atomic radius decreases across a period (increased nucleus attraction) and increases down a group (more energy levels).

Ionization and Ionic Properties

• Ionization refers to the formation of charged ions through electron gain/loss.
• Ionization energy decreases down a group (increased distance from nucleus) and increases across a period (higher nuclear charge).
• Cations are positively charged ions, while anions are negatively charged ions.

Valence Electrons and Electronegativity

• Valence electrons in the outermost energy level dictate chemical behavior and bonding capacity, with their count relating to group numbers.
• Electronegativity, indicating an atom's ability to attract electrons, increases across a period and decreases down a group, with fluorine being the highest.

Element Blocks in the Periodic Table

• The "s" block contains Groups 1 and 2, primarily active metals with +1 or +2 oxidation states.
• The "p" block includes Groups 13-18, which mainly consist of non-metals, variable oxidation states, and acidic oxides.
• The "d" block (Groups 3-12) features transition metals known for variable oxidation states and colorful compounds.
• The "f" block (lanthanides and actinides) consists of inner transition elements with special properties such as radioactivity and unique chemical behaviors.